A THERMODYNAMIC AND KINETIC INVESTIGATION OF B UNE

FORMATION FROM ISOBUTYRALDEHYDE AND PRIMARY ALKYLAMINES

A DISSERTATION

Presented in Partial Fulfillment of the Requirements for

the Degree Doctor of Philosophy in the Graduate School

of The Ohio State University

By

Francis Anthony Via

The Ohio State University 1970

Approved by

Advisor Department of Chemistry PLEASE NOTE;

Some pages have small and indistinct type. Filmed as received.

University Microfilms ACKNOWLEDGMENT

I wish to thank Dr. Jack Hine for his patient guidance, informative discussions, and helpful suggestions during the course of this research. I wish to thank the appropriate government a- gencies and this Department of Chemistry for financial aid during the execution of this research. It is also appropriate to thank all my laboratory colleagues and other associates for their con­ structive influence.

IX VITA

November 50, 19^5 • ...... Born, Frostburg, Maryland.

June, 1961...... Graduated Beall High School, Frostburg, Maryland.

June, 1965 ...... Graduated West Virginia Univer­ sity, Morgantown, West Virginia.

September, I9 6 5 ...... Entered The Ohio State University as a Teaching Assistant

August, 1 9 6 8 ...... • Received M.S. Degree, The Ohio State University, Columbus, Ohio.

1x 1 TABLE OF C0ETEBT8

Page

ACKWOWLEDOyiENTS...... il

VITA ...... ill

LIST OF TABLES...... vi

LIST OF FIGUEES...... x

Chapter

I. INTRODUCTION...... 1

II. HISTORICAL...... 6

Kinetics of Imine Formation......

Reported Bifunctional Catalysis......

Hydration of Isobutyraldéhyde......

The Micro-pKa’s of the Diamines......

III. EXPERIMENTAL...... 5 6

General......

IV. RESULTS ...... 46

Hydration of Isobutyraldéhyde......

Determination of the Micro-pKa Values of the Diamines

Imine Formation......

V. DISCUSSION...... 1 9 9

Hydration of Isobutyraldéhyde......

The Micro-pKa Values ......

Imine Formation......

iv TABLE OF CONTENTS (Cont'd.)

Page

VI. CONCLUSION ...... 230

APPENDIX ...... 231»-

LITERATURE CITED...... 311 LIST OF TABLES

Page

1. Estimated Structure of Monoprotonated N,N-Dimethyl- $4 ethylenediamine.

2. observed Rate Constants for the Hydration-Dehydration 48 Reaction.

3- Results for the Hydration-Dehydration Reaction with 52 Isobutyraldehyde.

4. Catalytic Constants for Hydration of Isobutyraldéhyde. 53

5 . Rate Constants for Dehydration of Isobutyraldéhyde 55 Hydrate.

6. Chemical Shift of Protonated Methylamine. 6l

7. Chemical Shift of Protonated Trimethylamine. 63

8. Chemical Shift of Protonated Triethylamine. 65

9 . Chemical Shift of Protonated n-Propylamine. 67

10. Chemical Shift of Protonated 2-Methoxyethylamine . 69

11. Chemical Shift of Protonated ,N'-Tetramethyl- ' 71 ethylenediamine.

12. Chemical Shift of Protonated N,H,H’,N’-Tetramethyl- 73 1,3 -propane diamine.

1 3 . Chemical Shift of Protonated H -Dimet hy let hylene- 75 diamine.

14. Chemical Shift of Protonated -Dimethyl-1,5-propane- 78 diamine.

1 5 . Chemical Shift of Protonated N-Dimethyl-1,4-butane - 8l diamine.

1 6 . Chemical Shift of Protonated H,H-Dimethyl-1,5-pentane- 84 diamine.

Vi Page

17- observed and Estimated Chemical Shift Values. 88

1 8 . Evaluation of f for MesN( CHg)« 90

19* Micro-pKa Values for the Diamines. 92

20. Rate and Equilibrium Constants for the Reaction of 1Q5 n-Propylamine and Isobutyraldéhyde.

21. Equilibrium Constants for the Reaction of 3-Methoxy- 106 propylamine and Isobutyraldéhyde.

22. Rate Constants for the Reaction of 3-Methoxypropylamine 107 and Isobutyraldéhyde.

23 . Equilibrium Constants for the Reaction of 2-Methoxy- 109 ethylamine and Isobutyraldéhyde.

24. Rate Constants for the Reaction of 2-Methoxy ethyl- 110 amine and Isobutyraldéhyde.

2 5 . Equilibrium Constants for the Reaction of 2,2- 113 Dimethoxyethylamine and Isobutyraldéhyde.

2 6 . Rate Constants for the Reaction of 2,2-Dimethoxy- ll4 ethylamine and Isobutyraldéhyde.

2 7 . Equilibrium Constants for the Reaction of 2,2,2- , II6 Trifluoroethylamine and Isobutyraldéhyde.

2 8 . Rate Constants for the Reaction of 2,2,2-Trifluoro- 117 ethylamine and Isobutyraldéhyde.

2 9 . Equilibrium Constants for Formation of Carbinolamines II9 and Imines from Isobutyraldéhyde and Aliphatic Primary Monoamines.

3 0 . Evaluation of Linear Functions of the Form y = mx + d. 123

3 1 . Rate Constants for Monoamines. 125

3 2 . Equilibrium Constants for Imine Formation "with n,K- 132 Dimethylethylenediamine.

35- Equilibrium Constants for Carbinolamine Formation 135 ■with UjR-Dimethylethylenediamine.

vii Page

Rate Constants for Imine Formation with N,N-Dimethyl- 158 ethylenediamine.

55* Equilibrium Constants for Imine Formation with l4l Dimethyl-1,5 -propane diamine.

3 6 . Equilibrium Constants for Carbinolamine Formation l44 with Ef ,N-Dimethyl-1,3 -propanediamine.

31. Rate Constants for Imine Formation with -Dimethyl- l46 1,3 -propanediamine.

5 8 . Equilibrium Constants for Imine Formation with l48 Dimethyl-1,4-butanediamine.

39' Equilibrium Constants for Carbinolamine Formation 151 with N,N-Dimethyl-l,4-butanediamine.

40. Rate Constants for Imine Formation with E,E-Dimethyl- 155 1 .4-butanediamine.

41. Equilibrium Constants for Imine Formation with E,E- 155 Dimethyl-1,5-pentanediamine.

4 2 . Equilibrium Constants for Carbinolamine Formation 158 with E ,E-Dimethyl-1,5 -pent ane diamine.

4 3 . Rate Constants for Imine Formation with E,E-Dimethyl- 160 1 .5 -pentanediamine.

44. Calculation of Taft Substituent for Diamines. 162

45 . Calculation of Equilibrium Constants for Protonated l64 Diamines.

46. Equilibrium Constants of Diamines. I69

4 7 . Calculation of the Ionization Constants for Mono- 176 protonated Carbinolamine.

48. Estimation of the Ionization Constants for Protonated 178 Carbinolamines of the Diamines.

4 9 . Calculated Rate Constants for Imine Formation from 179 Isobutyraldéhyde and the Four Diamines.

5 0 . Greatest Percent Contribution of Each Rate Term to 18O Observed Rate Constants. viii Page

51- Evaluation of kg^. I98

5 2 . Rate Data for Acid-Base Catalyzed Hydration of Iso- 200 butyraldéhyde.

55r Correlation of Internal Acid-Catalysis. 226

5^. Major Infrared Bands of the Diamines. 242

55 • Example of Kinetic Run for Hydration of Isobutyral- 245 dehyde.

5 6 . Chemical Shifts of Protonated Trimethylamine. 245

57- Chemical Shifts of Protonated N-Dimethylethylene - 247 diamine.

58 . Data for Run Number 210.805 • 250

59* Data for Run Number 210.664 . 255

6 0 . An Illustration of the Sensitivity of Kq to Change 260 in Percent Transmittance.

IX LIST OF FIGURES

Page

1. Intramolecular Acid Catalysis. 5

2. Logarithm of the First-Order Rate Constants for the 10 Hydrolysis of Substituted Benzylidene-1,1-dimethyl- ethylamine as a Function of pH.

3. Enzymic Catalysis of Transamination. IT

A Plot of the Observed Rates of Hy dr a t ion -Lehydrat i on 50 for Isobutyraldéhyde vs. the Total Buffer Concentration.

5* . A Plot of the General Base Catalyzed Rate Constant for 5^ Dehydration vs. the pKa of the Protonated Catalyst.

6 . Correlation of Chemical Shift "with Percent Unprotonated 62 Methylamine.

T • Correlation of Chemical Shift with Percent Unprotonated 64 Trimethylamine.

8 . Correlation of Chemical Shift with Percent Unprotonated 66 Triethylamine.

9- Correlation of Chemical Shift with Percent Unprotonated 68 n-Propylamine.

10. Correlation of Chemical Shift with Percent Unprotonated 70 2-Methoxyethylamine.

11. Correlation of Chemical Shift with Percent Unprotonated 72 ITjN ,N ’ '-Tetramethylethylenediamine.

12. Correlation of Chemical Shift with Percent Unprotonated 74 N ,11 ’ -Tetrame thyl-1,5 -propanediamine.

13. Correlation of Chemical Shift with Percent Unprotonated jS U,H-Dimethylethylenediamine.

14. Evaluation of Equation 42. 77 Page

15. Correlation of Chemical Shift •with Percent Unprotonated 79

1 6 . Evaluation of Equation 4-2. 80

1 7 . Correlation of Chemical Shift with percent Unprotonated 82 U -Dimethyl-1 -hut ane.

1 8 . Evaluation of Equation ^2. 83

19- Correlation of Chemical Shift -with Percent Unprotonated 85 U -Dimethyl-1,5 -pentanediamine.

2 0 . Evaluation of Equation 42. 86

2 1 . Plot of pKa vs. log Kj for Monoamines. 120

2 2 . Plot of pKa vs. log Kq of Monoamines. 121

2 5 . Plot of a* vs. log for Monoamines. 122

2h. Plot of pKa vs. log ko for Monoamines. 127

2 5 . Plot of cr* vs. log ko for Monoamines. 128

2 6 . Plot of pKa vs. log kg for Monoamines. 129

2 7 . Plot of pKa vs. log Kgko for Monoamines. 130

2 8 . A Plot of pH vs. Kjq-^s for MesNCCHs)aNHg- 133

2 9 . A Plot of pH vs. KqoIjs for Mes^CCH2)2KH2 * 136

3 0 . A Plot of pH vs. Kjobs Me2M(CH2)sHH2 . 142

3 1 . A Plot of pH vs. for Me2N(CHg)3l3H2 . 145

3 2 . A Plot of pH vs. for Me2N(CH2)4HH2 . 149

3 3 . A Plot of pH vs. KgQ^g for Me2N(CHg)4HH2 - 152

3^. A Plot of pH vs. KjQijg for Me2N(CH2)5UH2 - 156

3 5 . A Plot of pH vs. KQQi3g for Me2N(CH2)sNH;2 . 159

3 6 . A Plot of pKa vs. log for Diamines. 165

XI Page

5T- Plot of a* vs. log for Diamines. l66

3 8 . Plot of pKa vs. log Kg. I67

3 9 . Plot of a* vs. log Kq for Diamines. 168

40. A Plot of pH vs. ksots Me2ÎT(CH2)2im2 . I82

41. A pH-Rate Profile for Me2N( CH2 ) 2HH2 'with Calculated I83 Contribution for Each Rate Teim.

42. A Plot of pH vs. RsqTjs for Me^(CH2)3HH2 " l84

4 3 . A pH-Rate Profile for Me2N( CH2 )sHH2 "with Calculated 185 Contribution for Each Rate Term.

44. A Plot of pH vs. ksQ-bs Me2H( CH2) . I86

4 5 . A pH-Rate Profile for Me2H( CH2 )4ÏÏH2 with Calculated I87 Contribution for Each Rate Term.

46. A Plot of pH vs. k2obs Me2lî(CH2)sNIfe- I88

4 7 . A pH-Rate Profile for Me2N( CH2 ) sNH2 with Calculated I89 Contribution for Each Rate Term by Assuming General- Acid Catalysis.

48. A pH-Rate Profile, for Me2W(CH2 )sKH2 with Calculated I90 Contribution for Each Rate Term by Assuming No General-Acid Catalysis.

4 9 . Plot of log koKg vs. pKa for Mono- and Diamines. 192

5 0 . Plot of pKa vs. log kpj^ or log kg. 193

5 1 . Plot of pKa vs. log Kgkg for Specific Acid Catalysis. 195

5 2 . Plot of pEa vs. log Kgpjg (k^^ + kp^). 197

53 . A Br^nsted Plot for the Acid-Catalyzed Hydration of 201 Isobutyraldehyde.

5 4 . A Br^nsted Plot for the Base-Catalyzed Hydration of 202 Isobutyraldéhyde.

5 5 . Plot of [EMlPt]/[EMip] vs. n. 206

xii Page

5 6 . Plot of i-.' ^ and pKaj^^ vs. n. 209

57* Plot of pKSj^^ and pKa^yg^ vs. n. 210

5 8 . Plot of time vs. In Ag/Kg (A-Ag) for Hydration of 2hh Isobutyraldéhyde.

5 9 . Chemical Shift of Protonated Trimethylamine. 246

6 0 . Chemical Shift of Protonated N,N-Dimethylethylene- 248 diamine.

6 1 . A Copy of the Photo for Run 210.801 - 210.805 • 250

62- The Observed Transmittance vs. Time Plot, Obtained 251 from the Stopped-Flow Photo.

6 3 . Plot of f(A) vs. time -where t = f( A). • 252

64. Plot of the Calculated Concentrations of Aldehyde, 255 Hydrate, Carbinolamine and Imine vs. Time.

6 5 . A Copy of the Photo for Run 210.661 - 210.664. 255

6 6 . Observed Transmittance vs. Time Plot, Obtained from 256 the Stopped-Plo-w Photo.

6 7 . Plot of f(A) vs. time where t = f(A). 257

6 8 . Plot of the Calculated Concentrations of Aldehyde, 258 Hydrate, Carbinolamine and Imine vs. Time.

6 9 . An Exaggerated Illustration of the Mixing for the 262 Stopped-Flow.

Xlll CHAPTER I

Introduction

An area of research of growing interest in recent years is bifunctional, polyfunctional, or "enzyme-like" catalysis. A bi­ functional catalyst is one that uses two groups acting simultan­ eously on different parts of the substrate. It is proposed that the efficiency of enzymatic catalysis may be mechanistically 1 rationalized as a form of bifunctional catalysis. In a study of

O'-hydrogen exchange of isobutyraldéhyde-2 -d in the presence of methylamine and méthylammonium ions the reaction was reported to 2 follow third order kinetics, eq. 1 , with the rate determining

V = kCMegCDCHO] [MeHEg] [MeH% t] (l)

step involving the attack of the methylamine base on the conjugate

acid of the H-methylimine to give an enamine, eq. 2. The rates of

Me2CDCH=HHMe + MeNHs — --f Me2C=CHNHRe + MeMgD'*’ (2)

dedeuteration of i sobutyr aide hy de -2 -d with such bifunctional species

as w-amino-n-alkanoic acids and Ui-dimethylamino-alkylamines, which

contain a primary amine capable of forming the imine intermediate

1. J. H. Wang, Science, l6l, $28 (I968 ) and a basic site capable of intramolecular dedeuteration, did not 3 indicate the occurrence of bifunctional catalysis. These results

■were rationalized in terms of the limitations imposed by the stereo- electronic requirements of the transition state. Similar work with acetonewas initiated by John C. Kaufmann and is continuing with

Michael S. Cholod. The dedeuteration of isobutyraldehyde-2-d with the polyfunctional catalyst polyethylenimine investigated by Francis

E. Rogers and Robert Notari is reported to have "enzyme-like" be- 4 havior. Jesse Lynn began to evaluate the structural requirements for this dedeuteration reaction by using more rigid bifunctional catalysts of the form Me2NCH2C=C( CHa) •

■ To increase our understanding of this dedeuteration reaction, the thermodynamics and kinetics of the formation of the R-alkylimine were investigated. The thermodynamic study was conducted by C. Y.

Yeh, who determined the equilibri'um constants for imine formation

2. (a) J. Hine, J. Mulders, J. G. Houston and J. P. I doux, J. Org. Chem., 32, 2205 ( 19^7) .

(b) J. Hine, B. C. Menon, J. H. Jensen, and J. Mulders, J. Amer. Chem. Soc., 8 8 , 35^7 (1966).

(c) J. Hine, P. C. Kokesh, K. G. Hampton and J. Mulders, J. Amer. Chem. Soc., 8 9 , 1205 (I967).

(d) J. Hine, J. G. Houston, J. H. Jensen and J. Mulders, J. Amer. Chem. Soc., 8 7 , 5050 (I965).

3 . J. Hine, B. C. Menon, J. Mulders, and R. L. Flachskam, Jr., J. Org. Chem., 3^, 4o83 (1969)*

4. J. Hine, F. E. Rogers, and R. E. Notari, J. Amer. Chem. Soc., 9 0 , 3 2 7 9 (1968). 5 from isobutyraldéhyde and various primary alhylamine's. The kinetic study v!as conducted in this investigation.

The scope of this investigation is twofold: first, to determine

structural effects on the rate of formation of N-isobutylidenealkyl- amines, secondly, to determine if the rates of formation of imines from i sobutyr aldehyde and u)-dimethylamino -alkylamine s are subject to

intramolecular catalysis. The accepted mechanism for this reaction, eq. involves a bimolecular addition to form the carbinolamine in­ termediate and the unimole cular dehydration of this intermediate.

0 HO i-ErC-H + feNE z " i-Pr-C-HHE i-Pr-C#E + HsO (5) ^ - à mu - i

Under neutral and basic conditions, the solution conditions of this

investigation, the dehydration of the carbinolamine intermediate is 6 .rate determining. The rate constants of imine formation from iso- 7 butyraldéhyde and n-propylamine, methylamine, 3 -methoxypropylamine,

2 -met hoxye thy lamine, 2 ,2 -dimethoxyethylamine and 2 ,2 ,2 -trifluoro­

ethylamine can yield a Taft relationship plot, which can test the

assumed mechanism and permit estimation of rate constants that are

5. (a) J. Hine and C. Y. Yeh, J. Amer. Chem. Soc., 89 , 2669 (I9 6 7 ).

(b) J. Hine, C. Y . Yeh,and F. C. Schmalstieg, J. Org. Chem., 35, 3^0 (1970).

6 . U. p. Jencks, Frogr. Phys. Org. Chem., 2, 63 (196*»-).

7. J. Hine, F. A. Via, J. K. Gotkis and J. C- Craig, Jr., J. Amer. Chem. Soc., in press. k 6 not or can not be experimentally determined. Jencks and coworkers have demonstrated that the dehydration of the carbinolamine is

subject to acid catalysis. Therefore, a bifunctional amine con­ taining an nucleophilic site (an unprotonated primary nitrogen)

and an acidic site (a protonated tertiary nitrogen) may intramole-

cularly catalyze imine formation. Pig. 1. The occurrence of intra­ molecular acid catalysis can be determined by comparing the

experimentally observed rate constants of imine formation from

i sobutyr aldehyde and monoprotonated-tw-dimethylamino-alkylamines (l)

to the rate constants estimated from the Taft relationship. The

equilibrium constants for carbinolamine formation were also deter­

mined. Me .Me Me Me II fast H^©'^CH2 C + H © CHo = = ^ HO I i Pr" '"H CHg ' |-Pr — C —N

H it M^ /Me Me^ ^ Me

" ^ ^"2 fo^ HOH

" x / C H 2 — e / C H g C—N C—H l-Pr ' l-Pr

Figure 1. Intramolecular Acid Catalysis. v/i CHAPTER II

Historical

A. Kinetics and Thermodynamics of Imine Formation

Imines, azomethines, anils, or Schiff bases result from nucleo­ philic addition at the carbonyl center by the lone pair electrons of a nitrogen atom bound to t-wo dissociable protons. The first step R-C-H + R ’EHs ^ = ± R-C-H-R'u + HgO (4) involves the removal of a proton from nitrogen and the addition of a proton to oxygen to form the tetrahedral addition intermediate, a carbinolamine. The second step, also involving the removal of a proton from nitrogen and the addition of a proton to oxygen, results in the formation of a water molecule and the imine. Experimental results indicates that there are at least two consecutive steps and one intermediate in this reaction mechanism.

As early as I907 Acree and Johnson explained their observed acid catalysis of acetaldoxime formation by proposing the existence of an intermediate, eq. 5- Although they were not completely lucid

+ + (CIfe)2C0 + NHa-OH -- > (CH3)2C0'N%'0H (5) as to the structure of this intermediate, the concept of the two 7 8 step process "was initiated. The follo-wing year Barrett and Lap- 9 ■worth working on the formation of acetone and acetaldehyde oximes

reported a bell-shaped pH-rate curve. A reaction thus character­

ized must necessarily proceed in t-wo steps, an addition and a de­

hydration step, in which under different experimental conditions 10 one or the other step is rate determining. Bodfoss -was the first

to provide definite experimental evidence for a carbinolamine inter­ mediate, He showed that the rate of disappearance of phenylhydrazine

is much faster than the rate of appearance of the phenylhydrazone

product ; hence the addition compound must have been formed rapidly 11 and in relatively high concentrations. In 1927 A. Olander kine-

tically demonstrated the existence of a carbinolamine intermediate. 12 Willi and Robertson were the first to undertake a kinetic

investigation of imine formation and hydrolysis. By measuring the

rates of hydrolysis of. substituted benzylideneanilines in neutral

solutions these authors concluded that the observed rate constants

contained terms for specific acid and general acid catalysis. 13 Willi postulated (from p* -1.^5) that the rate-limiting step under

8 . S. F. Acree and J. M. Johnson, Chem. Abs., 2_, 1127 (I908).

9 . F. Barrett and A- Lapworth, J. Chem. Soc., 93, 85 (1908).

10. S. Bodfoss, Z. Phys. Chem. (Leipzig), 109, 225 (1924).

11. A. Olander, Z. Phys. Chem. (Leipzig), 129, 1 (1927).

12. A. V. Willi and R. E. Robertson, Can. J. Chem., 31, 3^1 (1953) »

1 3 . A. V. Willi, Helv. Chim. Acta, 39, 1193 (1956). 8 these conditions is the attack of hydroxide on the conjugate acid of the imine and the acid catalyzed addition of -water to the imine.

By application of the principle of microscopic reversibility, one can infer under neutral conditions that carbinolamine formation is 14 not rate determining. Jencks reports that near neutral pH the reaction of excess hydroxylamine, methoxyamine, hydrazine, and semicarbazide with a number of carbonyl compounds results in a rapid decrease in the ultraviolet absorption of the carbonyl group that is not accompanied by a rapid appearance of the absorption peak of the product. The magnitude of this rapid initial decrease in the carbonyl concentration was proportional to the concentration of nitrogen base and used in the calculation of equilibrium con­ stants for carbinolamine formation, eq. 6 . Jencks determined that

q OH I I fast I , R-C-H + HsNR* ; ^ = = ± R-C-ip' (6 ) H H the formation of the oxime of pyruvate ion under neutral conditions with excess hydroxylamine is kinetically independent of the hy­ droxylamine concentration because the pyruvate is completely con­ verted to the carbinolamine addition intermediate: With relati-vely low concentrations of hydroxylamine only a fraction of the pyruvate

is converted to the intermediate and the reaction rate is dependent on the concentration of each of the reactants. Thus, the dehydration of the. carbinolamine intermediate is rate-determining near neutral pH.

l4. W. p. Jencks, J. Aner. Chem. Soc., 8l, (1959) 9 In acidic solutions the nitrogen base becomes protonated and this decrease in the concentration of free nucleophilic reactant causes a decrease in the concentration of the intermediate carbin­ olamine; -whereas, the dehydration step is acid-catalyzed. Thus, these t-wo effects should produce an observed rate that is pH inde­ pendent. Ho-wever, it is often found that the observed rate de­ creases -with decreasing pH and therefore, must reflect a change in 15 the rate-determining step. Additional information confirming the two-step process for imine formation includes subs-tituent-effect studies and relative sensitivities to acid catalysis on the two sides of the pH-rate maxim-um. 16 Cordes and Jencks ( see Figure 2) found that the rates of hydrol­ ysis of p and m-substituted benzylidene-l,l-dimethylethylamines are independent of pH above pH 9 and increase with electron-donating power of the polar substituent (p = -0.2l). These results suggest that the attack of hydroxide ion on the protonated Schiff base is rate- determining. Under more acidic conditions in which an appreciable fraction of the Schiff base exists as the conjugate acid, the hydroly­ sis rates increase with decreasing pH for Schiff bases possessing an electron withdrawing substituent and decrease with decreasing pH for

Schiff bases containing an electron-donating substituent (p = l.Jl) •

The changes in the rate with pH can be correlated to the conversion of

15- ¥. p. Jencks, J. Amer. Chem. Soc., 8 o, 4^81 (1958). l6. E. H. Cordes and W. P. Jencks, J. Amer. Chem. Soc., 8 5 , 2Sk3 (1963). 10

p-HO.

m-Br

p-Cl

0.1

p-OCH.

0.01

2 4 6 8 10 12 pH

Figure 2. Logarithm of the First-Order Rate Constants for Hydrolysis of Substituted Benzylidene-1,1- dimethylethylamine. 11 the Schiff bases to their conjugate acids, indicating that the'pre­ dominant reaction "under these conditions is the attack of water on the protonated Schiff base. The application of the principle of microscopic reversibility reveals that the dehydration of carbin­ olamine proceeds primarily by an acid catalyzed mechanism under these conditions. Below pH 4, the hydrolysis rates for all of the Schiff bases decrease and eventually become linear ■with re­ spect to hydroxide ion concentration and. can be correlated by ct'*’ substituent constants with a of 2.17. This behavior indicates a transition in the rate-deteimining step to the nucleophilic addition of the amine to the substituted benzaldehyde. 17 Cordes and Jencks reported that the formation of N-p- chlorobenzylideneaniline is subject to general acid catalysis with the catalytic constant at pH 2.5 for acetic acid about tenfold greater than at pH 5-I- These authors interpret this differing, susceptibility of Schiff base formation to general acid catalysis on the two sides of the pH- rate maximum as evidence for a change in rate-determining step.

Catalysis of Imine Formation

Rates of formation of imines are susceptible to general acid- base catalysis. The relative significance of these catalyzed pathways is dependent upon the nucleophilicity of the amine base.

IT* E. H* Cordes and W. P. Jencks, J. Amer. Chem. Soc., 84, 852 (1962). 12

The rate of addition of a weakly basic amine to a carbonyl group

can be accelerated by general acid and/or general base catalysis, whereas the rate of addition of a rather basic amine, alkylamine,

is not effected by the presence of these catalysts. The dehydra­ tion of a carbinolamine near neutral pH proceeds predominately by an acid-catalyzed pathway. However, at alkaline pH the rate of

dehydration of carbinolamines formed from weakly basic amines, 6 oximes, semicarbazines, exhibits general base catalysis. Since there is incorporated into the scope of this research program an

effort to evaluate the effectiveness of an internal general acid

catalyzed path for dehydration of a carbinolamine, the mechanistic

aspects of the catalysis of imine formation will be considered.

The addition reaction: The acid catalysis of the attack of

semicarbazide on ^.-nitrobenzaldehyde could occur by proton donation

to the carbonyl group, eq.. 7 , or by general base catalysis of the

attack of semicarbazide on the conjugate acid of the aldehyde, eq.

8 . Incorporation of the ionization constants for the aldehyde.

HgldH ^products (7) Vy = kylEmslC )C=o][A-H]

H ^ . fast ^ V n I . C=0 + H+ HO— C < --- products + I (8) Vs = kslENHsDC ^C=OH][a ‘] ^

^ M ’ the catalyst, into the expression for Vg illus­

trates that mechanisms 7 and 8 are kinetically indistinguishable. 13 i.e., the observed rates of reaction can be defined in terms of a constant, the aldehyde, amine and catalysis concentrations, eq. 18 10 and 7. Jencks and Cordes determined that catalysis of car-

[C=Q][H+] ^ [A-][H+] . . %A - [c=OH+] %-H " A H

Vs = ksCRNHs] Ca h ] k ^AA. A-fi k* %A.H nr. V C m h 2 ][c=o][ah] 'AA binolamine formation proceeds via mechanism 7* These researchers determined the rate of product formation, Vyorg, values for and and then calculated a value for kg of eq. 10. This calcu­ lated rate constant, kg =2.7 x 1 0 ^ M~^sec"^, is larger than the rate constant for diffusion controlled reactions in water, 1.4 x

10^^ M ^sec and thus demonstrates that mechanism 8 cannot account for the observed rate of seraicarbazone formation. Other 20 workers also believe that general acid catalysis of addition of a weakly basic amine to a carbonyl proceeds via mechanism of eq. 7*

General base catalysis for the addition of a primary amine across

10. E. H. Cordes and W. P. Jencks, J. Amer. Chem. Soc., 84, 4^19 (1962).

1 9 • M . Eigen and L . DeMaeyer, Z. Electrochem., 59, 986 ( 1955) •

20. C. G. Swain and E. R. Thornton, J. Amer’. Chem. Soc., 8 3 , 3884 (1961). 14 6 a carbonyl bond is not observed.

The dehydration of carbinolamine at intermediate pH values 21-23 is subject to general acid-general base catalysis. Mechan­ istically the general acid catalysis involves the protonation of the carbinolamine, eq. 11, as the rate-limiting step. General base catalysis involves deprotonation of a protonated carbinol­ amine, eq. 1 2 , as the rate-limiting step.

\ I _ + fast . \ , A + H-0 + y = m ^ AH + HsO C=H\ (ll) R H R R

I fast _ 1 ^ , HA + HOCNH H-0— C-H-H A ;^I± HOH + C=N\ + AH (12) R H ' i ' E 24 Cordes and Jencks observed general base catalysis for the hydrol­ ysis of benzylidene-t-butylamines under condition in which the Schiff base is completely protonated and therefore, must react by mechanism 25 11. Using a molecular model approach, Reimann and Jencks found that the Br^z^nsted slope (0.7T) for general acid catalysis of oxime

21. W. p. Jencks and E. H. Cordes, J. Amer. Chem. Soc., 84, 8$2 (1962).

22. R. L. Reeves, J. Amer. Chem. Soc., 84, 3352 (1962).

2 5 . A. V. Willi and R. E. Robertson, Can. J. Chem., 31, 36I (1955) •

24. E. H. Cordes and W. P. Jencks, J. Amer. Chem. Soc., 8 5 , 2843 (1963).

2 5 . J. E. Reimann and W. P. Jencks, J. Amer. Chem. Soc., 8 8 , 3973 (1966). - 15 and nitrone formation from p-chlorobenzaldéhyde -were identical.

The fact that the carbinolamine formed from H-methylhydroxylamine contains no nitrogen bound proton and that similar catalysis is seen for both compounds rules out mechanism 12. In alkaline pH the dehydration of a carbinolamine containing electron-•withdrawing 26-28 substituents on the nitrogen atom is subject to base catalysis.

Ho-wever, it is not kno’wn -whether these reactions proceed by gen­ eral or specific base catalysis.

B. Intramolecular and Bifunctional Catalysis.

A significant aspect of molecular biochemistry currently in­ volves attempts to remove the "magic" from enzymatic catalysis of natural processes and replace this with a mechanistic elucidation defined in terms of basic kinetic and thermodynamic principles.

The various ramifications of these findings are of pharmacological significance. Transamination reactions, those which "in vitro" convert keto acids to amino acids, are of biological importance.

Evidence suggesting the involvement of imine and carbinolamine 29 >30 intermediates in several enzymatic transamination reactions

26. B. M. Anderson and W. P. Jencks, J. Amer. Chem. Soc., 82, 1TT3 (i960).

2 7 . B. L. Hastening, et al., Z. Electrochem., 60 , 130 (1956).

2 8 . D. H. R. Barton, R. E. O'Brien, and S. Sternhell, J. Chem. Soc., 1962 (470).

2 9 . E. E. Snell and W. P. Jenkins, J. Cell Comp. Physiol., 54, 161 (1959). 16 has focused attention on the formation and hydrolysis of these

compounds. Transamination of pyridoxal-5-phosphate bound to an enzyme moiety at pH 8.10 occurs ten million times faster than does 31 that of pyridoxal, Fig. 3- In the preceding section we have seen that the dehydration of carbinolamine is rate determining at pH

8.10 and subject to general acid catalysis. Since the enzyme con- 31 tains a lysine linkage in the vicinity of the reaction center, application of kinetic principles indicates that the enzyme may function via intramolecular general acid catalysis. Intramolecular reactions usually occur more rapidly than the analogous intermole-

cular reactions because the catalytic center is located in the same molecule and will spend more of its time in a position conducive to

catalytic interaction than if it were a separate molecule. In this example the inter- and intramolecular reactions that are being com­ pared are not strictly analogous, for such a comparison is compli­

cated by the fact that the intermole cular reaction will usually be bimolecular with a rate constant kg, which may be expressed in units of M“^sec“^, while the intramolecular reaction will be uni- molecular with a rate constant ki, which may be expressed in units

of sec“^. The ratio ki/kg has the units of molarity and is termed the "effective concentration." This would be the hypothetical

50. D. S. Metzler, M. Ikawa, and E. E. Snell, J. Amer. Chem. Soc., 76, 648 (1934).

31- E. H. Fisher, A. B. Kent, E. R. Snyder, and E. G. Krebs, J. Amer. Chem. Soc., 80, 2906 (1958). ,-Enzyme 'Enzyme

PO3 ! x'^°3NH2

? I k' Î hc^ ' ^ ' r”" "Y DH + H2N-C-COOH — »

tN© x CH% " , ■ CHg H ^ H

k ' = 1 0 ^ Jl mole"' sec"'

k" = 1 / mole"' sec"' .

i-Pr CH-COOH

HOHaCUyOH ^ h^n -C-COOH

S® CHj \ V y C H ;

Pigui’e 3* Enzymic Catalysis of Transamination. 18 concentration at -which the catalyst in the "bimolecular reaction

■would have to be present in order to obtain a pseudo-first-order rate constant equivalent to the true rate constant for the intra­ molecular reaction. Thus the transaminase enzyme has an "effective concentra-fcion" of lO’^M. Mechanistically this "effective concentra­ tion" probably results from a summation of those factors attributed 32 to enzymatic behavior. Of course these concentrations may not be obtainable, but the me-thod serves to show the very large enhance­ ments obtained and is equivalent to the comparison of values.

The scope of this research includes an evaluation of the feasability of internal catalysis and intermediate stabilization as methods of enzymatic catalysis. The broad catalytic principles that are currently used to describe the mechanistic behavior of enzymes are outlined under the classifications of internal general acid catalysis, internal general base catalysis, internal nucleo- philic catalysis, internal metal ion catalysis, internal hydrogen bonding, and bifunctional catalysis.

Internal General Acid Catalysis

In one of the first accounts of the catalysis of ester hydrol- 33 ysis by other than lyate species, Garrett proposed an intramole-

cula-^ general acid assisted nucleophilic attack on the pyrrolidyl-

acetylsalicylate (ll) ester by an acetate ion.

32. W. p. Jencks, "Catalysis in Chemistry and Enzymology," McGraw- Hill Book Co., Ne-w York, 1969* 19

C-OCHs

0— q-CHs

AcO -/I...

II

34 Hansen illustrated the catalytic effects of an internal general acid in the hydrolysis of thioesters.

P. ^ e l II H+ , CHbCHaC-SCHsCHsNMea 2kO

R + CHbCHaC-SCHsCHsNMes 1

The important difference "between these two esters is that one con­ tains a quaternary ammonium ion while the other contains a tertiary ammonium ion that apparently functions as a general acid catalyst

R- u s-

III

53* E. R. Garrett, J. Amer. Chem. Soc., 79, 5206 (1937) «

5^- B. Hansen, Acta Chem. Scand., 12, 324 (1958). 20

T. G. Bruice, -who has Been most prolific in his investiga­ tions of "enzyme-like" catalysis, reports that internal general acid catalysis significantly increases the rate of transamination 35 of 3 -hydroxypyridine-4-aldehyde by glutamic acid and by BL- 3sa,b alanine (IV). Bruice also reports the existence of internal general acid catalysis in the cyclization of 3 -hydroxypyridine-4- , , 37 aldehyde with histamine (V) .

CHs Çt -H :B

Internal General Ease Catalysis

Only a few examples of intramolecular general base catalysis 38 have been mechanistically substantiated. In I966 Menger reported

55* J. W. Thanassi, A. R. Butler, and T. C. Bruice, Biochemistry, k, 1463 (1965).

3 6 . a) D. S. Anld and T. C. Bruice, J. Amer. Chem. Soc., 8 9 , 2090 (1967). b) ibid., 89, 2098 (196771

37" T. C. Bruice and A. Lombardo, J. Amer. Chem. Soc., 91, 5009 (1969).

3 8 . F. M. Menger, J. Amer. Chem. Soc., 88, 3081 (1966). 21 that the ami dinolysis of p-nitrophenyl acetate with benzamide was assisted by internal general base catalysis. However, Anderson 39 and coworkers demonstrated that Menger could not distinguish between internal and bimolecular general base catalysis. The 40 hydrolysis of aspirin and its derivatives has been reviewed and is believed to be subject to both intramolecular nucleophilic 41 catalysis and, as Fersht and Kirby demonstrated for acetylsali- cylic acid (Vl), is subject to internal general base catalysis.

Bruice demonstrated that the aminolysis of 8 -acetoxyquinoline is

— C — C H j3

C—_ n 0 - I

subject to internal general base catalysis for a Bryinsted plot reveals two lines, one for proton bearing nucleophiles, primary and secondary amines, and one for non-proton bearing nucleophiles, 42 tertiary amines. Bruice has also demonstrated that the

39' H. Anderson, C. Su, and J. W. Watson, J. Amer. Chem. Soc., 91, 483 (1969).

40. M. L. Bender, Chem. Reviews, 60, 55 (I960).

41. A. R. Fersht and A. J. Kirby, J. Amer. Chem. Soc., 8 9 , 4855 (1967).

42. T. C. Bruice and S. M. Felton, J. Amer. Chem. Soc., 9 1 , 2J99 (1969). 22 hydrolysis of 8 -acetoxyquinoline (VIl) and 4-(2'-acetosyphenyl)- imidazole (VIIl) is assisted ty intramolecular general "base catalysis.

H-OH

I ^ 1 0=C (— rURs

VII VIII

Internal Encleophilic Catalysis

Of all forms of intramolecular catalysis the nucleophilic 40 attack has the most abundant claims in the literature. In 195^ 43 Sondheimer and Holley reported that carbobenzossy-L-glutamine hydrolyzes via the internal nucleophilic attack of the amide group, eq. 15- Bender employed labeling techniques to verify Ï\ CHs-CHHg CH2 -C\ .KH Cdz-NH-CH-COCCHa CdzM-CH— (13) ^ 0 internal nucleophilic catalysis for hydrolysis of phthalamic 44 45 acid, eq. 1 4 , and methyl hydrogen phthalate.

4$. E. Sondheimer and R. W. Holly, J. Amer. Chem. Soc., j6, 2467 (1954).

44. M. L. Bender, y. L. Cho-w, and F. Chloupek, J. Amer. Chem. Soc.,

8 0 , 5 3 8 0 (1958). ' 23

(14)

CO^H

46)47 Bruice and co-workers postulate a similar internal attack by an imidazolyl group for the hydrolysis of 2 -( 4 * -imidazolyl)phenyl acetate, eq. 15.

C-CHb

(15)

rds, + CHsCOOH

Several investigators report that the hydrolysis of a mono-

anion of esters of dlcarboxyl!c acids proceeds through a cyclic 48,49 anhydride intermediate. A somewhat different type of

4 5 . M. L. Bender, F. Chloupek, and M. C. Neveu, J. Amer. Chem. Soc., 80, 5584 (1958).

46. G. L- Schmir and T. C. Bruice, J. Amer. Chem. Soc., 80, II73 (1958).

h j . U. K. Pandit and T. C- Bruice, J. Amer. Chem. Soc., 82, 3386 (i960).

48. E. Gaetjens and H. Morawetz, J. Amer. Chem. Soc., 82, 5328 ( i960) . — 2k 50 intromolecular catalysis is reported, by M. S . Ne-wman for the hydrolysis of methyl-2-benzoyl-6-methylbenzoate (IX). Here the bimolecular step, eq. l6 , occurs before the internal one, eq. 17.

COOCHs

-OH (16)

IX

11

HO c.

CHsC

^9* (a) T* C. Bruice and U. K. Pandit, J. Aaer. Chem. Soc., 82, 5 8 5 8 (1960).

(b) T. C. Bruice and U. K. Pandit, Proc. Natl. Acad. Sci., U.S., 46, 402 (i960).

5 0 . M. S. Ne-wman and S. Hishida, J. Amer. Chem. Soc., 84, 3582 (1962). 25

Internal Metal Ion Catalysis

Although it is not possible for metal ion catalysis to occur in this investigation, the mechanism is of relative biochemical significance. Speculation concerning the role of a metal ion in 51 enzymic reactions is reviewed. Generally, catalysis occurs when the metal can chelate with the reaction site of the substrate, thus rendering this substrate more susceptible to nucleophilic attack. This mechanism could also be classified as internal Lewis acid catalysis for the metal accepts a pair of electrons frcm the substrate. Metal ion catalysis has been frequently reported for 52 53 the hydrolysis of amino acid esters and monoester of diacids.

Internal Hydrogen Bonding

This mechanism is distinguished from internal general acid catalysis in that the catalytic proton is not divorced from its initial site. This type of catalysis can occur in a molecule con­ taining a proton bearing, basic group in the vicinity of an appro- 54 priate reaction site. Pascual and coworkers were among the first

51' F. H. Westheimer, Trans. U. Y. Acad. Sci., l8 , 15 (1955) «

52. (a) H. Kroll, J. Amer. Chem. Soc., 74, 2054 (1952).

(b) M. L. Bender and B. W. Turnquest, ibid., 79, .I889 (1957) •

( c) W. L. Koltun, M. Fried, and P. R. R. Gurd, ibid., 8 2 , 255 (i960).

55* J' H. Hoppe' and J. E. Prue, J. Chem. Soc., 1775 ( 1957) «

5 4 . J. Pascual, J. Sistare, and A. Regas, J. Chem. Soc., 1943 (1949). 26 to suggest that hydrogen bonding stabilization of a transition state accounts for an increase in reactivity. They observed that the aminolysis with p-toluidine of cis-2 -hydr oxycyc lohexane car - boxylic acid (x) proceeds faster than the trans compound. Simi-

OH

X

55 larly, Henbest and Lovell found, that with coprostane 3-acetoxy-

5-hydroxy, steroids the ester group that was cis to the hydroxyl group underwent solvolysis faster than the ester group which was trans. Acceleration of the base-catalyzed methanolysis of cis- cyclohexane-l,3-diol monoacetates was attributed to intramolecular 56 hydrogen bonding. However, mechanistically there are uncertain­ ties. Several investigators consider that hydrogen bonding to 55 either oxygen atom, eq. lo, would catalyze the reaction; others

Product (l8)

CHa

55- H. B. Henbest and B. J. Lovell, J. Chem. Soc., 1965 (1957)*

56. S. M. Kupchan and C. R. Haraganan, J. Amer. Chem. Soc., 8l 1913(1959)- 27 believe bonding only to the carbonyl oxygen of the ester ■would be 57 58 effective. Bruice and Fife postulate from results of alkaline hydrolysis of a number of cyclopentane and norbornane acetates and diol monoacetates that the possibility exists that the internal solvation of the transition state by neighboring hydroxyl groups may be due to more subtle factors, like microscopic solvent changes. The proximity of the hydroxyl group to the ester group may change the medium surrounding the ester group. Theoretical interpretation of the nature of a microscopic solvent change is not reliable.

Bifunctional Catalysis

S'wain and Brown were among the first to report the significance of bifunctional catalysis. They found the mutarotation of 2 , 3 6 - tetramethyl-D-glucose would be 7000 times faster in the presence of

10""^ M 2-hydroxypyridine, eq. 19, than with an equivalent

or / \ (19)

57- H. G. Zachan and W. Karan, Ber., 93, 1830 (1960).

5 8 . T. C. Bruice and T. H. Fife, J. Amer. Chem. Soc., 84, 1973 (1962) . 28 59 concentrations of phenol and pyridine. Catalysts that have tvo or more reactive centers and give large positive deviation from a

Br^nsted relationship are usually said to be bifunctional catalysts.

The most frequently reported catalysts to illustrate this behavior 60 are carboxylic acids and phosphates. The work of Sander and Jencks illustrates this mechanism of catalysis. Carboxylic acids (Xl) and hydrogen phosphates (XIl) give large positive deviations and an a =

0 for the Br^nsted relationship of the general acid-base catalyzed addition of peroxide to aldehydes and dimethylphosphate does not.

Similarly bifunctional catalysis has been reported for phosphoric 61 acid anions in the hydrolysis of iminolactones, the reaction of 62 urea with formaldehyde, the hydrolysis of 1 ,3 -diphenyl-2 -imidazo-

of\ ,\ ^ O H R ’-CH TV ^- DR'-CH I _r«iT 0-' ^OH o f ^ O H XI XII

39- C. G. Swain and J. F. Brown, J. Amer. Chem. Soc., T4, 2558 (1952).

60. E- G. Sander and W. P. Jencks, J. Amer. Chem. Soc., 90s ^377 (1968). ^

61. B- A. Cunningham and G. L. Schmir, J. Amer. Chem. Soc., 88, 551 (1966).

62. B- Glutz, Angew. Chem. Intern. Ed. Engl., 4, 44o (19^5). 29 63 64 linium chloride, the cyclization of glutamic acid esters, and 65 the conversion of glutamine to pyrrolidonecarboxylic acid, and for carboxylic acids or carbonates in methoxyaminolysis of phenyl- 66 67 acetate, hydrolysis of 2 acetimidate esters, ethanolysis of 68 sec-butyl borate, and hydrolysis of 2,2,2 -t r if luor o -N -met hyl- 69 acetanilide. For other miscellaneous examples of bifunctional 70 catalysis see Rony's article on polyfunctional catalysis.

This review of the mechanistic principles associated with enzymic reactions is designed to familiarize the reader with mech­ anisms similar to those studied in this investigation; for compre- 71 32 hensive reviews see Bruice and Behkovic and Jencks. Some examples of intramolecular catalysis are not included because of

65. D. R. Robinson and W. P. Jencks, J. Amer. Chem. Soc., 89 , 7 0 8 8 (1967). ------

64. A. J. Hubert, Helv. Chim. Acta, 1429 (1 9 6 3) .

65. A. Meister, J. Biol. Chem., 210, 17 (1954).

66. L. do Amaral, K. Koehler, D. Bartenback, T. Fletcher, and E. H. Cordes, J. Amer. Chem. Soc., 8 9 , 5537 (19 67).

6 7. R. K. Chaturvedi and G. L. Schmir, J. Amer. Chem. Soc., 90, 4415 (1968).

68. C. L. Denson and T. I. Crowell, J. Amer. Chem. Soc., 79, 5^56 (1 9 5 7). : '

69. • R. L. Schowen and G. W. Zworick, J. Amer. Chem. Soc., 8 8 , 1 2 2 5 (1966). ------

7 0 . p. R. Rony, J. Amer. Chem. Soc., 91, 6090 (19^).

7 1. T. C. Bruice and S. J. Benkovic, Bioorganic Mechanisms, W. A. Benjamin Co., Inc., Rew York, 1 9 6^ 50 72 mechanistic uncertainties. Bruice and covjorkers report that 3- hydroxylpyrididine-ij--aldehyde forms imines 6o times faster than does pyridine-4-aldehyde. However, it is not clear whether the hydroxy group behaves as a general acid XIII, a general base XIV, causes ring transmitted expulsion XV, or hydrogen bonds to the leaving group.

RHN-C

XIII XIV XV

C. Hydration of Isobutyraldéhyde

In the pursuit of our interests, that of studying the rate of formation of imines from the reaction of various primary alkyl­ amine s with isobutyraldéhyde, it became evident that the rate of formation of the aldehyde by dehydration of its hydrate may be­ come comparable to the rate of formation of imine. This signifi­ cantly affects the kinetics of imine formation and requires experimental measurements of the concentration of imine as well

72. (a) D. S. Auld and T. C. Bruice, J. Amer. Chem. Soc., 89 2085 (1967)-

(b) T. C. French, D. 8 . Auld, and T. C. Bruice, Biochemistry, 77 (1965) • 51 as that of aldehyde at various times. In an effort to avoid this additional kinetic complexity an investigation to obtain a complete rate expression for the hydration of isobutyraldéhyde -was under­ taken. With this rate data it might be possible to avoid pH ranges in ■which the rates of formation of imine and of dehydration of aldehyde are comparable.

The existence of a hydrated form of aldehyde in aqueous solu- 73 74 tions "Was first suggested by Ramsay and Young (l88b) and Perkin

(1887) • They found upon mixing acetaldehyde and water an immediate but small absorption of heat which was followed by a much larger evolution of heat lasting for several minutes. By quantitative 75 evaluation of this heat of mixing. Brown and Pickering (lo97) found that the degree of hydration varies with temperature and dilution, and that a considerable fraction of aldehyde remains 76 unhydrated. Herbert and Lander (1958) demonstrated by means of the isotope 0 ^® that there is rapid exchange of oxygen between water and acetaldehyde, which was attributed to the reversible formation of the hydrate. Bell and coworkers have employed

75' W. Ramsay and S. Young, Phil. Trans., 177, 71 (1886).

7 4 . W. H. Perkin, J. Chem. Soc., 51, 808 (I887).

75* H. T. Brown and P. S. U. Pickering, J. Chem. Soc., 71, 774 (1894). ------

76. J. B. M. Herbert and I. Lander, Trans. Faraday Soc., 34, 1219 (1958). 52 77)78 79 SO dilatometry and thermal and spectrophotcmietric methods

to determine the rate of hydration of acetaldehyde. Bell reported

that the addition of -water to the carbonyl group follows first

order kinetics in aqueous solutions, and is subject to general acid- 79 81 base catalysis. Hine and Houston used nmr to study the specific

acid catalyzed rate of hydration (35*^) of isobutyraldéhyde. At the

time of the initiation of this study, the complete expression in­

cluding the general and specific base catalyzed rate constant as

■well as the general acid catalyzed rate constant for the hydration

of this aldehyde -was unkno-wn to us. After completion of our inves- 82 tigation Pocker and Dickerson published the complete rate expres­

sion for the hydration-dehydration of isobutyraldéhyde at 0° C in

aqueous solution. Nevertheless, to fulfill the purpose of this

present investigation the rate data for hydration-dehydration of

isobutyraldéhyde at 35° "was required.

77* R* p. Bell and B. d. B. Darwent, Trans. Faraday Soc., 46, 54 (1950).

7 8 . R. P. Bell and J. C. Clunie, Proc. Roy. Soc. (London), A212, 55 (1952).

79" R" P" Bell, M. H. Rand, and K. M. A. Wynne-Jones, Trans. Faraday Soc., 52, 1100 (1956).

8 0 . R. p. Bell and p. G. Evans, Proc. Roy. Soc. (London), 2 9 7 (1 9 6 6 ). ------

8 1 . J. Hine and j. G. Houston, J. Qrg. Chem., 3 0 , 1328 (1965).

8 2 . y. C. Pocker and D. G. Dickerson, J. Phys. Chem., 75 , 4005 (1969). 53

D. Determination of Micro-pKa Values for the Diamines.

Complete evaluation of observed kinetic and equilibrium con­ stants for imine formation for the unsymmetrical dime thy lalkyl- diamines requires structural knowledge as to the time averaged location of the proton on these monoprotonated diamines. Dis­ sociation constants for basic amines in aqueous solutions do not afford a reliable estimation of the structure of a monoprotonated diamine. Evaluation of existing data for model compounds of E,N- dimethylethylenediamine results in an unacceptable degree of un­ certainty. This is shown by comparing three methods of estimating the relative basicities of the two nitrogen atoms of N, N-dimethyl- ethylenediamine.

First: Considering the pKa of the primary nitrogen of the 83 diamine to be that of ethylamine, 10.33 at 35° and the pKa of the tertiary nitrogen to be that of N,IT-dimethylethylamine, 9.80 83 at 35°, then the proton on monoprotonated E^N-dimethylethylene- diamine would be located 22.7% of the time on the tertiary nitro­ gen and 7 7 .35^ of the time on the primary nitrogen.

Second: In order to account for inductive affects the electro­ negative nitrogen atoms may exert on each other in this diamine,

consider the pKa of the primary nitrogen to be that of ethylene- 83 diamine, 9*307 at 35° and the pKa of the tertiary nitrogen to ^ S3 be that of N,]lT,N*,W’-tetramethylethylenedlamine, 8.50 at 35 ,

8 3 . D. D. Perrin, Dissociation Constants of Organic Basis in Aqueous Solutions, Butterworths, London, 1965» St- then the proton on monoprotonated W,N-dimethylethylenediamine -would be located 15-5^ of the time on the tertiary nitrogen and 86.$^ of the time on the primary nitrogen.

Third: In order to evaluate the inductive effects of the tertiary amine "beta" to a primary amine and the primary amine

"beta" to a tertiary amine, consider the pKa of the primary ni­ trogen to be that of ethylamine minus the difference in pKa between

B,IT-dimethylethylamine and ,IJ’-tetramethylethylenediamine, 83 10.55 - 1-50 = 9-05 at 55°, and the pKa of the tertiary nitrogen to be that of dimethylethylamine minus the difference in pKa be- 83 tween ethylamine and ethylenediamine, 9*80 - 1.02 = 8.78 at 55 , then the proton on monoprotonated N ,N-dimethylethylenediamine would be located 5 ^ of the time on the tertiary nitrogen and 64^ of the time on the primary nitrogen.

Table 1

Estimated Structure of Monoprotonated E -Dime t hyle thyle ne diamine

assumed pKa Location of Proton l°nitrogen 3°nitrogen l°nitrogen 3°nitrogen %

First 10.55 9.80 77.5 22.7

Second 9-507 8.50 86.5 15.2

Third 9.05 8.78 6 k 56 55

It appears that a reasonably reliable determination of the

structure of the monoprotonated unsymmetrical diamines can be obtained by observing the changes in the chemical shifts of the

N-methyl protons 'with changing acidity of the aqueous amine solu- 84 tion. In 1957 Grun-wald and coworkers reported that for methyl, dimethyl, and trimethylamine there is a linear relationship be­ tween the fraction of protonated species and the chemical shift. 85 According to accepted concepts of nuclear magnetic resonance the experimentally observed chemical shift for the N-methyl pro­ tons of an unsymmetrical dimethylalkylamine is a linear function of the fraction of tertiary nitrogen and fraction of primary nitrogen that is protonated. To determine the effects of pro­ tonation at these sites, model compounds are used.

84. E. Grunwald, A. Loewenstein, and S. Meiboom, J. Chem. Phys., 27, 641 (1957).

8 5 . L. M. Jackman, Applications of Nuclear Magnetic Resonance Spectroscopy in Organic Chemistry, Pergamon Press, New York, (1963) • CHAPTER III

EXPERIMEiraAL a ) Chemicals

Aminoa cat aldehyde Dimethylacetal ( 2,2-Dimethoxyethylaiaine), ob­

tained from Mat he son Coleman and Bell, was distilled through

a spinning band distillation column under reduced pressure:

bp 3^° (58 mm) . Analysis by glpc on a Carbowax 2CM column

at 90° indicated the amine was > 99^ pure.

H,N-Dimethyl-1,4-butanediamine, synthesized by Dr. James H. Jensen

from the nitrile, was distilled through a vigreux column un­

der reduced pressure: bp 43-50° (12 mm). Analysis by glpc

on a Carbowax 2CM column temperature programmed from 55 to

170° at a rate of 5°/min and on a Apiezon L column similarly

temperature programmed indicated the amine was > 95^ pure.

See appendix for ir data.

N,N-Dimethylethylenediamine, obtained from Ames Laboratories, Inc.,

was distilled under nitrogen through a vigreux column at at- 86 mospheric pressure: bp 104-108°; literature value 108°

86. A. Lespagnal, E. Cuingnet, and M. Debaert, Bull. Soc. Chim. France, 2, 383 (I960).

3 6 5T

Analysis by glpc on a Carbo'wax 2QM column at 100° indicated

the amine was > 99^ pure. See appendix for ir data.

N,N-Dimethyl-1,5-pentanediamine, synthesized by Dr. James H- Jensen

from the nitrile, was distilled through a vigreux column under

reduced pressure: bp 104-106° (66 mm). Analysis by glpc on a

Carbowax 2(M column temperature programmed from 55 to 170° at

a rate of 5°/inin and on a Apiezon L column similarly pro­

grammed indicated the amine was >99^ pure. See appendix for

ir data.

N,N-Dimethyl-1,5-propanediamine, obtained from Aldrich Chemical Co.,

was distilled under nitrogen through a column packed with tan­

talum heligrid: bp 129-151° at atmospheric pressure; litera­

ture 150-155 . See appendix for ir data.

Isobutyraldéhyde, obtained from Aldrich Chemical Co., was distilled

under nitrogen through a vigreux column not more than an hour

previous to usage: bp 65°.

Methanol was obtained from J. T. Baker Chemical Co.

2-Methoxyethylamine, obtained from Aldrich Chemical Co., was dis­

tilled under nitrogen through a vigreux column at atmospheric

8 7 . S. Kawakara and H. Zawakami, Yakugaku Zasski, 8 1 , 1^9 (I961). Chem. Abs. 62: 1529^a. 38

pressiare: bp 91°. Analysis by glpc on a Carbowax 2QM colnnm

at 80° indicated the amine -was > 99^ pure.

3-Methoxypropylamine, obtained from Aldrich Chemical Co., was dis­

tilled under nitrogen through a vigreux column at atmospheric

pressure: bp Il4-ll6°. Analysis by glpc on a Carbowax 2CM

column at 90° indicated the amine was > 99^ pure.

Methy lamine was obtained from Ma the son Company.

Ïï-Methylmorpholine, obtained from Eastman Organic. Chemicals Co.,

was distilled under nitrogen through a vigreux column at at- 88 mospheric pressure: bp 112°, literature ll6°.

Perchloric Acid solutions were obtained from and standardized by

the Q.S.U. Lab Stores Department.

N-Propylamine was obtained from Eastman Organic Chemical Co.

Analysis by glpc on a Carbowax 2CM column at 6o° indicated

the amine was > 99^ pure.

Sodium Chloride was obtained from J. T. Baker Chemical Co.

Tetrahydrofuran, obtained from J. T. Baker Chemical Co., was dried

over sodium hydroxide for 48 hr., distilled over lithium

aluminum hydroxide and stored under nitrogen.

88. D. D. Reynolds and W. 0. Kenyon, J. Amer. Chem. 80c., 72, 1597 (1950). 39

,F'-Tetramethylethylenediamine, obtained from the Aldrich

Chemical Co., vas distilled under nitrogen at atmospheric

pressure: bp 119-121°. Analysis by glpc on a Carbowax 2CM

column at 100° indicated the amine was > 99?^ pure.

N,N,Tf' -Tetramethyl-1,3-propanediamine was obtained from Ames

Laboratories, Inc. Analysis by glpc on a Carbowax 2CM

column at 100° indicated, the amine was > 99^ pure.

Triethylamine, obtained from Eastman Organic Chemicals Co., was

distilled under nitrogen at atmospheric pressure: bp 87-88°; 89 literature 89.4°. Analysis by glpc on a Carbowax 2CM column

at 70° indicated the amine was > 99Ç& pure.

2,2,2-Trifluoroethylamine, obtained from Peninsular Chemicals Co.,

was distilled through a spinning band distillation column

under nitrogen: bp 36°. Analysis by glpc on a Carbowax 2CM

column at 70° indicated the amine was ~ 9^ pure.

Trimethylamine was obtained from Matheson Scientific Co.

Water was obtained from O.S.U. Lab Stores.

8 9 . A. W. von Hofmann, Ann., 73, 91 (I850) 4o

B) Instrumental

Infrared Measurements

All infrared ( ir) spectra were recorded with a Perkin-Elmer

Grating Infrared Spectrophotometer, Model 337, and calibrated with polystyrene.

Gas-Liquid Chromatograph

The purity of amines used in this study was measured with a

F & M Dual Column Temperature Programmed Gas Chromatography,

Model 720, with a Carbowax 2CM or Apiezon L column.

Constant Temperature Bath

Solutions for which the pH was determined were immersed in a

Sargent (S-848cç) high precision thermostatic water bath equili­ brated at 35 ±0.01°.

Weight Determination

All measurements were made on a Mettler Model B5H26 balance with 1 mg sensitivity. pH Measurements

The pH of all solutions was determined at 35 .0° with a

Beckman Research pH Meter (No. IOI9OO) which was standardized with

two Beckman Buffers of a pH within the range anticipated of that

for the solutions to be measured. 4l

C) Hydration of Isobutyraldéhyde

Solution Preparation

Buffer Solution: An example of the procedure used in prepar­ ation of buffer solution is given. Into a 100 ml volumetric flask,

1.0431 g (0.01031 mole), of N-methylmorpholine and 24.53 ml of

0.2038 N perchloric acid (O.OO5OO mole) was added. Sodium chlor­ ide previously weighed in a 50 ml beaker, 0.288 g (0.00492 mole), was added by use of a powder funnel. That remaining on the beaker and funnel was washed into the flask with sufficient water to in­ crease the liquid volume to the 100 ml line. Properties of this solution are— pH of 7.440, measured at 55°, free buffer concentra­ tion of 0.053 M and protonated buffer concentration of 0.050 M.

Aldehyde Solution: A 1 to 1 by volume mixture of isobuty­ raldéhyde and tetrahydrofuran was prepared by syringing 1.0 ml of each solution into a 5 ml Bantum ware flask. The purpose of this solution is to increase the rate at which isobutyraldéhyde dis­ solves in the aqueous buffer solution at the initiation of each kinetic run.

Method

The rate of reaction was observed spectrophotometrically by recording the change of optical density with time at the .maximum 0 5 b n -* 71* absorption for isobutyraldéhyde, = 2850 A. All spectrophotometric measurements were made with a Gary (Model l4)

Spectrophotometer. Into two Perkin-Elmer #17677 matched 1.0-cm 42 uv cells, -which were thermostated at 35° ty use of a Haake constant temperature circulating bath, -was placed 5 -0m l of aqueous buffer solution. Into the reference cell 25 m-1 of tetrahydrofuran was syringed. After establishment of thermal equilibrium at 55°» a process requiring approximately 15 minutes, the reaction was ini­ tiated by syringing 50 |il of the i s obutyr aide hy de -t et r abydr of ur an solution into the sample cell, gently shaking the sample cell, placing the sample cell into the spectrophotometer, and finally triggering its recording system. This initiation process prevented the experimental observation of the first seven to fifteen seconds of the reaction. The data consists of a plot of the optical den­ sity, the ordinate, vs time. The initial aldehyde concentration was calculated from the equilibrium constant, Kjj, and the ob­ served aldehyde concentration at equilibrium, Ag. time scale, the abscissa, was determined by first evaluating the chart speed,

5 .00 * 0.01 sec per chart unite or O .1692 cm/sec, and secondly, by extrapolating the observed (optical density)/(time) curve to the calculated initial optical density and designating this point on the abscissa at t^. In order to determine the most represen­ tative rate constant, a least squares method, utilizing eight readings from each kinetic run, was employed. See appendix for computer program. h-3 d) Determination of the Micro-pKa Value of the Diamines

Preparation of Solutions

Solutions were prepared by employing standard, accepted tech­ niques. Concentrations of amine stock solutions were determined by potentiometric titration utilizing a Beckman Zerometric II pH meter. See tables— for solution compositions.

One drop of the standard, a 10^ methanol/water solution was added to the nmr tube of only the first solution of each amine investigated. The nmr spectra of the remaining solution of each amine were compared to the standard signal generated in the first solution. After recording the nmr spectra for these remaining solutions of the monoamines and symmetrical diamines, one drop of the standard methanol solution was added to each and the nmr spec­ tra were again recorded. The results obtained from both methods were within experimental uncertainty, ±0.5 ops or l.C^ of

Because the adding of the standard solution to the nmr tube of the amine solutions changes the acid concentration, the results from the first method were recorded. When indicated one drop of a sodium hydroxide solution is also added to the first solution of each amine.

Solutions of H-dimethyl-1,4--butanediamine and N,N-dimethyl-

1,5-pentanediamine were prepared by weighing the anhydrous amine in a 10 p,l Hamilton syringe and adding it to an nmr tube. Water, a sodium chloride solution, and perchloric acid were added to the nmr tube by use of a 100 m-1 Hamilton syringe, see Tables 15-16 for details. kk

All nmr measurements were recorded on a Varian, Model A-So,

NMR Spectrometer, with the probe thermo stated at 55 ^ 1°, em-

, 90 ploying methanol as a standard, 6 = 5 •

E) Imine Formation

Collection of Spectrophotometric Data

The method of collecting data for determination of carbinol- amine equilibrium constants, Kq , carbinolamine-imine equilibrium

constants, Kj, and the rate constants for dehydration of carbinol-

91 amine, kg, have been described previously.

All kinetics and equilibrium data were collected with a

Durrum-Gibson Stopped-Flow Spectrophotometer. The kinetic data,

consisting of a transmittance vs time plot electrically displayed on a recording cathode-ray oscilloscope, were photographically recorded. The instrument was operated as described in the opera­ tion and service manual. During all operations the cuvette and bath were maintained at 55 0.1°. For a schematic description

of the Durrum-Gibson instrument see reference 9I.

Collection of pH Data

The pH for each set of solutions was detennined. The pH of

an e qui volume mixture of the amine solution and water, measured at

90. N. S. Bhacca, L. F- Johnson, and J. N. Shoolery, Varian NMR Spectra Catalog, Varian Associates, Palo-Alto, Calif., 1962.

91. F. A. Via, A Kinetic Investigation of the Reaction of Iso­ butyraldéhyde and Methylamine, M.S. Thesis, The Ohio State University, 1968, pp 9 -12. 55°, is considered to be the initial pH at to. The pH of an equivolume mixture of the amine and the aldehyde solution, measured at 55° after establishment of all equilibria, is con­ sidered to be the final pH at tg. The average of the initial and final pH is considered to be the approximate pH for each set of solutions in the cuvette during a kinetic run. This average pH is recorded "with its range, e.g. if the initial pH were ten and the final pH were nine then the average would be recorded as 9*5

± 0.5. CHAPTER IV

RESULTS

A. The Hydration of Isobutyraldéhyde

Derivation of kinetic expression

It is assumed that the hydration of this aldehyde is a

first-order reversible reaction.

Symbols:

A = concentration of aldehyde at time t

Ag = concentration of aldehyde at equilibrium

Aq = concentration of aldehyde at time to

H = concentration of hydrate at time t

k = rate constant for hydration

k* = rate constant for dehydration

A H (20) «tTT-

Kd = ^ (21)

^ = kA - k*H (22)

Since only A is present at to

H = Ao - A (23)

46 47

Thus, eq.. 22 becomes

- %dt = (k + k»)A - k»Ao (24) Integration of eq. 24 yields

When equilibrium is established

kAe = k’He = k*(Ao - Ae) (26) Ao = Ae

Substituting eq. 26 into eq. 25

(k + k«)t = In (ref. 92) (2?)

Using eq. 2 1 , eq. 23 , and k^^g = k + k*.

The experimentally observed rate constants, k^-^g, which are

the sum of the rate constant for hydration of the aldehyde and the

rate constant for dehydration of the hydrate, are listed in Table 79 2 . As reported in Chapter II, Bell and coworkers state that this

rate constant, for studies of a cet aldehyde, can be expressed by

eq. 29, which includes terms for acid and base catalysis.

92. A. A. Frost and R. G. Pearson, Kinetics and Mechanisms, 2nd ed., John Wiley and Sons, Inc., N. Y., I961, p. 1Ü6 . TABLE 2 48

Observed Rate Constants for the Hydration-Dehydration Reaction

Solution § of Runs k STD obs sec”^

64-1 8 0.06978 0.0036

70-1 5 0.04873 0.0041

70-2 4 0.03892 0.0010

64-2 10 0.02641 0.0045

64—5 2 0.01406 0.00043

64—6 5 0.01189 0.0025

64—3 9 ' 0.0460 0.0025

70-3 5 0.02882 0.0013

70-4 5 0.02116 0.0077

64—4 5 0.1073 0.0106

70-5 5 0.08133 0.0070

70-6 4 0.06808 0.0027

70-7 5 0.05425 0.0012

70-8 5 0.03614 0.0025

70-9 5 0.02730 0.00088 k9

^obs " ^AO ^ '*' (^9)

■where :

k^O ~ the uncatalyzed rate constant

^AH ~ the specific acid catalyzed rate constant

k^OH “ the specific base catalyzed rate constant

^AB ~ the general base catalyzed rate constant

^ABH “ the general acid catalyzed rate constant

B = concentration of free buffer

BH = concentration of protonated buffer

Bj = B + BH

A plot of the observed rate constant against the total buffer

concentration. Figure 1 , indicates the applicability of eq. 2 9 -

The intercept, -which can be expressed by eq. 3 0 , and increases

■with increasing pH, indicates that the hydration of this aldehyde

is subject to specific base catalysis. The existence of general

acid-general base catalysis is indicated by the different slopes

of the lines -with increasing buffer concentration at constant pH.

intercept ky^tH"*"] + k^QgCoH] + k^Q (50)

81 Hine and Houston determined the rate constant for specific acid

catalysis, k ^ = l.lT x 10^ M “^sec"^. In this investigation,

conducted under neutral to basic solution conditions, pH 6.^1--8 .0 ,

the contribution of specific acid catalysis to the observed rate

constant -was "within the experimental uncertainty; and therefore

eq. 50 is rearranged into eq. 51 and the remaining four rate 7.791 0.10

obs

7.44

7.179

6,486

0.02 0.04 0.08 0.10 Total Buffer Concentration

Figure 4 . A plot of the Observed Rate of Hydration-Dehydratlon for Isobutyraldéhyde vs. the Total Buffer Concentration vn O 51

^obs ■ “ ^AO '*' + ^ABH^®^ (5l) constants are evaluated, see Tables 3 and 4 , by employing a least squares program, determining the smallest percent standard devia­ tion, for a four-termed linear equation. The general base catalyzed rate constants for the dehydration of isobutyraldéhyde hydrate by the aliphatic primary and tertiary amines used in this investiga­ tion of imine formation are estimated graphically (see Figure 5)*

This BrjzJnsted plot was constructed by placing a line of slope 0.5 at the point for N-methylmorpholine. This slope wa.s used because 93 82 Bell reports for 1 ,3 -dichloroacetone and Pocker for isobuty­ raldéhyde at 0° slopes of 0.5 and 0 .4-5, respectively.

9 3 * R- P- Bell and M. B. Jensen, Proc. Roy. Soc. A, 26l, 38 (1961). :------— TABLE 3 Results for the Hydration-Dehydratlon Reaction of Isobutyraldéhyde and Its Hydrate

k . obs^ pH [B]^ Cb h '^]^ — ± # sec sec M M

1 0.06978 0.06973 7.440 0.05210 0.04915 2 0.04873 0.04868 7.440 0.02084 0.01966 3 0.03892 0.03887 7.440 0.01042 0.00983 4 0.02641 0.02593 6.486 0.01022 0.08847 5 0.01406 0.01358 6.486 0.00204 0.01769 6 0.01189 0.01141 6.486 0.00082 0.00708 7 0.04600 0.04587 7.045 0.03008 0.06881 8 0.02882 0.02869 7.045 0.01203 0.02752 9 0.02116 0.02103 7.045 0.00602 0.01376 10 0.10728 0.10726 7.791 0.06891 0.02949 11 0.08133 0.08131 7.791 0.02756 O.OII80 12 0.06808 0.06806 7.791 0.01378 0.00590 13 0.05425 0.05415 7.179 0.03742 0.05898 14 0.03614 0.03604 7.179 0.01497 0.02359 15 0 .027.10 0.02720 7.179 0.00748 0.01180 — Where k^=l, 4ô7 M ^ sec — Accounting for dilution by addition of 50]il VI of aldehyde solution to 3•0 ml of buffer solution. tV) TABLE 4 Catalytic Constants for the Hydration of Isobutyraldéhyde in Aqueous Solution at 35°C

where X = A X = h X = d

-3 -1 -3 -1 7.69 X 10 ^sec ^ 2.3 X 10 sec 5.4 X 10 sec

3.20 X 10"*^M”^sec“^ 9.6 X 10^M“^sec"^ 2.24 X lO^M ^sec ^ \ o H 1.47 X lO^M^^sec"^ 4.4 X lO^M ^sec ^ 1.03 X 10^M“^sec"^ XH -1 -1 0 . 22 M sec 0.51 M"“^sec”^ > 4 0.724 M “^sec”^ k» 9.83 X 10 ^sec ^ 3.0 X 10 ^sec ^ 6.9 X 10 M sec XBH

— Values from ref. 81. — Values for N-methylmorpholine

« 1.0

0.6

0.2

9

- 0.2

- 0.6

— 1.0

pKa

Figure 5 * A plot of the General Base Catalyzed Rate Constant for Dehydration vs. the pKa of the Protonated Catalyst. TABLE S 55 Rate Constants for Dehydration of Isobutyraldéhyde Hydrate

AMINE- R-NH^ ^ -1 « 35- log k sec

CHjCH^CHj 10 .21^ 1.03 10.8

MeOCCH^)^ 9.83^ 0.91 8.04

MeOCH CH 9.09^ 0.51 3.24

(MeO)jCHCH2 8.35^ 0.15 1.42

FjCCH^ 5.52^ -1.40 0.04

Me^N(CHp^ 9.45^ 0.67 4.68

Me^N(CH^)g 10.05® 0.93 8.52

Me^NCCH^)^ 1 0 .29® 1.08 12.1

Me^NCCH^)^ 10.51® 1.17 14.8

Me^m(CH^)^ 6.45® -0.77 0.17

Me^mCCH^)^ 8.00® -0.02 0.96

Me^&CCH^)^ 8.67® 0.30 2.00

Meg&(CEg)^ 9 .52® 0.70 5.02

N-Methylmorpholine 7.27 -0.30® 0.507

— Determined in present investigation. — Values from ref. Sh. — J. H. Jensen^ unpublished obsearvations 56

B. Determination of Micro pKa Values for Diamines

Derivation of expression for evaluation of ia>iR data

Symbols :

Ce m t ] = total diamine concentration

Ce m UP] = concentration of unprotonated diamine

[EMlp] = concentration of monoprotonated diamine

[EMlPp] = concentration of primary protonated diamine

CEMLP-t] = concentration of tertiary protonated diamine

f = fraction of monoprotonated diamine in "which the tertiary nitrogen is protonated

LeM2P] = concentration of diprotonated diamine

AU = nmr chemical shift of the methyl protons of an unprotonated amine

®bbs - the absolute value of AU - Aobs

Ap^ = chemical shift of the methyl protons (using Au as standard) n atoms distance from a protonated primary nitrogen

ATn = chemical shift of the methyl protons (using AU as standard) n atoms distance from a protonated tertiary nitrogen atom; when n = 0, subscript not always.used

ADjj = chemical shift of the N-methyl protons using AU as standard) of a diprotonated unsymmetrical dia­ mine containing n number of methylene groups

^obsn “ experimentally observed chemical shift using methanol as standard

[a1] = concentration of perchlorate ion, which is the anion of the acid used as proton source

Derivation of expression to determine [EMUP], CemIP], and [em2p] . 5T

From definition

[EMI] = ÜMÜP] + [EMlP] + tEM2P] (3 2 )

From definition of pKa

K. ■ (»)

Eg = [M-PKh*] (3 4 ) CeM2P]

The following approximation of electrical neutrality can be made for the concentration of OH' or H"*" is usually less than 10'* M whereas A, [emIP], and [em2p] range from 5 x 10"^ to 5 x 10"^ M.

2[em2p] + [smp] + =i o a r î + Ui] (35)

Dividing eq. 53 by eq. 3^:

^ ^ [EMUP][EM2P] Ks Ce m i p ]^

Solving eq. 35 for [EM2p] , eq. 3 ^ for [EMUP] and substituting in eq. 32

(3T)

Solving eq. 37 for [emIP] : minus sign eliminated for CeMIP] is a positive number.

-[.e m t ] + \ Cemt]^ - - 2) - CemtDCai]) [EMIP] = -- :------(38) U(Ki/Ka) - 1.0

Derivation of expression evaluating f. 58

From eq. 52 and the assumption of the linearity of chemical shift

From definition of f; ' ■ S

From definition of Au, eq. 5 9 , ^0, and 4l:

[EMT]Do13s = fÜMlP] (At) + (1 - f) [EMIP] (ApJ + [eM2P] (AdJ (42) [emt]Dobs - [EM2P] (ADn) = [EMIP] ( f( At - Ap^) + ApJ

Experimentally [EMT], Dobs, [EMIP], [EM2P], and ADn can be deter­ mined and a plot of [EMT] Dob g “ [EM2P] AD^ vs [EMIP] wotild yield an intercept of zero and a slope which can be expressed by eq. A5.

Employing the concept of molecular linearity, i.e., choosing model

slope = f(AT - APJ + AP^ (45)

compounds, AT and Ap^ can be estimated and the expression, eq. 4 5 ,

can be evaluated for f .

Changes in chemical shift, relative to methanol, with changing mole fraction of protonated amines, at constant total amine con­

centration, were observed and recorded. Preliminary experiments,

checking the standard and evaluating the salt affects, indicate

the acceptability of this experimental method. The chemical shift 59 of the carbon-bound protons for methanol, relative to external tetramethylsilane, remained constant as the acid concentration of the solutions varied from 0.0 to O.5 molar. The acid- containing solutions -were prepared by diluting a 1.000 molar solution of perchloric acid to the desired concentration. One- half ml of these prepared solutions was placed in an nmr tube and one drop of a IC^ methanol/water solution was added. Although the addition of the standard solution to each nmr tube would de­ crease the acid concentration by approximately IC^, this small change in concentration is considered insignificant in respect to the purpose of this experiment, and also the acid concentrations were varied over a relatively large range, 6.0 to O.5 molar. The

"delta" chemical shift, ^obss Table 6 for the third solution of methylamine, =57.0 cps, and the chemical shift for methanol, the carbon-bound protons, relative to external tetramethylsilane, remained constant when the concentration of sodium chloride was varied from 0.0 to O.5 molar. Thus, it appears that the observed

chemical shifts are not sensitive to changes in ionic strength or proton concentrations that occur under these experimental condi­ tions. The change in chemical shift resulting from converting an

unprotonated amine to a totally protonated amine, for a pri­ mary amine, AT^ for a tertiary amine and AD^ for a diamine, for

the model compounds and for the unsymmetrical diamines are given

in Tables 6 -1 0 . These amines were generally protonated to the

extent of 75-98^ and AT^, Ap^, and AD^ were estimated by 60

extrapolating to lOC^ protonation. If the amine is completely

protonated the resulting nmr spectrum generally contains a very broad doublet, due to spin-spin splitting by the nitrogen-bound

proton; these results are not recorded. Data on N,lT-dimethyl-

alkyl diamines are used in eq. 44 to calculate the slope and data

on the model compounds are used to estimate ATi and i^Pn+a*

slope^ = f(ATi - APn+s) + Apn+g (4 4 )

ATi is the change in chemical shift for the o'-methyl protons

bound to a tertiary amine with protonation of the amine. ATi in

eq. 4 4 , is the contribution to the observed change in chemical

shift of the nitrogen-bound methyl protons , this is caused by + protonation on the tertiary amine, e.g., Me^NHCHaCHsWHs. Tri-

methylamine yielded a value of 4l .4 cps for AT% and was chosen as

the representative model of N-dimethylalkyl diamines, ATi. Per­

haps dimethylethylamine would be a "better" model than is tri-

methylamine; however, it was not readily available and as shown on

Table IT, the difference between model values and observed values

for AD^ and AD^ is small.

Ap^ is the change in chemical shift for methyl or methylene

protons, n atoms distance from a primary amine, with protonation

of the amine, AP^ in eq. 44 is the contribution to the observed

change in chemical shift of nitrogen-bound methyl protons that is

caused by protonation on the primary amine of a N,lT-dimethylalkyl + diamine. E.g., for MegNCHaCHaRHa, Ap is P4, 2 -Methoxyethylamine TABLE 6

Chemical Shift of Protonated Methylamine*

Volume of 1.00 J6J Aobs- HCIO, Free amine 1 Number % cps

1 0.0 96 66.4 2 3.54 75 62.0 3 7.09 50 57.0 4 10.63 25 52.3 5 12.40 12.5 49.7

Vol. of sol'n Wt. of NaCl Cone, #3 gr. of NaCl Aobs^ ml M

10.0 0.058 0.10 57.0 10.0 0.146 0.25 57.0 10.0 0.292 0.50 57.0

— Each solution contained 5*0 ml of a 2.835 M amine solution and a total volume of 55»0 ml. 62

65

60

A?2 = 20 ops 55

50

20 60 80 Percent Unprotonated Amine

Figure 6# Correlation of Chemical Shift with Percent Unprotonated Methylamlne TABLE 7 Chemical Shifts of Protonated Trimethylamine ~

Volume of Volume of Free amine Aobsi # I.OO4N HCIO 1.0 M NaCl % ml ml

1 0 .0^ 15.3 100 42.4

2 3.83 11.48 75 31.4

3 7.65 7.65 50 21.84

4 11.48 3.83 25 11.64

5 15.00 0.31 2 1.8

— Each solution contains 5*0 ml of a 3.O6OM amine solution and a total volume of 100 nil. — Nmr tube contained one drop of 0.2^ NaOH in O.J ml volume 6k

40

30

20 AT, = 41.4 cps

10

20 60 80 Percent Unprotonated Amine

Figure 7 . Correlation of Chemical Shift with Percent Unprotonated Trlmethylamlne TABLE 8 Chemical Shifts of Protonated Triethylamine^

Volume of • Free Amine Aobs. b AobSg 1 . 004 N HCIO ^ amine X gr # ml 4 % ops cps

1 1.0953 0.0® 100 54,0 49.5

2 0.6940 1.72 75 42.2 44.2

3 0.7086 3.50 50 32.2 40.7

4 0.8000 5.93 25 20.0 35.0

5 0.7644 6.59 12.5 15.0 33.2

— Each solution has a total volume of 50.0 ml. — T, and obs^ refer to the a-methylene protons- — Contains one drop of 12 N NaOH solution in 0.5 ml volume. 66

50

40 AT2 - 18,0 cps

3 (0 I 30

a T = 43*5 ops 20

100 Percent Unprotonated Amine

Figure 8* Correlation of Chemical Shift with Percent Unprotonated Triethylamine TABLE 9 Sj Chemical Shifts of Protonated n-Propylamine^

Volume of Free Aobsi Aohs 2 Aobs 2 1 .004 N HCIO, amine # ml 4 % cps cps cps

1 0.0 96^ 67.9 138.0 166.0

2 2.52 75 63.5 135.3 165.0

3 5.04 50 57.3 132.5 163.7

4 7.55 25 51.5 128.0 162.4

— Each solution contained 10.0 ml of a 1 .0 0 7 M amine solution and a total volume of 50.0 ml. ^ Calculated from pKa 10.22 at 35°, ref. 5b. 68

60 60

o 0 H AP, = 24.lops 50 r4 0} g I 1 40

30 30

10 5030 70 90 Percent Unprotonated Amine

Figure 9 * Correlation of Chemical Shift with Percent Unprotonated n-Propylamlne TABLE 10 69 Chemical Shifts of Protonated 2-Methoxy ethyl amine

Volume of Free Aobs]_ Aobs 2 Aobs^ 1.004. N HCIO, amine # ml * % cps cps cps

1 0.0* 100 70.2 29.1 6.43

2 2.47 75 64.5 26.3 5.95

3 4.94 50 58,6 23.4 5.44

4 7.42 25 52.5 21.2 5.15

5 8.90 10 4,68

— Each solution contains 10.0 ml of a 0 .989,25 amine solution in a total volume of 50,0 ml, — nmr tube contained one drop of 0 .2 N NaOH solution in 0.5 ml volume. 70

70

60 6.0 AP- = 23*6 cps

50 5.0

o I 40 4.0

30 3.0

20 AP o — 11*1 cps 2.0

10 Percent Unprotonated Amine

Figure 10. Correlation of Chemical Shift with Percent Unprotonated 2-Methoxyethylamine TABLE 11 71

Chemical Shifts of Protonated Nj N * j N * —T et r amethyl ethylene diamine^

Volume of l.OOOM Free c perchloric acid amine Aobs^ # ml % cps

1 0.0 99.3 75.3

2 3.0 87.9 68.9

3 5-99 75.8 62.1

4 8.99 63.7 56-7

5 11-98 51.6 50.5

6 14.98 39.5 43.9

7 17.97 27.4 37.8

8 20.97 15.3 30.1

9 23.96 3.2 —

— Each solution contains 10.0 ml of a 1.238 M amine solution in a total volume of 60.0 ml. — Refer to the nitrogen-bound, methyl protons — Calculated from pKb of 5 -33> ref. 2d 72

Aobs

40 a T-j^ = 52.6ops

10 30 50 70 90 Percent Unprotonated Amine

Figure 11 . Correlation of Chemical Shift of the Nitrogen- bound Methyl Protons with Percent Unprotonated, N,N ,N* »N*-Tetramethylethylenediamine TABLE 12 ^ Chemical Shifts of Protonated — Tetramethyl-1, 3~propanediamine*

Volume of 1 .004; N Amount of % Free Aobs^ # HCIO. acid NaCl amine cps A gr %

1 0.0® 1.18 100 60.8

2 2.54 1.03 87.5 55.86

3 7.61 0.738 62.5 4 5.2

4 10.14 0.590 50.0 39.2

5 12.68 0.443 37.5 32.9

6 17.75 0.148 12.5 21.84

7 19-87 0.0236 2.0 17.8

— Each solution contains 10.0 ml of a 1 .018’M amine solution in a total volume of 50.0 ml. — The nitrogen-bound, methyl protons — Nmr tube of this solution contains one drop of a 0.2 N NaOH solution in 0.5 ml of solution T4

60

50

40

30

20

20 40 60 80 100 Percent Unprotonated Amine

Figure 12 . Correlation of Chemical Shifts of the Nitrogen- bound Kethyl Protons with Percent Unprotonated Amine, N,N,N* ,N*-Tet3ramethyl-l,3-Propanediamine TABLE 1 3 75

Chemical Shifts of Protonated N, N-D ime thy lethyl ene diamine ^

Volume of Free 1.004 2i Aobs HCIO. amine ^ [EM2P] AD g # ml4 % cps

1 0 .0** 100 49.6 0.0 0.0

2 2.53 87.5 45.3 0.05067 0.9118

3 5.05 75 40.8 0.1012 1.820

4 7.58 62.5 36.14 0.1512 2.749

5 10.10 50.0 31.0 0.1908 3.528

6 12.63 37.5 23.86 0.1512 2.814

7 15.15 25.0 16.84 0.1012 1.883

8 17.68 12.5 9.36 0.05067 1.006

— Each solution contains 10.0 ml of a I.OI4 M amine solution in a total volume of 50.0 ml and enou^ NaCl for }i=0.4 — Nmr tube contains one drop of 0.2 N NaOH solution in 0.5 ml volume 76

40

30

Aobs ADp = 4?.3ops 20

10

10 30 50 70 90 Percent Unprotonated Amine

Figure 13. Correlation of Chemical Shift with Percent Unprotonated N ,N-Dimethylethylenediamine (5 )

3.0

(6 )

2.5

CM

(7 )

(3 )

1.5

1.0 (8) slope = 18*47 ops (2)

0.02 0.06 0.10 0 . 1 4 0 . 1 8 ( EHIP) Figure 14. Evaluation of Equation 42 TABLE 14 78 Chemical Shifts of Protonated N,N-D imethyl-1, 3 ” Propane diamine*

Volume of 1.004 1 Free HCIO4 Aobs [e m p i ] amine Ce M 2P] ADg ml # % cps M

1 0.0^ 100 68.6 0.0 0.0

2 2.66 87.5 65.4 0.0529 0.674

3 5.32 75 61.8 0.1046 1.404

4 7.98 62.5 57.5 0.1523 2.189

5 10.64 50.0 52.5 0.1790 2.621

6 13.30 37.5 45.6 0.1523 2.189

7 15.96 25.0 37.6 0.1046 1.489

8 18.62 12.5 29.4 0.0529 0.737

— Each solution contains 10.0 ml of a I.O64 M amine solution in a total volume of 50*0 ml and enough NaCl to yield an ionic strength of O.4O. — Nmr tube contaas one drop of 0.2 N NaOH solution in 0.S ml volume 79

60

50

Aobs

40

30

10 Percent Unprotonated Amine.

Figure 15» Correlation of Chemical Shift with Percent Unprotonated N,N-Dimethyl-1 ,3-Propanedlamlne 8o

2.5

CM

CM

(7 )

(3)

1.0 slope = 14.29 cps

(8) (2)

0.5

0 . 0 4 Ô7Ô8 0.12 0.16 (EMIP)

Pigtire 16. Evaluation of Equation 42 81 TABLE 15 Chemical Shifts of Protonated N, N-D imethyl-1, 4~butanediamine®

Vol. of Vol. of CEMT]l>obs - 1.004 N 1.0 N Free [e m i p ] Aobs [eM2P] AD^ HCIO. NaCl amine # ml ml % cps M

1 0.0 0.1096 100^ 45.0 0.0 —

2 0.0137 0.0959 87.5 41.2 0.02098 0.316

3 0.0274 0.0822 75' 37.5 0.04023 0.587

4 0.04113 0.0685 62.5 33*6 0.05551 0.7916

5 0.0548 0.0548 50.0 29.0 0.06189 0.8673

6 0.0685 0.0411 37.5 23.0 0.0556 0.8034

7 0.0822 0.0274 25 16.0 0.04023 0.629

8 0.0959 0.0137 12.5 9.5 0.02108 0.337

— Each solution has a total volume of O.64 ml and a total amine concentration of O.O8 6 M. — Nmr tube contains one drop of 0 .2 N NaOH solution 82

40

30

Aobs

ADi. =42.0 ops 20

10

10 30 50 70 90 Percent Unprotonated Amine

Figure 17. Correlation of Chemical Shift with Percent Unprotonated K , N-Dimethy 1-1,4-Butanedlamlne 83

0.8 (6) (4) f & g I (7) * .o I (3 ) slope = 14.52 cps 0.5

0.4

0.3 (2)

0.02 0.03 0.04 0.05 (EMIP) Figure 18. Evaluation of Equation 42 TABLE 16 8k Chemical Shifts of Protonated N,N-Dimethyl-1, 5“Pentanediàmine M.

Vol. of Vol. of Free # Aobs [e m i p ] 1 . 0 0 4 N 1.0 N amine [EM2P%ADg) HCIO. NaCl % cps. M ml ^ ml

1 0 . 0 “ 0 . 1 0 2 1 0 0 4 4 . 0 0 , 0 —

2 0.0129 0 . 0 8 9 9 8 6 . 5 4 1 . 7 0 . 0 2 0 3 0 . 1 5 7

3 0.0257 0 . 0 7 7 0 7 3 . 1 3 8 . 4 0 . 0 3 6 7 0 . 3 1 3

4 0 . 0 3 8 6 0 . 0 6 4 2 5 9 . 7 3 5 . 0 0 . 0 4 7 0 . 0 . 3 5 2

5 0 . 0 5 1 3 0 . 0 5 1 3 4 6 . 3 2 9 . 2 0 . 0 4 9 0 0 . 4 1 7

6 0 . 0 6 4 2 0 . 0 3 8 6 3 2 . 9 2 2 . 0 0 . 0 4 2 3 0 . 4 1 2

7 0 . 0 7 7 0 0.0257 1 9 . 4 1 4 . 1 0 . 0 2 8 3 0 . 3 1 2

8 0 . 0 8 9 9 0.0129 6 .0 6 . 4 0 . 0 0 9 5 0 . 1 0 0

— Each solution has a total volume of 0 . 5 9 3 1 ml and a total amine concentration of O.O809 M. — N.m.r. tube contains a drop of 0.02 M NaOH solution in a total volume of O.5 ml. 85

40

Aobs AD< = 41,3 ops 20

10

10 30 50 70 90 Percent Unprotonated Amine

Figure 19, Correlation of Chemical Shift with Percent Unprotonated K , N-Dimethyl-1, S-ï’entanediamine 86

(5 ) (6) 0.4

(4 ) (7 ) # (3 ) CM

O (2)

0.1 slope = 8.72 ops

0.01 0.02 0.03 0.04 (EMIP)

Figure 20. Evaluation of Equation 42 87

■was used to determine Ap^, AP5 , APg, and APy. In this model com­ pound the methyl protons are three carbon and one oxygen atom distance from the protonated primary amino group, in the dimethylethylene diamine the methyl protons are three carbon atoms and one nitrogen atom distance from the protonated primary amino group. However, differences between model values and observed values for some of the diamines are small, indicating that 2 - methoxyethylamine is a reasonable choice for a model compound.

AP4 is determined as 1.95 cps (see Table IT), APg is estimated from AP4 and the fall-off factor, the ratio Ap^/APg = 2 .1 2 . APg and APy are estimated similarly, APn = APn-i/2.12. This calcula­ tion uses the ass'umption that the observed change in chemical shift is a linear function of the number of atoms located bet'ween the protonated nitrogen and the methyl protons, that is Ap^yAPg =

AP4/AP5 = APg/APg = APg/APy. Fall-off factors have been experi­ mentally determined, 2 -methoxyethylamine, AP^/Ap^ = 2.12 and n- propylamine, Ap^/APg = 1 .7T, APg/APg = 2 .67, with an average of

2 .1 9 . The fall-off factor determined for 2-methoxyethylamine is reasonably close to the average and was therefore used.

ADj^ is the change in chemical shift of the If-methyl protons from an unprotonated to a completely diprotonated H,If-dimethyl- alkyl diamine with n being the number of methylene groups between the two nitrogen atoms. From the assumption of the linear addi­ tivity of molecular properties, eq. 45, AD^^ can be estimated.

ADn = ATi + A%4g (45) TABLE 17 88 Observed and Estimated Chemical Shift Values^

Amine ^ 6 (1 ) lit 5 uf tif (7)

Methyl 19.7 a (est.) n-Propyl 24-1 13.6 5.1 1.76 b

2-Methoxy- (est.) (est.) (est.) ethyl 23-6 11.1 1.95 0.91 0.43 0.20

All

Trimethyl 41.4

Triethyl 43.5 18.0

AD^ AD calculated from n obS n

N, N, N ' , N ' -Tetramethyl- ethylenediam ine 2 52.6 45.4 "2 + *^5 N,N,N *,N * —Tetramethyl- 1,3-propanediamine 44-2 3 44.4 ®2 + N, N-D imethyl ethyl ene- diamine 2 47.3 43.4 ^2 + ^5 N, N-D imethyl—1 , 3~ propanediamine 47.6 3 42.3 '2 + N, N-D imethyl-1,4“ butanediamine 42.0 41.8 4 ®7 N, N-D imethyl-1 ,5“ pentane diamine 5 41.6 41.3 <*2 +

~ All values are in ' units of cps 89

Differences between estimated and observed values of for iy,N-dimethyl-l,A-butanediamine, (obs. = 42.0 cps, est. = 4l.S cps) and for N,E-dimethyl-l,5-pentanediamine, (obs. = 4l.5 cps, est. = 4l.6 cps) are small and indicate a degree of validity in the method of employing model compounds. However, differences between estimated and observed values for for N,H-dimethyl-

1,3-propanediamine (obs. = 4J.6, est. = 42.3) and N,N-dimethyl- ethylenediamine (obs. = 47.3, est. = 43.4) are comparatively large. In solving eq.. 45 for the latter two amines, the values for ATi and Ap^+g are made larger in proportion to their percent contribution to AD^, eq. 46 and 47-

The calculated values for f, the fraction of monoprotonated diamine protonated at the tertiary amine, are given in Table l8.

The micro-pKa values for the diamines were calculated in the following manner. All symbols are defined in the Appendix. The first experimentally observable Ka of a diamine is defined by eq.

48. The desired Ka for protonation of the primary nitrogen is

defined by eq. 4 9 . Dividing eq. 4$ by the ratio of [emip„]/[EM1P] TABLE 18 90 Evaluation of f for Me.N(CH ) NH L 2 n 2

Estimated values proportionally corrected n slope fAll ^ n +2

2 45-1 2.1 18.47 0.381^

3 46.6 1.02 14.29 0.291^

uncorrected

4 41.4 0.43 14.52 0.344^

5 41.4 0.20 8.72 0.207^

2 41.4 1.95 18.47 0.419

3 41.4 0.91 14.29 0.330

â ' ■ — Because of a greater probability of absolute uncertainty in AT^ than in AP these values are believed to be a more accurate representation of f. 91 Kêl HgO + Me2N(CH2)niraa ^ > MegN( CHg) BEg + Bq O^ (4$) defines in terms, of experimentally observed values. This identical approach is utilized for calculation of the Ka for ■V. calculation of the Ka for protonation at the tertiary nitrogen

■with the final form defined by eq.. 5 1 -

Ka. = Ka. = Ch +3[EMÜp 3 x ®Mlt ^ [EMlPk] [EMlPt3

The second experimentally observable Ka for a diamine is defined by

eq. 52. The desired Ka for diprotonation at the primary nitrogen is

^ i»>

defined by eq. 53. Multiplying eq. 52 by the ratio [EM1E(;3/[EM1P3 de-

HaO + Me2M(CH 2 )nmîs'^ MegimC CE2) NH2 + H3O+ (53) ' fines Kaj^pin terms of experimentally observed values. The identical

V. - Tf« [EMlPb3 _ [H+3[EMlPb3 ^®M2p ^ [e m p 3 LEM2p 3 ^ method is employed in calculating the Ka for diprotonation at the tertiary nitrogen -with the final form defined by eq. 5 5 * [emip^3 [h'*'3[emipJ , ^ K&M2t “ [EM1P3 " [EM2p 3

These micro-Ka are listed as their negative log, pKa in

Table I9 . 92 TABLE 19

The Micro-■pKa Values for the Diamines “

- [EMlPp] n Ïe m i p ; ^^®Mlt

2 9-448 9.240 9.031 0.619

3 10.05 9.901 9.511 0.709

4 10.287 10.104 9.824 0.656

5 10.512 10.411 9.804 0.793

[EMlPtl P*"^2p ^^®M2t Ce m i p ]

2 6 .446 6.866 6.654 0.381

3 8.00 8.536 8.150 0.291

4 8 • 866 9.330 9.050 0.344

5 9.518 10.201 9.620 0.207

— At ionic strength of around 0.1 95

C. Kinetics and Thermodynamics of Imine Formation

1. Methods of calculation

When aqueous amine is added to an aqueous solution of iso- o butyraldéhyde, the optical density at 2850 A, produced by the

absorbance of the aldehyde, decreases. It is assumed that the

carbinolamine, as all secondary amines, does not absorb appre­

ciably at this wavelength. The rate of formation of imine can thus be determined by following the rate of change of optical „ o density at 2850 A. The Durrum-Gibson Stopped-Flow Spectrophoto­ meter affords rapid mixing of two aqueous solutions and electron­

ically records the transmittance vs time changes of the resultant

solution. Consequently, by conversion of transmittance values to 94 optical densities the rate of change of isobutyraldéhyde con­

centration can be determined. The kinetic expression for deter­ mination of the rate constants ( kg and k_g) by incorporation of

the rate of change of aldehyde concentration assumes the reaction

can be defined by eq. 5 6 .

0 ^ OH ^ (CH3)2CHCH + EKH2 (CHs)2CH-C-m

(56) (CIfe)2CHCH=mi + H2O

K. Base, Chem. Ber., 9 0 , 1251-8 (1957)- (b) M. S. Mellon, Analytical Absorption Spectroscopy, John Wiley and Sons, Inc., Hew York, W. Y., 1950, p. 93* 94

Derivation of kinetic expression.

Symbols used in the expression:

A = isobutyraldéhyde concentration at t, e = 20.86®®

Ag = isobutyraldéhyde concentration after estab­ lishment of equilibrium with imine. e = 20.66

Am = total initial aldehyde concentration ■ A + H + C + [CH+] + I + [IH+]

a = A - Ag

C = carbinolamine concentration at time t

= protonated carbinolamine concentration at time t

[EMT] = the total amine concentration

Le MUP] _ the fraction of the amine that is free (un- LEMTJ protonated

H = concentration- of isobutyraldéhyde hydrate at time t

I = imine concentration at time t

[iH^] = concentration of protonated imine at time t

^2obs “ the observed rate constant for dehydration of carbinolamine

ksHobs = the observed rate constant for dehydration of carbinolamine calculated by assuming that dH/ dt = 0 during the reaction

K^Qal “ the average equilibrium constant for carbinol­ amine formation, (C + Cc H'^'D/a Ce m t ], for each run and is used in the kinetic expression

95- C. Y. Yeh, Equilibrium Constants for Imine Formation from Isobutyraldéhyde and Primary Alkylamines, M.S. Thesis, Georgia Institute of Technology, 19&5, P- 65. 95 Kj) = dehydration equilibrium constant of the aldehyde

Kjjjal “ the average equilibrium constant for carbinol­ amine-imine formation, + ^iMcal

K _ - = the average equilibrium constant for imine formation, [l + IH‘‘']/a Ce m t 3, used in the kine­ tic expression

Average Kq = the average value of the true equilibrium con­ stant for carbinolamine formation, which ac­ counts for protonated amine, c/a Le MUP]

Average K-j- = the average value of the true equilibrium con­ stant for carbinolamine-imine formation, which accounts for protonated amine, [l + C]/a [EMUP]

Mm = total initial amine concentration, [EMT] + I + [lH+] + [C + CH+]

“ " ^Ccal ^ c a l

p = 1 + 1/Kg

6 = Mrp - A^

The mechanism can be expressed

H A (57) ' g A + [EMT] , C + [CH+] I + [IH+]

This investigation reveals that in basic conditions, when pH is greater than 1 0 , it is reasonable to assume that the dehydration

(Kg) and the carbinolamine (Kq ) equilibria are established rapidly relative to the formation of imine. This assumption will be given further consideration.

Kg = ~ (58) 96

From definition of 5, Mj and

6 = [EMT] - A - H = [EMT] - A - (6o) %

Rearranging eq.. 60 [EMT] = 6 + PA (6l)

From (59) and (6l)

C = KccalAto^ = Kccal^C ^ + PA) = KccalC^A^ + ^A) (62)

From differentiation of (58) and (62)

ÉA = (65) dt dt

# = Ko=al(2PA + 6) M (64) dt dt

From the assumed mechanism and the La'w of Conservation of Matter

# + # + § + § = ° («)

Rearranging and substituting (63) and (61|-) into (65),

= 0 (66)

The remainder of this derivation, that of evaluating ^ and in- 96 te grating to yield eq. 67 has been previously described.

^ ^ c a l JS-sXi — (6T)

In eq. 67 ,

= >1 (offi + p)2 + 4opA^ (68)

9 6 . Francis A. Via, M.S. Thesis, The Ohio State University, 1967, pp. 20-2 5 . 97 and

m = ^ [g 4. SccalG + ^ ~ ~ ] (69) •''■Ccal The preceding kinetic expression eq. 67 is modified for those amines -whose rate of formation of imine is rapid relative to the rate of dehydration of aldehyde hydrate. This situation depends on the value of the rate constant, k^, and the pH. For this situation it is assumed that during the reaction,

~ =OandAj=A+C+I. (70)

This assumption requires that eq. 65 be defined as eq. 71 and eq.

60 be defined as eq. 72

# + § + # = ° (Tl)

6 = [EMT] - A (72)

The initially derived kinetic expression eq. 67 can be adapted to the restrictions of eq. 71 and 72 by calculating the total alde­ hyde concentration, A , with the real extinction coefficient of isobutyraldéhyde, = 20 .8 6 , and defining P = 1 instead of

P = 1 + 1/Kjj. Rate constants calculated using this modification are labeled kaHobs-

For condi-tions, pH = 8-9, in which it is not reasonable to assume dH/dt = 0 and the rate of formation of imine is not rapid relative to the rate of dehydration, rate constants are calcu­ lated using both assumptions, kaobs Eind kaHobs, and averaged. 98

To determine which of these three methods of calculation to employ, the ratio of the rate of dehydration, V^, to the rate of imine formation, Vj, is calculated, eq. 75 and 7^* The specific base catalyzed rate constant for dehydration of the aldehyde hy­ drate, k(3Qijj an estimate of the general base catalyzed rate con-

Vd = kaog + k m ^ (75)

V i = k^obs KcM (7^) stant for dehydration of the aldehyde hydrate, k^, for various aliphatic amines (see Table 5)j the estimate of the rate constant for formation of each imine, kg^^g, using eq. 67, and determination of each equilibriiun constant for carbinolamine formation, Kq, come from this study. Under the reaction conditions of this investiga­ tion (pH 5 -0-15) the specific acid catalysis of the rate of dehy­ dration, of the aldehyde is insignificant. When the ratio is greater than five (V^Vj > 5) the first method, assuming Kg es­ tablished is employed; when this ratio is one fifth or less, the second method, assuming dH/dt = 0, is used; and when this ratio is less than five but greater than one fifth (5 ^ V^Vj > 1/5 ) both calculations are made and averaged.

Since the dehydration of isobutyraldehyde and dehydration of carbinolamine are reversible reactions, comparison of the rates of approach to equilibrium would be a more accurate means of de­ termining which of the two methods of calculation of kg to employ than is the comparison of the rate of dehydration of 99 isobutyraldéhyde hydrate, to the rate of dehydration of car­ binolamine, Vj. However, there are three factors which permit the latter comparison. One, the differences in the two methods of calculation are small, k^gobs is usually about 15^ smaller than the value of k^obs calculated from the identical experimental data.

Two, the larger rate constants are compared in both reactions. The rate constant for dehydration of the hydrate, ka, is about 2.3 times as large as the rate of hydration of the aldehyde, kg. The rate constant for dehydration of carbinolamine, kg, is about 15 times larger than the rate of hydration of imine, k-g. Three, the ratio range, 5 > Vd/V% > 1/5, over which both kgghs kggobs are calculated is rather large and consideration of reversible steps would not invalidate the calculations of kg; because the majority of solution conditions that qualified in this range have a ratio near the center of the range, Vjj/Vj ~ 1 .0 .

The equilibrium constants, Kq and Kj, were determined in con­

junction with the kinetic runs. The concentrations of aldehyde,

initial, after establishment of Kq , and final, after establishment

of Kj were determined in the stopped-flow spectrophotometer. The 5 method of calculation is similar to that of C. Y. Yeh, for deter­ mining apparent equilibrium constants in which the aldehyde con­ centration is the sum of the free and the hydrated aldehyde, K-jVg^

except that corrections for protonated amines are made with the

observed pH and the appropriate pKa. Another necessary modifica­

tion of Yeh's method is the utilization of the real extinction ' 100 coefficient of isotutyraldehyde in determining the carbinolamine equilibrium constants, Kq, ■when the hy drat ion-dehy drat ion equi­ librium, Kjj, for aldehyde is not established. The calculated pseudo first-order rate constants for dehydration of the aldehyde hydrate eq. 7 3 } affords an estimate of the half time for dehydra­ tion. This half time can be compared to the time at •which obser-

t = 0-693 g rate constant vations are recorded and if the dehydration reaction has proceeded as long as its second or third half life, the equilibrium is as­

sumed established and an apparent equilibrium constant is calcu­ lated, KQg_. If the observations are recorded •well before the first

half life of dehydration, the change in hydrate concentration with

respect to change in time is assumed as zero, dH/dt = 0, and the

equilibrium constant K q is calculated.

2. Monoamines

Imine formation from isobutyraldéhyde and primary alkyl mono­

amines was studied. The first step in this reaction involves "the

stoichiometric addition of the amine across the carbonyl bond

forming the intermediate, a carbinolamine. The rate of carbinol­

amine formation is too fast to observe with the stopped-flow

spectrophotometer, but the equilibrium constants, K q , were deter­

mined. The first-order rate constants for dehydration for imine

formation were determined. The constants are given in Tables 20-

2 8 . All symbols are defined in the Appendix. 101

a. The reaction of n-propylamine and isoTautyraldehyde.

The l4 recorded determinations of K^obs averaged and corrected

for protonated amine, eq. 75 > "by using the average pH and the pKa sb of 10.21, \determined at zero ionic strength. These and other re­

sults obtained with n-propylamine are listed in Table 20. The ten

recorded observations for determination of were treated as

ave. Kj = (75)

7,91 previously described assuming that the hydration-dehydration

equilibrium, Kjj, for isobutyraldéhyde is established before the

observations can be recorded. From this investigation it is

learned that the major contributing term to the rate of dehydration

of i sobutyr aldehyde hydrate at pH 11.4 is that of hydroxide-ion

catalysis ( symbols are defined in the Appendix). This pseudo

Vd =

Vd = Y (2.24 X 10*^ ^ sec ^)(5 x 10 \) Aq + very small term

Vd = 48 sec”^ Ao

first-order rate constant affords an estimate of the half time for

dehydration of 0.014 sec.

' ÜT&? = O.Olkseo.

Thus, the first observations for Kg^^g, extrapolated to zero time 102 from data taken after about 0 .2 -0.5 sec, occur after the seventh half-life of dehydration and it can be assumed that this equili­ brium is established. Observed values of were averaged and corrected for protonated amine, eq. j6. The determined value for ave Kj is modified by averaging it with the two values determined ave Kq = K ( ® ) i A (76) sa previously by Yeh. These average equilibrium constants (ave

KjQbs ave Kqq^q) are used to obtain the calculated constants which are employed in the kinetic expression.

Kccal = (77)

KiMcal = ave Ki • - Kc^al (78)

The 232 recorded observations for determination of kgobs "were 7 treated as previously described assuming that the dehydration equilibrium. Kg, is established rapidly relative to establishment of the carbinolamineimine equilibrium, Kj. Under these conditions, pH =11.^, the rate of dehydration, to the rate of imine formation, Vj, is greater than ten, Vg/V]- = 48/3.1 =15*

Vj = ksKQcai^M

V- = 9-95 sec“H4.77 M ‘^][0.066 m]a

Vj = 3 "1 A sec ^

The reaction of n-propylamine and isobutyraldéhyde was studied TABLE 20

Rate and Equilibrium Constants for the Reaction of n-Propylamine and Isobutyraldéhyde

Run 4 bs H Cobs # 84 MM # # jj-1

1-5 0.041 0.066 14 96. 5±0.7 10 4.77-0.64 7-10 0.041 0.088

11-14 0.040 0.066 P = 0.046

15-17 0.040 0.066 P = 0.306

18-20 0.040 0.066 P = 0.412

Fe m u p ] ^IMcal *^Ccal ^2 obs Run [EMT] # pH M-1 M-1 sec”"^

1-5 8 11.39^0.13 95.1 4.77 9.99-1.3 0.938

7-10 4 11. 40=*"0.12. 95.2 4.78 9.52-1.5 0.939

11-14 8 II.30-.I 93.6 4.70 9.03-0.28 0.924

15-17 4 II.30-.I 93.6 4.70 9.75-0.51 0.924

18-20 5 II.30-.I 93.6 4.70 10.1-0.94 0.924

Average = 102.9Ï1.0 Average = 5 • 09lo.9

2 Calculated from the average of this study and that (K =106.5 M”^) reported by Yeh^^ ick

over a limited pH range ; hence it was not possible to evaluate

catalytic affects. The ionic strength of the solution was varied

by ten-fold, O.cA-6 to 0.4l M, and the observed rate constant in­

creased by only 13^, see Table 20.

b . The reaction of $-methoxypropylamine and isobutyraldéhyde.

The 50 recorded determinations of were averaged and cor­

rected for protonated amine by using the average pH and the pKa of 5b 9 .8 5 . The 18 recorded determinations of Kgotig observed in solu­

tions of pH 11.4 to 10.6 were treated assuming that the hydration-

dehydration equilibrium, Kjj, for isobutyraldéhyde is established before observations are recorded. The pseudo first-order rate

constant for dehydration of the aldehyde hydrate at pH 10.6, J.J -1 sec , affords an estimate of a half time of O.O9 sec. Thus,

observations for K^obs’ extrapolated to zero time from data taken

after about 0 .2 -0-3 sec, occur after the second to third half-life

of dehydration and it was assumed that this equilibrium is estab­

lished. The remaining 26 recorded determinations of K(jobs>

observed in solutions of pH 9*T to 9*2, were treated assuming

that the change in hydrate concentration with respect to change

in time is zero, dH/dt = 0. Under these conditions (pH = 9*7)

the half time for dehydration is at least 0.4 sec, and the ob­

servations for K^Q^jg are recorded before as much as 13^ of the

hydration-dehydration equilibrium can be established. Observed

values of KQjÿjjg were averaged and corrected for protonated amine. 105 5b This average Kq and the literature value of Kj, 92-3 M "were used to obtain the calculated constants, eq,. 77 and j8, -which are employed in the kinetic expression, see Tables 21 and 22.

The l84 recorded observations for determination of ksobs in.

solutions of pH 11.4 - 10.6 -were treated assuming the hydration-

dehydration equilibrium is established rapidly relative to the establishment of the carbinolamine -imine equilibrium, Kj. At pH

10.6 the ratio of the rate of dehydration, to the rate of

imine formation, V%, is greater than t-wenty-three.

The 244 recorded observations for the determination of k^obs^^ solutions of pH = 9 *7-9*2 -were treated assuming the change

in hydrate concentration with respect to change in time is zero,

and also treated assuming that the dehydration equilibrium is es­

tablished rapidly relative to establishment of carbinolamine -

imine equilibrium. These t-wo rate constants, kgobs and ksHobsj

are averaged when the observed rate constants are evaluated as a

function of pH (see Table 22). As the pH was varied from 9.2 to

9.7} the ratio of the rates of dehydration to imine formation,

V(3/Vj, varied from five to eight-tenths, respectively.

c. The reaction of 2 -met hoxye t hylamine and isobutyraldéhyde.

The 4o recorded determinations of were averaged and corrected Sb for protonated amine by using the average pH and the pEa of 9 - 09 «

The 21 recorded determinations for observed in solutions of

pH 1 1 .5 -10.5 were treated assuming that the hydration-dehydration io6 TABLE 21 Equilibrium Constants for the Reaction of 3-Methoxypropylamine and Isobutyraldéhyde

Run 4 ^T Kiobs ^ o b s IlEMUP] 240 [EMT] M M # 1-5 0.0183 0.0200 4 75.8-3.5 2.70^.49 0.973

6-10 0.0183 0.0200 5 7 2 .0 -1,6 2.09-0.5 0.961

11-15 0.0183 0.0200 3 60.3-2.1 1.48^0.51 0.867

16-19 0.0183 0.0200 6 63.7^.4 1.24^0.42 0.849

20-23 0.0183 0.0200 5 28.9-3.2® 1.82^0.21 0.432

24-28 0.0219 0.0300 6 42.1-2.1 1.91-0.10 0.580

29-32 0.0219 0.0400 6 32.8^0.3 2.68I0.IO 0.426

33-37 0.0219 0.0800 4 26.4^1.5^ 0 .306± 0 .15‘* 0.447

38-41 0.0219 0.0800 5 12.4-0.1® 0 .466± 0.04 0.198

Average = 74 *6-2 .6

Average = 3.04^0.77

— This entry not averaged TABLE 22 107 Rate Constants for the Reaction of 3-Methoxypropylamine and Isobutyraldéhyde

*^IMcali^^Ccal ^2 obs ^2Hobs Run pH -1 -1 240 # average sec sec

1-5 6 11.38+0.02 86.8 2.96 6.18+0.11 ——

6-10 7 11.22+0.02 85.7 2.92 6.19+0.45 —

11-15 6 10.64+0.1 77.4 2.64 6.44+0.14 —

16-19 6 10.58+0.06 75.5 2.85 5.86+0.16 —

20-23 5 9.71+0.12 38.6 1.31 14.62+0.55- 13.6+0.5-

24-28 6 9.97+0.13 51.8 1.76 8.06+0.56 7.20+0.50

29-32 6 9.70+0.122 38.0 1.29 7.55+0.68 6 .61+0.60

33-37 6 9.74+0.10 39.9 1.36 7.06+0.13 5.72+0.11

38-41 5 9.22+0.11 17.7 0.602 10.62+0. 38 9.07^0.33 -1 — Calculated with K_ — 92.3 M from ref. 5a Tj ■“ This entry not averaged 108 equilibrium, K^, for isobutyraldéhyde is established before ob­

servations are recorded. The pseudo first-order rate constant

for dehydration, at pH 10.5, 6.1 sec"^, affords an estimate of

a half-time of 0.11 sec. Thus observations for extrapo­

lated to zero time for data taken after about 0 .3 -0.4 sec, are based on measurements made after the third half-life of dehydra­ tion and it -was assumed that this equilibrium is established.

The remaining 19 recorded determinations of K^obs> observed in

solutions of pH 9-4-8 .6 , -were treated assuming that the change

in hydrate concentration ■with respect to time is zero, dH/dt =

0. Under these conditions the half-life for the dehydration is

greater than 1.5 seconds, and the observations for "were recorded before as much as l49^ of the hydration-dehydration

equilibrium could be established. Observed values of were

averaged and corrected for protonated amine. The average Kq and sa the literature value if Kj, 54.0 M , were used to obtain the

calculated constants, eq. 77 and j8, which were enployed in the

kinetic expression (see Tables 23 and 24).

The 152 recorded observations for determination of ksobs in

solutions of pH 11.5-10.5 were treated assuming the dehydration

equilibrium, Kjj, is established rapidly relative to establishment

of carbinolamine-imine equilibrium, K-j-. (The ratio of the rate

of dehydration, V^, to the rate of imine formation, Vj, is greater

than t-wenty-six.) The l44 recorded observations for determination

of ksobs In solutions of pH 9.4-8.6 were treated assuming the TABLE 23 Equilibrium Constants for the Reaction of 2-Methoxy ethyl amine and Isobutyraldéhyde K_ , lobs Cobs E m i p J Run 4 # [EMT] 225 M M

1-5 0.0267 0.0436 8 45.8+0.9 0.996

6-10 0.0267 0.0431 7 54.7± L .5 3.18+0.51 0.991

11-14 0.0267 0.0439 6 52.9± 2.4 2.86+0.63 0.966

15-19 0.0267 0.0600 7 29.1± 0.3 1.53+0.28 0.654

22-25 0.0267 0.1025 6 15.6+0.8 0.807±0.0118 0.415

26-28 0.0267 0.1743 6 8.71+0.13 0 .455± 0.051 0.250

Average = 53* 9 ^ - 1

Average = 2 -34^ 0,35 TABLE 24 Rate Constants for the Reaction of 2-Methoxyethylamine and Isobutyraldéhyde

Run pH ^IMcal^ ^Ccal ^ 2obs ^ 2Hobs Average 225 # M -1 M -1 sec”l sec-1

1-5 7 11.473-0.005 51.3 2.44 2.18-0.06 ——

1-10 6 11.113-0.001 51.1 2.43 2-17-0.15 “

11-14 6 10.544-0.04 49.8 2.37 2.24-0.14

15-19 6 9.366-0.076 33*7 1.60 2.60-0.43 2.23-0.40

22-25 6 8.940-0.09 21.4 1.02 3.37-0.59 2.81^0.50

26-28 6 8.611-0.08 12.9 0.614 4.90-0.29 4.03^0.24

— Calculated with — 54-0 M ^ ref. 5a Ill change in hydrate concentration with respect to time is zero, and also treated assuming the dehydration equilibrim, Kp, is estab­ lished rapidly relative to establishment of carbinolamine-imine equilibrium. These two calculated rate constants, k^obs

^2obs) averaged, for the ratio of the rates of dehydration to imine formation, Vg/Vp, varied from two to one-fourth.

d. The reaction of 2,2-dimethoxyethylamine and isobutyral­ déhyde . The 56 recorded determinations of were averaged and corrected for protonated amine by using the average pH and 5b the pKa of 8.35 • The 27 recorded determination of ob­ served in solutions of pH 11.3 to 10.0 were treated assuming that the hydration-dehydration equilibrium, Kp, for the aldehyde is established before observations are recorded. The pseudo first- order rate constant for dehydration of the aldehyde hydrate at pH 10.0, 1.92 sec affords an estimate of a half time of O .36 sec. Thus observations for Kg^pg, extrapolated to zero time from data taken after about 0.4-0.8 sec, occur largely after the second half-life of dehydration and it was assumed that this equilibrium is established. The remaining 29 recorded determin­ ations of KgQ'j^g) observed in solutions of pH 9*6 to 7-8, were treated assuming that the change in hydrate concentration with respect to time is zero, dH/dt = 0. Under these conditions

(pH = 8 .6 ) the half time for dehydration is approximately four seconds, and the observations for recorded before as 112 much as 5^ of the hydration-déhydration equilibrium can be es­

tablished. Observed values of were averaged and corrected

for protonated amine. This average Kq and the literature value sa of Kj, 50.85 M , were used to obtain the calculated constants,

eq. T7 and 7 8 , which are employed in the kinetic expression (see

Tables 25 and 26).

The 183 recorded observations for determination of ksobs

solution of pH 11.5-10.0 were treated assuming the dehydration

equilibrium, Kp, is established rapidly relative to establishment

of carbinolamine-imine equilibrium, Kj. At pH 10.0 the ratio of

the rate of dehydration, V^, to the rate of imine formation, Vj,

is greater than thirty.

The 2hh recorded observations for determination of ksobs

solution of pH 8.6-7.8 were treated assuming the change in hydrate

concentration with respect to time is zero, and also assuming the

dehydration equilibrium, Kp, is established rapidly relative to

establishment of carbinolamine-imine equilibrium. These two

calculated rate constants, kgobs ^^Hobs averaged. This

calculation was necessitated for the ratio of the rates of dehy­

dration to imine formation, Vg/Vp, varied from two to one-fifth.

e. The reaction of 2,2,2-trifluoroethylamine and isobuty­

raldéhyde . The 25 recorded determinations of were averaged

and corrected for protonated amine with the average pH and the pKa 5b of 5 «52. The nine recorded determinations of Kp^i^g observed in 113 TABLE 25

Equilibrium Constants for the Reaction of .2,2-Dimethoxyethylamine and Isobutyraldéhyde

4 “ t ^lobs Run M 260 M #

1-3 0,0219 0.0400 6 26.8±0.5 1.33*0.42 0.999

4-7 0.0219 0.0400 7 27.0±0.7 1 .08± 0.27 0.998

8-11 0.0219 0.0400 7 25.3-0.5 0.884±0.252 0.982

12-15 0.0219 0.0400 7 27.3-0.3 1.46^0.42 0.980

19-21 0.0219 0.0400 7 17.5-1.3 0.550*0.126 0.667

15-18 0.0219 ,0.0600 4 11.3- 0 .2 0.457*0.074 0.508

23-27 0.0219 0.0800 5 10.9-0.3 0.481*0.155 0.508

28-31 0.0219 0.120 7 11.3-0.1 0.325*0.063 0.562

32-35 0.0219 0.160 6 5.38=^0.16 0.158*0.046 0.247

Average Average = 1 .07-0.17 TABLE 26

Rate Constants for the Reaction of 2,2-Dimethoxyethylaniine and Isobutyraldéhyde

Run pH ^IMcal^ ^Ccal K2obs V. ^ 2Hobs 260 # average sec“^ sec"^

1-3 6 11.288-0.02 27.3 1.07 1.01-0.08 —

4-7 6 11.12-0.01 27.3 1.07 1.02-0.07 ——

8-11 6 10.08^0.06 26.9 1.05 1.02-0.08 ——

12-15 6 10.03-0.04 26.8 1.05 1.08-0.06 —

19-21 6 8 .65±0.05 18.2 0.714 2.36-0.24 2.08-0.21

15-18 6 8.36-0.04 13.9 0.544 4.43-0.27 4.08-0.25

23-27 6 8.364-0.04 13.9 0.544 4.34-0.36 3.78-0.31

28-31 6 8.458-0.03 15.4 0.601 4.11-0.13 3.43-0.11

32-35 4 7.87-0.05 6.76 0.264 IL 32-1.01 9.75-0.87

i Calculated with Kj = 30. 85 ref. 5a 115

solutions of pH 12 to 11 were treated assuming that the hydration-

dehydration equilibrium, Kp, for isobutyraldéhyde is established

before observations are recorded. The pseudo first-order rate

constant for dehydration, 19-2 sec“^, at pH 11, affords an esti­ mate of a half time of O.056 sec. Thus, observations for

extrapolated to zero time from data taken after about 0 .5 -1.0 sec,

occur after the 15th half-life of dehydration and it was assumed

that this equilibrium is established. The remaining l4 recorded

determinations of K^obs’ observed in solutions of pH 6.5 to 5*5,

were treated assuming that the change in hydrate concentration

with respect to change in time is zero, dH/dt = 0. Under these

conditions, the half-time for dehydration is approximately 60

seconds and observations for were recorded before as much as

3^ of the hydration-dehydration equilibrium can be established.

Observed values of K^obs were averaged and corrected for protonated 5a _ amine. This average Kq and the literature value of Kj, 6.29 M ,

were used to obtain the calculated constants, eq. 77 and 78 , which

are employed in the kinetic expression ( see Tables 27 and 28).

The 88 recorded observations for determination of ksQ-^g in

solution of pH 12-11 were treated assuming the hydration-dehydration

equilibrium. Kg, is established rapidly relative to the establish­

ment of the carbinolamine-imine equilibrium, Kj. At pH 11 the

ratio of rate of dehydration, V^, to the rate of imine formation,

Vj, is greater than one thousand. The 128 recorded observations

for determination of kgQ-bg in solutions of pH = 6.5 were treated TABLE 27 n 6 Equilibrium Constants for the Reaction of 2, 2, 2--Trifluoroethylamine and Isobutyraldéhyde

EMIP, Run ^lobs OODS [EMTJ 230 M M # M“^

1-4 0.0297 0.3815 4 5 . 4 0 ^ . 0 1 0.220^0.01 1.000

5-7 0.0297 0.3815 5 5 .01± 0.34 0.187-0.04 1.000

8-11 0.0297 0.3815 6 5.65-0.13 0.157-0.08 0.906

13-17 0.0297 0.3815 3 4.49-0.01 0.153-0.025 0.731

21-23 0.0297 0.3590 5 3.17-0.06 0.194-0.038 0.519

Average = 5 .36-O.48 Average = 0 .238^0.036 TABLE 28 Rate Constants for the Reaction of 2j2,2-Trifluoroethylam±ne and Isobutyraldéhyde

Run pH - ^IMcal ^ ^Ccal 2obs ^2Hohs 230 # average —1 -1 M-' sec sec

1-4 5 12.339-0.03 6.05 0.238 0.0208^0.0014 ——*

5-7 6 10.925^0.041 6.05 0.238 0.0224-0.0015 ——

8-11 6 6 .502± 0.07 5.48 0.216 — 36.(Æl .4

13-17 6 5.954+0.002 4.42 0.174 — 137.0±2.2

21-23 4 5.553+0.007 3.14 0.124 —— 358.0+9.1

a Calculated with K^=6 .29 m"^ ref. 5a 118

assuming the change in hydrate concentration with respect to time

is zero. The ratio of the rate of dehydration to imine formation,

Vd/Vj, is less than one one-thousandth.

Equilibrium constants for carbinolamine (Kq) and carbinol­

amine-imine (Kj) formation for six monofunctional primary alkyl-

amines are given in Table 2 9 . The literature value fof and

observed values of Kq are plotted against the acidity constants

of the respective alkylammonium ions, pKa and the Taft substituent

constants. Figures 21-2$. Results of a least squares treatment of

these linear functions, of the form y = mx + d, are listed. Table

5 0 .

The observed rate constants for these primary alkylamines are

evaluated by assuming that the formation of imine is subject to

only specific acid catalysis, eq. 79, and subject to general acid

catalysis, eq. 8 0 . To determine the general acid catalysis rate

kobs = ko + kjj [H+] (79)

kots = ko + kjj [H+] + kBH [BH+] (8o )

constant, the concentration of the protonated amine, must be known.

Although buffers other than the reacting primary amines were not

used, the concentration of protonated amine, that is the concen­

tration of the general acid, probably does not vary significantly;

for it is the free amine that reacts. Most reactions were experi­

mentally observed for the first two half-lives; and the concentra­

tion of the protonated amine at the first half-life, which is not TABLE 29

Equilibrium Constants for Formation of Garbinolamines and Imines from Isobutyr­ aldéhyde and Monoamines in Aqueous Solution at 35° log Amine pKa- Code Kja ic" a " ÿ sec n-Propyl 10.21 107.3 1 102.4 2.031 5.09 0.707 -0.115

Methyl 10.31 127.9 2 128.1 2.107 8 .5o£ 0.942 0.0

3“Methoxy- propyl 9.83 92.3 3 76.3 1.966 3.04 0.483 0.08

2-Methoxyethyl 9.09 54.0 4 53.9 1.732 2.45 0.389 0.23

2 f2-D imethoxy- ethyl 8.35 30.9 5 26.7 1.490 1.07 0.029 0.46

2 f2 ,2-Trifluoro- ethyl 5-22 6.29 6 5.36 0.799 0.238 -0.624 0.92

— Values from ref. 5 * “ Values from this investigation. — Values from ref. 7» 120

1). (3)

1.7 W

M (5)

1.3

(6 ) 0.8

10 pKa

Figure 21. Plot of pKa vs. log for Monoamines 121

1.0

(3 ) (1) 0.6

(4 ) I 0.0 (5)

—0.5 slope = 0 . 2 8 (6)

10 pKa

Figure 22. Plot of pKa vs, log Kq for Monoamines 122

1.0

0.5

0.0

-0.5

0.0 0.2 0.4 0.6 0.8

Figure 2$. Plot of a* vs. log Kq for Monoamines. 125 TABLE 30

Evaluation of Linear Functions of the Form y = mx + d

Plot Slope Intercept Figure y/x m d

21® Kj/pKa 0.254 — 0.56

22® Kç/pKa 0.28 -0.22

23® V 0 -* -1.42 0.69

24^ k y p K a 0.51 — 0.04

25® - 2 . 5 7 0.88 0.81 25^ V O ’* -1.79 27® k^Kc/pKa 0.793 6.45

— All podLnts of egch figure are used in the least squares calculation. — Point number six (2^2^2-trifluoro­ ethylamine) was omitted for this result. 124

significantly different from the initial protonated amine con­

centration, ■was used in solving eq. 8o. The calculated constants are listed in Table 31* The investigation of imine formation with n-propylamine was conducted in aqueous solutions of a limited pH range, 11.4-11.5, in which less than eight percent of the amine existed in the protonated form. These eijperimental conditions do not permit e-valuation of general acid catalysis. Although the in­ vestigation of imine formation with 2 ,2,2-trifluoroethylamine was

conducted in aqueous solutions of a moderately large pH range,

12.0 -5 .5 , little protonated amine (pKa = 5 *22) existed in the

solutions which were more basic than pH 6 .6 . The observed rate

constants in solutions of pH range 6.5 to 5*5 are large because of the expected contribution from specific acid catalysis. The eval­

uation of general acid catalysis, eq. 80, yielded a negative rate

constant. A least squares trea'fcment of the observed rate constants

for imine formation from 3 -methoxypropylamine with eq. 80 also yielded a negative rate constant, which indicates the absence of

general acid catalysis. Also, the observed rate constants for

imine formation from 2 ,2 -dimethoxyethylamine remained unchanged

when the amine concentration was doubled at constant pH of about

8.4. The observed rate data for 2-methoxyethylamine and 2,2-

dimethoxyethylamine is not correlated with a significantly smaller

fractional standard deviation with eq. 80 than with eq. 7 9 . Fur- 97 thermore, Hine and coworkers did not observe general acid TABLE 31 125 Rate Constants for Monoamines k Amine Code 0 -1 STD sec“^ M sec M"^ xec"^ 1 9.763

2 6.20 3 .2x10^

MeO(CH^)g- 3 6.13 -43.2 1.04xloJ° 0.036 6.06 5 .75x 10^ 0.068

MeOCCHg)^ 4 2.16 9.04 3.30x10® 0.068 2.17

(MeO)^CHCHg- 5 0.991 5.52 6 .43x10? 0.045 0.991 7 .00x10 0.067

F3CCH,- 6 0.0209 -264.0 1 .45x10? 0.019 0.0207 1 .21x10 0.046

— Value from ref. 7. 126 catalysis for the hydrolysis of ïï-isob-utylidenemethylamine.

Therefore, the observed rate constants for the six monofunctional amines, which are used to determine stnictural affects on reac­ tivity of the primary amines in which internal catalysis is not possible, are evaluated with eq. T9, which assumes no general acid catalysis. The calculated rate constants for the uncatalyzed, ko, and the specific acid catalyzed, kg, reaction are plotted a- gainst the acidity constants of the respective amines, pKa, and the Taft substituent constants, Figures 2h to 26. The specific acid catalyzed rate constant for 3 -methoxypropylamine determined mostly from the kinetic data for one set of conditions is consid­ ered unreliable and is not illustrated. Figure 26. Results of a least squares treatment of these linear functions, of the form y = mx + d are listed in Table 30. The first order, uncatalyzed

dehydration of carbinolamine to yield imine can be written as a

second-order reaction in terms, of the isobutyraldéhyde and amine

concentrations. A plot of this second-order rate constant, k^Kg,

Rate Forward = koC = k^K^AM = kosAM (8l)

against the acidity constants of the respective amines yields a

linear function. Figure 2J.

97. J. Hine, J. C. Craig, Jr., J. G. Underwood, II, and P. A. Via, J. Amer. Chem. Soc., in press. 127

1.0

0.5

0.0

slope = 0.51

- 1.0

-1.5

6 7 8 9 10 PKa

Figure 2 h . Plot of pKa vs. log kg for Monoamines. 128

1.0

slope = -1.79

0.0

slope = -2.57

0.00 0.25 0.50 G.75

Figure 25. Plot of a* vs* log k© for Monoamines. 6 7 8 9 pKa

Figure 26. Plot of pKa vs. log kg for Monofunctional Amines 2.0

1.0

0.0

1.0

pKa

Figure 27# Plot of log vs. pKa for Monoamines

o 351

5. Diamines

Imine formation from isotutyraldehyde and four œ-dimethyl- amino-alkylamines was studied. Equilibrium constants for car­ binolamine formation with free (Kq ) and with monoprotonated diamines (KQp), the equilibrium constant for carbinolamine -imine formation with free (Kj) and monoprotonated diamines (Kjp), and the rate constants for dehydration of carbinolamine were deter­ mined. The equilibrium constants for the diamines were calculated from eq. 82 and 83 using a least squares method for determining a

%obs = (82)

Koobs = (83) minimum fractional standard deviation. The standard deviation is calculated from the sum of the squares of the deviations, eq. 84.

These constants and rate data are listed in Tables 32-43 •

_ Ce m t o ] _ [emlp] '^p_ deviation = (%) ten] [EM]

a . The reaction of N,ïï-dimethylethylenediamine and iso­ butyraldéhyde . Equilibrium constants for carbinolamine-imine formation; The 66 recorded determinations of were used to calculate the average Kp and Kpp by employing eq. 82 , the final 98 pH, the pKai of 9.448 pKag of 6.446, (Table 32). A similar calculation using the calculated vales of Kj and Kpp yields the TABLE 32 Equilibrium Constants for Imine Formation with N, N-Dlmethylethylenedlamlne

At K_ . K_ [e m u p ] [e m i p ] Runs pH lobs lave # MM at tfo M M 24-1

210.78-81 5 0.02663 0.0409 11.202 0.04019 0.00071 49.96*0.47 48.28 210.74-77 6 0.02663 0.0409 10.718 0.03882 0.00208 48.04-0.64 47.15 131.27-30 a 0.0290 0.0449 9.33 29.76 31.00 210.67-73 7 0.02663 0.0513 8.634 0.0068 0.0442 18 .81±0.72 19.79 210.44-48 3 0.02663 0.0460 8.518 0.0048 0.04085 i7.56io.29 18.82 210.49-55 9 0.02663 0.0450 7.642 0.00065 0.04169 i4.78io.23 15.03 210.56-61 6 0.02663 0.0213 7.574 0.00026 0.01958 13 *99^0.68 14*81 210.40-43 5 0.02663 0.0425 7.368 0.00019 0.03697 l3.5io.5i 14,07 210.34-39 8 0.02663 0.0444 7.154 0.00031 0.03768 13.34^0.44 13.08 210.28-33 8 0.02663 0.069 6.406 0.00003 0.0329 8.18—0.76 7.39 210.24-27 9 0.02663 0.133 6.152 0.00002 0.0448 6 . 4 3 ^ . 0 7 5.22 170.23-25 a 0.0291 0.1738 5.984 5.22 3.08

— Not used to determine average Average K_ =48. 86 Average K „ =» 15*46 STD = 0.078 5.000E 0 1 - ♦ ^ * # # * cii:nuüi.ij()CjuJ*jijoüuiCüf/i # oocccu OOfl - cc I) .0 0 u 4.000k 01 - * ♦ ♦ ..... 0 ♦* . ♦ 0 0

— _ ' 0 _ o - u 3.000E 01 - ♦ ♦ 4 ♦ ♦ — ^ 0 ______0 0

K lobs : “ 2.000E 01 - ♦ ♦ 0 ♦ ♦ ♦ oil# - u • - ______ÜOU__ _ _ _ 1 ■ ' ' " ÔCO " ' ■ " ■ t c e# , 0# # 0# CO - __ 0 _ ___ _ l.OOOE 01 -...... * " ' d * ...... ♦ * ' " ♦ : Ü •00 0 _ _ _ - U ...... CO - cco ucou 0.0 11111rICQUCCCGCCÛQ01111111111111111111111111111111111II111 1111111II11 11111111 II11111111111111 11111111

\PNP- 0 I.OOUE CO S. 140k CO T.gdOk 00 9.4>üé 00 " l - l W 01 ' 1.3701 01 pH tl Figure 28. A Plot of pH vs. for MesN(CHa)sNHg. 134 average observed equilibrium constants, Kj^yg. A plot of the observed equilibrium constants vs^ pH is shown in Figure 28.

Equilibrium constants for carbinolamine formation; The eleven recorded determinations of Table 53} observed in solutions of pH 11.3 to 10.8 were treated assuming that the hy­ dration-dehydration equilibrium, Kp, for isobutyraldéhyde is established before observations are recorded. The pseudo first- order rate constant for dehydration, at pH 10.8, 12.5 sec"^ af­ fords an estimate of a half time of 0.06 sec. Thus, observations for extrapolated to zero time from data taken after about

0 .2 -0.3 sec occur after the third half-life of dehydration and it was assumed that this equilibrium is established. The remaining

55 recorded determinations of observed in solutions of pH

9 .4-6.0, were treated assuming that the change in hydrate concen­ tration with respect to change in time is zero, dH/dt =0. In this pH range the half time for hydration is greater than 1.5 sec and observations for were recorded before as much as 6^ of the hydration-dehydration equilibrium can be established. All values of K^obs used to calculate the average Kg and Kçp by employing eq. 83 and the initial pH. The reverse calculations using Kq and K^p yields the average observed equilibrium constants,

^ave* The observed equilibrium constant, is plotted against pH, Figure 29 .

98. J. H. Jensen, unpublished observations at zero ionic strength. TABLE 33 Equilibrium Constants for Carbinolamine Formation with N,N-Dimethylethylenediamine

Runs # 4 pH Ce m u p ] [e m i p ] ^C obs Cave MM at to M M M-1 M-1

210.78-81 5 0.02663 0.0409 11.259 0.04028 0.00062 2.02-0.31 2.03 210.74-77 6 0.02663 0.0409 10.790 0.03912 0.00178 2.21± 0.57 2.12 131.27-30 a 0.0290 0.0449 9.43 3.45-0.51 3.69 210.67-73 7 0.02663 0.0513 8.869 0.01071 0.04051 4 .41± 0 .é5 4.53 210.44-48 3 0.02663 0.0460 8.710 0.0071 0.0387 4.67^0.21 4.94 210.49-55 9 0.02663 0.0450 7.836 0.0010 0.0422 4.59-0.65 4.69 210.56-61 6 0.02663 0.0213 7.809 0.00046 0.0200 4.12-0.25 4.76 210.40-43 5 0.02663 0.0425 7.500 0.00044 0.0386 4.85^0.15 4.94 210.34-39 8 0.02663 0.0444 7.256 0.00025 0.0382 4 .95^ 0.43 4.50 210.28-33 8 0.02663 0.0690 6.462 0.00004 0.0351 3.62-0.61 2.66 210.24-27 9 0.02663 0.1330 6.232 0.00003 0.0504 2 .08^0.47 1.98 170.23-25 a 0.0291 0.1783 6.024 1.53 1.38

— This entry not averaged Average 1 . 9 7 6 Average K^p 5.217 STD = 0.11

« VI 6.000k 00

^ • • _ # occuouo______4.800E 00 - * ♦ OU ÜÜ ♦ 0 • G • 0 0 • 0 0 _ _ • _ 0 _____ - ■ ' Ü. Tj

0 0 u 3.600E 00 - * • O « *

' - .. ' " 0 ~ Ô

0 0 0 o - _ _ _ 0 ______% obs ’ - ' ' """ " Ô 0 00 2.400E 00 - * » » • Ü GO 0 00# • ------OGL^OQ’jguCAOO..... 0 ! ÜÛOürJCtOfiCOGü

0 ■ . I.200C 00 —: ♦ 0 ♦ ♦ ♦ ♦ 0

- 0 CO u ocu coco q.SlTE'UT 111111ICCUOCtCOI1111II11111111111111111II111111111111111111111111 111II1111111111111111IIIII1111111111

\PNPm 0 j.OOOt 00 * "i.UOE CO - t .z h OE'00 ' ' O.VzOE 00 ’ i.isoC 01 l.îfOt 01 pH H Figure 29 . A Plot of pH vs. for M62N(CH2)2NH2 . 157

Rate constants for dehydration of the carbinolamine; The equilibrivim constants for carbinolamine formation, used

in the kinetic expression were calculated with expression 83, the average pH, Kq , and K^p. The equilibrium constants for imine

formation, ) used in the kinetic expression were calculated with expression 85, the average pH, Kj, and Kjp, Table 3^* The

KlMcal = ^ ^ *IP ■ Kccal

l44 recorded observations for determination of kg in solutions of

pH 11.8-10.7 were treated assuming that the hydration-dehydration

equilibrium, Kp, is established rapidly relative to the establish­ ment of the carbinolamine-imine equilibrium, Kj. At pH 10.7 the

ratio of the rate of dehydration, V^, to the rate of imine forma­

tion, Vj, is greater than one hundred. The 48 recorded observations

■for determination of kg in the solution of pH 9.98 were treated

assuming that the change in hydrate concentration with respect to

change in time is zero, and also treated assuming that the dehy­

dration equilibrium is established rapidly relative to establish­

ment of carbinolamine-imine equilibrium. These two calculated

rate constants, ksobs ^nd kz^obs» w^re averaged, for the ratio of

the rate of dehydration to imine formation, Vg/Vp, is one-half.

The 452 recorded observations for the determination of kg in sol­

utions of pH 9 *^-8.0 were treated assuming that the change in

hydrate concentration with respect to change in time is zero.

The ratios of the rate of dehydration to imine formation, Vg/Vp, TABLE 34 Rate Constants for Imine Formation with N, N-Dlmethylethylenediamine k Runs # 4 “ x pH *%Mcal ^Ccal 2obs ^ 2Hobs M M average M-1 M -1 sec""^ sec

210.82-85 5 0.02663 0.0454 11.745^0.01 46.69 1.99 4.09-0.10 210.78-81 7 0.02663 0.0409 11.23*0.03 46.28 2.03 6.85+0.22 210.74-77 6 0.02663 0.0409 10.75-0.04 45.16 2.13 13.6±1.6 210.87-90 6 0.0298 0.0789 9.978+0.06 38.53 2.74 50.oil.4 40 .9ii.o 131.27-30 « 6 0.0290 0.0449 9.38^0.05 27.52 3.69 37.5*1.5 210.67-73 5 0.02663 0.0513 8.751Ï0.11 16.30 4.66 33.7*4.3 210.44-48 5 0.02663 0.0460 8.61+0.10 14.84 4.77 35.1*1.6 210.49-55 5 0.02663 0.0450 7 .739Ï 0.1 10.43 4.91 24.9*1.3 210.56-61 7 0.02663 0.0213 7.692^0.12 10.30 4.89 23.4*1.7 210.40-43 4 0.02663 0.0425 7 .434±q.o66 9.55 4.71 30.9*1.4 210.34-39 6 0.02663 0.0444 7.205^0.051 8.83 4.43 24.7*3.3 210.28-33 4 0.02663 0.0690 6.426+0.05 5.24 2.66 28.8*2.3 210.24-27 6 0.02663 0.133 6.19-0.07 3-67 1.87 23.9*1.4 170.23-25 a 6 0.0291 0.1738 6.OO4I0.O2 2.70 1.38 25.6*3.5

CD TABLE 34 Continued

Runs . # 4 pH *IMcal ^Ccal 2obs. *^2Hobs MM average M-1 M-1 sec-1 sec-1

210.82-85 5 0.02663 0.0454 11.746+0.01 46.71 1.972 4.69+0.11 210.78-81 7 0.02663 0.0409 11.23+0.03 46. 36 1.953 7.11+0.22 210.74-77 6 0.02663 0.0409 IO.754+.04 45.26 1.9 01 14.8+1.6 210.87-90 6 0.0298 0.0789 9.978+0.06 39.64 1.600 104.±3. 9 4 .±3 131.27-30^ 6 0.0290 0.0449 9.38+0.05 30.13 1.08 128.+13 210.67-73 5 0.02663 0.0513 8.751+0.11 20. 37 0.589 271.+36 210.44-48 5 0.02663 0.0460 8.61+0.10 19.09 0.523 338.±16 210.49-55 5 0.02663 0.0450 7.739+0.1 15.01 0.329 434.±22 210.56-61 7 0.02663 0.0213 7.692+0.12 14. 87 0.324 440.±32 210.40-43 4 0.02663 0.0425 7 .434±.066 13.96 0.299 603.±50 210.34-39 6 0.02663 0.0444 7.205+0.051 12.98 0.275 480.±6 5 210.28-33 4 0.02663 0.0690 6.426+.05 7.74 0.154 654.±98 210.24-27 6 0.02663 0.133 6.19+0.07 5.43 0.112 474.±27 170.23-25- 6 0.0291 0.1738 6.004+.02 4.00 0.083 622.±87 a This entry not averaged. — This rate data was calculated using an estimated K of 0.313. The value of K^p was obtained from eq. 112 using pKa^i ~ 8.65, P^®nit ~ 9 *031, 98 and [EMlPt]/[EMlP] = O.38. ^ Kc “ 1 VO 140 were less than one-ninth.

An alternate treatment was employed for evolution of this rate data for N,lT-dimethylethylenediamine, Table $4 , continued.

The experimentally determined value for K^p (5 *22), as will be considered later, is unusually large and an estimated value

(0.515) was employed. This estimated was determined with [EMlPt] ^ ^ ^ . eq. 112, = 0.58, Kq - 1.98, pKaQ2 = 8.65 and pKajyQ_.j.

- 9 -051' The value for the first ionization constant for this carbinolamine was obtained from Table 48.

b . The reaction of N-dimethyl-1 ,3 -propanediamine and isobutyraldéhyde. Equilibrium constants for carbinolamine-imine formation; The . 64 recorded determinations of KjQ^g were used to calculate the average Kj and Kjp by employing eq. 82, the final - 98 pH, the pKai of 10.05, and the pKag of 8 .00, Table 5 5 - A similar calculation using the calculated values of and K^p yields the average observed equilibrium constants. The observed constants,

Kjobss plotted against pH, Figure 30 .

Equilibrium constants for carbinolamine formation: The 25 recorded determinations of observed in solutions of pH 11.5-

10.1 were treated assuming that the hydration-dehydration equili­ brium, Kg, for isobutyraldéhyde is established before observations are recorded. The pseudo first-order rate constant for dehydration at pH 10.1, 2.5 sec“^, affords an estimate of a half time of 0.27 sec. Thus, observations for extrapolated to zero time from data taken after about 0 .4 -0 .6 sec occur after the second half-life TABLE 35 , Equilibrium Constants for Imine Formation with N,N-Dimethyl-l,3-propanediamine

Run pH [e m u p ] [e m i p ] # 4 «I lobs lave 270 M M at t(a M M 1-5 6 0.0267 0.040 11.360 0.0381 0.0019 64.3-0.7 98.58 6-8 3 0.0326 0.040 10.866 0.0347 0.0053 109±0.6 91.49 9-12 5 0.0326 0.040 10.333 0.0263 0.0137 91.9 -1.9 74.08 13-16 5 0.0326 0.040 10.012 0.0190 0.0208 76.0-5.0 59.13 17-20 6 0^0326 0.040 9.731 0.0128 0.0267 60.9*1.3 46.16 21-24 5 0.0326 0.040 9.458 0.0079 0.0310 47.1*0.94 35.82 25-28 6 0.0326 0.020 9.011 0.0015 0.0168 20.78*1.0 24.72 43-46 4 0.0326 0.040 9.086 0.0037 0.0336 20.4-0.6 26.16 47-50 3 0.0326 0.040 9.017 0.0031 0.0336 21.5*1.6 24.83 29-32 5 0.0326 0.048 8.458 0.0009 0.0349 9.81*0.87 16.49 33-36 5 0.0326 0.056 8.208 0.0005 0.0343 7.98*0.38 13.16 37-39 6 0.0326 0.072 7.978 0.0003 0.0349 5.87-0.62 10.14 40-42 5 0.0326 0.104 7.646 0.0001 0.0319 3.62*0.15 6.26

Average 102.4 Average K 20.0 STD 0.25 l.SüOf- oi -

l.ZüPl 0/ -

iiDijninnuii'jrjuuuiUiuuua - , uua . ÜO — # 00 9.00UC 01- ♦ ♦ * 0 * ♦ " U '...... • o 0 ...... # K u

- ... _ 0 # 6.30PE 0 1 - * * * # 0 *

• 0 00 V ü“ 3.U00E - ♦ ♦ 00* * «• - - - - , uo , “ _ „ 00 uu UO ■ ' 0$ " ■ " ...... - ...... w CCO QCCOO 0,0 11 II 11 ! tüurciiutuoucucocourtOocounr.LGOui ii ii ii fi il ii ii i ii u 11 ii 11 ii 11111 ii iii i ii 11 il ii ii n t n ii i ii 111 ii NPNP« O 3^Ô6(,L OU f.ZHÔE 00 9.420E 00 ûlVbE 01 ’l.370É 01 Ph g- Figure 50. A Plot of pH vs. Kt for M e g N ( C H g ) . ■^obs 1^5 of dehydration, and it "was assumed that this equilihritmi is es­ tablished. The 5^ recorded determinations ^Cobs observed in solutions of pH 9 *5-7 «8 were treated assuming that the change in hydrate concentration with respect to change in time is zero, dH/dt = 0 . Under these conditions, pH 9 *5 > the half time for dehydration is greater than 1.8 sec and observations for were recorded before as much as ll^ of the hydration-dehydration equilibrium can be established. For convenience, the five re­ corded determinations of in the solution of pH 9 «67 were treated assuming that the hydration-dehydration equilibrium is established. Estimates indicate that this equilibrium is only established at the time the spectrophotometric observations were recorded. All values of were used to calculate the average

Kg and Kgp by employing eq, 85 and the initial pH, Table $6. The reverse calculation using Kq and Kgp yields the average observed equilibrium constant, K^ave’ which is plotted, with Kg^^^g against pH, Figure 5 1 »

Rate constants for dehydration of the carbinolamine; The equilibrium constants, Kg^gi ^IMcal? in the kinetic ex­ pression were calculated as previously described, eq. 85, Table 57 -

The 120 recorded observations for determination of kg in

solutions of pH 11.4-10.4 were treated assuming that the hydration-

dehydration equilibrium, Kp, is established rapidly relative to the

establishment of the carbinolamine -imine equilibrium, Kj. The ratio

of the rate of dehydration, V^, to the rate of imine formation, Vj, TABLE 36 Equilibrium Constants for Carbinolamine Formation with N,N-Dimethyl-'l, 3—propanediamine pH K m u p ] Le m i f ] Run # at to ^Cobs *Cave 270 M M M-1 1- 5^ 6 11.512 0.0387 0.0013 1.01 6-8 3 10.924 0.0353 0.0047 O.85IÏO.I40 0.985 9-12 5 10.514 0.0298 0.0102 0 .811± 0.167 0.947 13-16 5 10.204 0.0235 0.0165 0.847-0.290 0.902 17-20 6 10.096 0.0210 0.0189 1.22±0.232 0.885 21-24 5 9.672 0.0116 0.0278 1.57-0-20 0.813 25-28 6 9.239 0.0015 0.0168 1.97+0.30 0.749 43—46 4 9.292 0.0057 0.0326 0.789^0.609 0.753 47-50 3 9.260 0.0053 0.0329 0.876±0.080 0.747 29-32 5 8.582 0.0013 0.0370 0 .576±o .o 66 0.602 33-36 5 8.364 0.0008 0.0385 0.528^0.080 0.527 37-39 6 8.093 0.0004 0.0396 0.335-0.036 0.416 40-42 5 7.757 0.0002 0.0378 0 .242±0.007 0.272 Average = 1.017 — Entry not averaged Average K^p = 0.745 STD = 0. 24 £ I.soot

• 1.200c UO - * » «

qnoui'o ou UÜU < I.OOOO'Ul - ♦ ♦ ♦ DO o • uu • #uo 00 0 KC,Ob 8 _ - - - Q — 0 b.OOOE-Ul - ♦ ♦ U ♦ u# #

u D 0 u 3.00CE-UI - ♦ * 1 1 ♦ ♦ « • 0 0 ' 000 CO nooon o.c mill! ccorccccccccuccccnouncuooi m 11111 ii 11 m i ii m 11 ii m 11 m i ii 1111 ii 11 m ii 111 m 111 ii 11 m 111 O I.OÙuE 00 b.KOE 00 7.?R0E 00 0.42VE 00 l.lïoE 01 I.370E 111 pH

Figure 31* A Plot of pH vs. for Me2N(CH2 )3 NH2 . TABLE 37 Rate Constants for Imine Formation with N,N-Dimethyl~ 3-jpropane diamine ^ k Run ^IMcal ^Ccal 2obs ^ 2Hobs pH average 270 # M -1 M -1 sec-1 sec-1 1-5 5 ll.436lo.076 98.17 1.00 2 0 .2- 0.2 6-8 6 IO.895I0.O30 91.13 0.983 23.3I0.4 9-12 4 IO.423I0.O9O 76.88 0.935 24.2I0.5 13-16 6 I0.i08l0.096 62.86 0.887 27.2I0.8 24.0I0.7 17-20 5 9.913I0.I8O 53.60 0.796 3 1 .8± L .0 28.2I0.9 21-24 5 9.568I0.IO7 38.74 0.776 33.9I2.3 3O.4I2.O 25-28 6 9.I25I0.II4 26.25 0.725 48.4I7.6 46.5I7.3 43-46 5 9.I89I0.IO3 27.65 0.736 42.2I5 .3 38.4I4.8 47-50 4 9.I43I0.I26 26.63 0.728 46.7I1.I 38.9I0.9 29-32 5 8.520I0.060 16.73 0.582 4I.2I2.4 37.7I2.2 33-39 4 8.266I0.056 13.44 0.489 4 2 .9I2.4 39.0I2.2 37-39 4 8.O36I0.OI7 10.50 0.391 44.0I4.I 39.6I3.6 40-42 5 7.701I0.056 6.59 0.250 44.7I5.4 40.1I4.8

— All solutions contain enough NaCl to give |i ™ 0*17

o\ 147 is greater than five. The 592 recorded observations for deter­ mination of kg in solutions of pH 10.1-7.7 "were treated assuming that the change in hydrate concentration with respect to change in time is zero, and also treated assuming that the dehydration equilibrium is established rapidly relative to establishment of carbinolamine-imine equilibrium. These two rate constants, I^obs and were averaged. From pH 10.1 to 7-7 the ratio of the rate of dehydration to imine formation, Vg/V%3 varies between one and two.

c. The reaction of H ,H-dimethyl-l ,4 -butane diamine and iso­ butyraldéhyde. Equilibrium constants for carbinolamine-imine formation: The 77 recorded determinations of E^obs were used to

calculate the average Kj and K^-p by employing eq. 82, the final pH, 98 the pKai of 10.287, and the pKag of 8.866, Table 58. A similar

calculation using calculated values of Kj and Kjp yields the av­

erage observed equilibrium constants, which is plotted with

^obs pH, Figure 52.

Equilibrium constants for carbinolamine formation: The 25

recorded determinations of observed in solution of pH 11.5 -

10.5 were treated assuming that the hydration-dehydration equili­

brium, Kjj, for isobutyraldéhyde is established before observations

are recorded. The pseudo first-order rate constant for dehydration

at pH 10.5, 3-8 sec"^, affords an estimate of a half time of O.18

sec. Thus, observations for K^Q^gj extrapolated to zero time from

data taken after about 0 .2-0.4 sec, occur after the second half-life TABLE 38 Equilibrium Constants for Imine Formation with N^N-Dimethyl-l,4-t*utanediamine

pH [EMUP] [EMIP] K_ Run # ^lobs • lave 310 M M at tgo M M M-1 M-1

1-4 7 0.0258 0.04005 11.515 0.0378 0.0022 96.9+1.5 96.55 5— 8 6 0.0258 0.04005 10.950 0.0328 0.0071 85.4-2.2 86.17 9-13 8 0.0258 0.04005 10.534 0.0254 0.0144 72.llo.98 70.44 14-18 8 0.0258 0.04005 10.243 0.0186 0.0206 56.3I0.3 56.06 19-23 9 0.0258 0.04005 .9.864 0.0102 0.0271 37.6+0.5 37.76 24-27 6 0.0258 0.04405 9.527 0.0055 0.0316 25.ill.0 25.27 28-31 6 0.0258 0.0528 9.217 0.0029 0.0345 I8.1I2.2 17.11 33-35 6 0.0258 0.0560 9.176 0.0028 0.0357 I4.7I1.3 16.23 36-39 8 0.0258 0.0640 9.075 0.0023 0.0381 I3.6I0.I 14.17 40-43 9 0.0258 0.0720 8.862 0.00132 0.0352 II.5I0.8 10.45 44^46 4 0.0258 0.0800 8. 806 0.0012 0.0367 9.68I0.60 9.60

Average =• 101 2 Average K ,6 IF 17 STD “ 0.051

& I.ZOpE 02 -

ooaouonnoouonooauouuiiDroua ouoooou ♦ 00# ♦ ♦ 9.600E Oi - on 00 00

T.ZOOE 01 -

K lobs

4.Z00Ë 01

0 000 Z.400E 01 - O# * |)QU

ODCUO - 000 00000 ooooouuo 0. 0 oooouooiiiiiiiiiiiiiiiiiiiiiiiiiiniiiiiiiiiiiiiiiiiiiiiiiiiiiiiiiiiiiiiiiitt II >1111111111II iiiint II NPNP> 0 r.SOUE 00 B.OeOE 00 9.860E 00 I.IUE 01 I.Z9ZE 01 i.oroc 01 pH

Figure 32. A Plot of pH vs. for MegN(CH2 )4HH2 . 1 5 0 of dehydration, and it was assnmed that this equilibrium is es­ tablished. The remaining 59 recorded determinations of observed in solutions of pH 9 -9-8 »9 , were treated assuming that the change in hydrate concentration with respect to change in time is zero, dH/dt = 0 . Under these conditions, pH 9 *9 » the half time for dehydration is greater than 0-5 sec and observations for were recorded before as much as 206^ of the hydration- dehydration equilibrium can be established. All values of were used to calculate the average Eg and Kgp by employing eq. 83 and the initial pH, Tablé 59 * The reverse calculation using Kg and Kgp yields the average observed equilibrium constants, K^ave» which are plotted with Kg^^^ against pH, Figure 53 .

Rate constants for dehydration of the carbinolamine. The equilibrium constants, Kg^^^^ and used in the kinetic expression were calculated as previously described, eq. 85, Table

40.

The 200 recorded observations for determination of kg in solu­ tions of pH 11.5-10.5 were treated assuming that the hydration- dehydration equilibrium, Kp, is established rapidly relative to the establishment of the carbinolamine-imine equilibrium, K^. At pH 10.5 the ratio of the rate of dehydration, V^, to the rate of imine formation, V^, is greater than five. The 528 recorded ob­ servations for determination of kg, in solutions of pH 9 *9-8 *8 were treated assuming that the change in hydrate concentration TABLE 39 Equilibrium Constants for Carbinolamine Formation with N,N—Dimethyl—l,4“butanediamine

Le m u f ] Le m i p ] Run pH ^Cobs Cave 310 # at to M M M-1 tri

1-4 7 11.524 0.0379 0.0022 2.83-0.73 2.13

5-8 S 11.064 0.0343 0.0057 I.49Ï0.49 2.00

9-13 7 10.682 0.0284 0.0115 I.85I0.32 1.77

14-18 7 10.344 0.0310 0.0184 2.IIÏ0.35 1.47

19-23 8 9.990 0.0128 0.0254 1.07-0.39 1.14

24-27 2 9.688 0.0079 0.0314 0.924-0.21 0.882

28-31 5 9.362 0.0044 0.0367 0.999^0.085 0.655

32-35 5 9.264 . 0.0039 0.0380 0.846-0.130 0.613

36-39 6 9.156 0.0030 0.0403 0.525-0.129 0.531

40-43 9 8.990 0.0020 0.0400 0.591-0.124 0.439

44— 46 4 8.927 0.0018 0.0419 0.289-0.054 0.406 Average =* 2.22 Average K^p “ 0.678 tn STD = 0.26 H 7.1001: PO -

ouunoooii'iuoooouiiuiiciooaouwii} 000000 ■joa 000 0 00 0 00 1.8*0^ CO - * * eu * - o ‘ o - o 0 u - 0 0 • I.1B07 00 - ♦ ♦0 ♦ Ü 0 II 0 ^obs n# ■ - O u 0 4.700k-01 - ♦ 40 * ♦ ♦ * : # «” ; I o“ : ; . - 0 : • o“ 00 0 # 4.00Qk-OI - t u * t * • ' 0 00 0 _ 00 • Ou - 00 000 dOüO - 00000000 -7.)84k-07 001111111111111111111111i1111111111 II 111111 II Ii11111111111111M111 II 11111111111111111111111111i111111 NPMP- 0 /.lOUE 00 a.ssot 00 o.a&ot JO l.lltE Ol I.747E 01 1.170k 01 pH . H Figure 33' A Plot of pH vs. for MegN(CHg)4NH2 . ro TABLE 40 Rate Constants for Imine Formation with N, N~Dimethyl* 1,4-butanediamine ^

Runs pH ^IMcal ^Ccal ^2obs ^ 2Hobs 310 # average sec sec“^

1-4 8 ll.5i9io.OO6 94.5 2.13 13.7-0.5 5-8 6 1 1 .007^ 0 .06 85.8 1.97 13 .6± 1.4 9-13 5 10.608+0.08 72.1 1.71 I2 ,4i l .2 14-18 6 10.293+0.06 57.2 1.43 13.1+0.5 19-23 6 9.927+0.06 39.5 1.08 ii.sio.7 10.2^0.6 24-27 6 9.608+0.08 27.1 0.822 1 1 .4± 1.2 10.1+1.1 28-31 6 9 .270±0.05 17.7 0.598 13.1±2.6 1 1 .7^ 2.4 32-35 6 9 .235± 0.06 16.9 ' 0.577 11.6±2.2 1 0 .3^ 1.9 36-39 6 9 .075± 0.08 13.7 0.485 1 2 .0±O.S 10.7^0.5 40-43 6 8.926±0.06 11.1 0.405 13.4±1,1 12.oil.0 44—46 5 8.867io.06 10.2 0.375 I3.6+L.2 I2.lil.l

All solutions contain enough KaCl for n “ 0.12 154 with respect to change in time is zero, and also treated assming that the dehydration equilibrium is established rapidly relative to establishment of carbinolamine-imine equilibrium. These two calculated rate constants, were averaged. From pH 9-9 to 8.9 the ratio of the rate of dehydration to imine for­ mation, V^Vj, varied from three to one.

d. The reaction of H,H-dimethyl-1 ,^-pentanediamine and iso­ butyraldéhyde . Equilibrium constants for carb inolamine - imine formation; The kj recorded determinations of KjQ-j^g were used to calculate the average Kj and Kjp by employing eq. 82, the final 98 pH, the pKai of 10.^12, and the pKag of 9.518, Table 4l. The reverse calculation using the calculated values of and K^p yields the average observed equilibrium constants, K^^ve) ^^hich are plotted with Kp^^g against pH, Figure 54 .

Equilibrium constants for carbinolamine formation: The 24 recorded determinations of K- . , observed in solutions of pH Cob s' 12.5-10.3, were treated assuming that the hydration-dehydration equilibrium, Kp, for isobutyraldéhyde is established before ob­ servations are recorded. The pseudo first-order rate constant for dehydration at pH 10.5 , 5-8 sec"^, affords an estimate of a half time of O.18 sec. Thus, observations for , extrapo­ lated to zero time from data taken after about 0 .2-0.4 sec, occur after one and one-half lives of dehydration and it was assumed that this equilibrium is established. The remaining 17 recorded TABLE 41 Equilibrium Constants for Imine Formation with N,N-Dimethyl-l,5-pentanediamine

pH [e m u p ] Le m i p J K-r t, Run # ^T “ t lobs *^Iave 335 M M at tq. M M If M”

1-4 7 0.0232 0.0400 12.336 0.0394 0.0006 100.9-6.3 97.1 5-9 8 0.0232 0.0400 11.408 0.0354 0.0045 88.5-7.3 90.0 10-12 6 0.0232 0.0400 10.878 0.0276 0.0118 80.0-4.0 75.8 13-16 6 0.0232 0.0400 10.120 0.0098 0.0242 39.8^7.3 40.6 17-20 6 0.0232 0.0440 9. 802 0.0050 0.0257 29.6-0.6 27*2 21-24 6 0.0232 0.0528 9.489 0.0023 0.0244 17.2-0.3 17.0 25-27 3 0.0232 0.0640 9.363 0.0018 0.0256 11.5-0.5 13.8 28-31 5 0.0232 0.0800 9.155 0.0011 0.0239 1 1 .8^0.7 9.48 Average “ 98.1 Average K^p “ 27*4 STD “ 0.11

ü vn I.200C m -

rnon,:ur'f:co:ii(iui:c ;cn: 9.600E n - ♦ ♦ * onocococd» • ouou eu n# nn U eu o - ; » 0 L# 7.200E Cl - ♦ . ♦ 0 ♦ O 0U ^^obs o 0 II 0 4,800e 01 - ♦ ♦ c n ii#

0 n •1 u 2.AOOE Cl - ♦ CI ♦ ♦ ♦ . ♦ U CO • oc # 0* 000 CO cnoc ciicecüfiii 0.0 ücccuccrotccccr uiiriii11iin n 11ii11iin111iitit 1111111111n 11 u il11il il11ilt M 111iinIn11111r111ii

NPNP* o T .îO u F Oil rl.5*ÎO t 0 0 O.NAOC CO l.lUC 0| I.242»- 01 1.370*: U PH

Figure 3k. A Plot of pH vs. for MegN(CH2 )5 NH2 - 157 determinations of observed in solutions of pH 10.0-9 .5 » were treated assuming that the change in hydrate with respect to change in time is zero, dH/dt = 0 . Under these conditions, pH

10.0 , the half time for dehydration is greater than 0.5 sec and observations for were recorded, before as much as 2Cfi^ of the hydration-dehydration equilibrium can be established. All values of were used to calculate the average Kq and Kgp by employing eq. 83 and the initial pH, Table k2 . The reverse cal­

culation using Kq and Kgp yields the average observed equilibrium

constants, which are plotted with K^^^g» against pH, Fig­ ure 35 •

constants for dehydration of the carbinolamine; The

equilibrium constants, K^q^I ^^cal’ iu the kinetic

expression were calculated as previously described, eq. 85, Table

43.

The 174 recorded observations for determination of kg, in

solutions of pH 12.3 -10.2 were treated assuming that the dehydra-

tion-hydration equilibrium, is established rapidly relative to

the establishment of the . carbinolamine-imine equilibrium, Kj. At

pH 10.2 the ratio of the rate of dehydration, to the rate of

imine formation, V^, is greater than five. The 174 recorded

observations for determination of kg in solutions of pH 9 .9-9.2

were treated assuming that the change in hydrate concentration

with respect to change in time is zero, and also treated assuming

that the dehydration equilibrium is established rapidly relative TABLE 42 Equilibrium Constants for Carbinolamine Formation ■with N, N-Dlmethyl-1 f 5“pentanedlamlne

Le m u p ] Ce m i p ] Runs pH ^Cobs ^Cave # 335 at to M M M -1 M -1

1-4 7 12.338 0.0394 0.0006 1 *73^0.67 1.73

5-9 6 11.452 0.0358 0.0041 1.67*0.40 1.63

10-12 6 11.004 0.0300 0.0097 1.59^0.50 1.46

13-16 5 10.276 0.0132 0.0228 0.920*0.30 0.933

17-20 5 9.962 0.0076 0.0268 0.587*0.095 0.679

21-24 4 9.642 0.0038 0.0280 0.425*0.115 0.455

25-27 -a 3 0.122*0.045

28-31 5 9.286 0.0017 0.0289 0.341*0.054 0.262

a Entry not averaged Average “ 1.75 Average K^p “ O.62I STD = 0.12

00 2*?00( »J0 •

l«7oOE ?0 - * ocoîncmîiîonncuccun* OOnCOUCD^rjO \ # QOOf] 000 e no 00 o 00 o 0 1-320E on * a «

K,Cobs

6»80üc-ül

n 0 c 0 a • 0 II 4 .4 0 0 E - O I I# * # n Ü0 0 ou - oa cou nnon CLIXCCU •I.nnc-Oh uccciiocrooccci 111111111n 1111t i 11111111i t 1111111n 1111111i i 111111111111111n 1111111111111itiii1111111

r.JO&C no E.580E 00 O.B60F CO I.II4E 01 1.242E 01 l.itOe ,ll pH

Figure 35- A Plot of pH vs. for MeaN(CH2 )5UH2 . VO TABLE 43 Rate Constants for Imine Formation with N, N-Dimethyl— l>5“pentanediamine £

Run pH ^IMcal *Ccal ^ 2obs ^2Hobs # —1 —1 335 average )ri *ri sec sec

1-4 6 1 2 .337^ 0.001 95.3 1.73 20.5+0.7

5-9 6 11.430+0.01 88.7 1.63 21.1+1.5

10-12 5 10.941+0.06 76.7 1.43 2 1 .6+ 2.0

13-16 6 10.198+0.08 43.4 0.868 1 9 .3± 1.5

17-20 6 9.882+0.08 29.6 0.619 20.2+1.7 . 17.8+1.5

21-24 6 9.566+0.08 18.8 0.408 1 8 .5+1.8 16.4+1.6

25-27 5 9.418+0.05 14.7 0.326 2 1 .2+1.0 18.7+0.9

28-31 6 9.220+0.07 10.5 0.234 21.6+0.9 1 9 .2± 0.8

3.All solutions contain enough NaCl for (x = 0.12

S l6l to establishment of carbinolamine-imine equilibrium. These two

calculated rate constants, and k^^^g, were averaged. From pH 9-9 to 9*2 the ratio of the rate of dehydration to imine forma­ tion, varied from three to one.

Correlation of equilibrium data

The preceding equilibrium data for formation of carbinolamine and imine from isobutyraldéhyde and the four diamines are corre­ lated with the acidity of the primary protonated diamines, and the

Taft substituent constants for each diamine of the form R-HHs, where E is the substituent containing a free or a protonated ter- 5 tiary amino group. From existing data a relationship, eq. 86, between the pKa of a primary alkylamine and the Taft substituent

pKa = 9-95 - 5*650* (86)

constant of the alkyl group is obtained. The eight pKa’s and

substituent constants for the four ^-dimethylamino-n-alkylamines

are given in Table 44.

The equilibrium constants for carbinolamine, and car­

binolamine-imine, Kjp, formation from monoprotonated diamines and

the aldehyde are calculated from the observed equilibrium constants,

K^Qljg and K^Q^g, using eq. 82 and 85, respectively. This method

of calculation uses the total concentration of the monoprotonated

unsymmetrical diamines; however, by definition the previous

equilibrium constants evaluate the addition of unprotonated pri­

mary amines to the aldehyde. These amines protonated at the TABLE 44 Calculations of Taft Substituent Constants for Diamines

a O' * + PK*Mlp- (CH2 )„NMe^ ( C Hgi^MMe, Amine n

2 9.240 0.175 6.866 0.83

3 9.901 0.0014 8.536 0.384

MejHCCHp^NHj 4 10.104 -0.0422 9.330 0.170

MegNCCHgigNHg 5 10.411 -0.126 10.201 -0.069

“ From Table 19

& i 63 9 9 primary nitrogen are not reactive. Corrected equilibrium con­

stants are defined in terms of reactive monoprotonated diamine

concentration, eq. 8%. These constants, and are cal­

culated from the fraction of monoprotonated diamine protonated

at the tertiary nitrogen and Kjp and K^p, eq. 88 and 8 9, see Table S ■ ■ SSi <”

K = K • I^EMIP] = [CH] . [EMIP] (Qg) [EMlPb] [a ][EM1P] [EMlPt]

The equilibrium constants are plotted against the acidity con­

stants and Taft substituents constants of the diamines. Figures 36-

39, Table 46. The straight lines in these graphs are those obtained

from data on monoamines

An imidazolidine is a heterocyclic compound -which is formed

by condensing a primary or secondary diamine with two carbon atoms

between the amino groups with an aldehyde or ketone. The first

reported imidazolidines were those formed by adding ’-diphenyl- 100 101 ethylene diamine to formaldehyde and benzaldehyde, eq. 9 0 *

99- J* C. Craig, Jr. and F. A. Via, unpublished observations , with méthylammonium ion. The Ohio State University.

100. (a) C. A. Bischoff, Ber., 31, 3248 (I8 9 8). (b) G. T. Morgan, W. J . Hickinbottom, and T. V. Barker, Proc. Roy. Soc. ( London), AllO, 502 (1926). TABLE 45

Calculation of Equilibrium Constants for Protonated Diamines

[EMlPt] *1? *IPE ^CFE *IMPE Amine [e m i p ] M -1 M-1 M-1

Diamine-2 15.46 5.217 0.381 40.6 13.7 26.9

Diamine-3 20.0 0.745 0.291 68.7 2.56 66.1

Diamine-4 17.6 0.678 0.344 51.2 1.97 49.2

Diamine-5 27.4 0.621 0.207 132.4 3.00 129. 165

2.2

2.0

21 20

1 .4 -

78 9 10 pKa

Figure 56. A Plot of pKa vs. log Kj for Diamines "with the Linear Function from Figure 21 . 2.0

hi # 21 20

I

1.0

- 0.2 - 0.0 0.2 O.tl- 0.6 0.8 T*

Figure 37 - Plot of o* vs. log K for Diamines with Linear Function Calculated from Least Squares Treatment of the Monoamine Data.

& 167

21

1.0

0.6

31

p

0.0

-0.5

6 T 8 9 10 pKa

Figure 5 8 . Plot of pKa vs. log Kg. 168

21 1.0

0.5

20

-0.5

0.0 0.2 0.4 0.6 0.8 a*

Figure 3 9 - Plot of a* vs. log Kq for Diamines with Linear Function from Figure 23 . TABLE 46

Carblnolamine and Carbinolamine-Imine Equilibrium Constants, Taft Substituent Constants and pKa Values for the Free and Monoprotonated Diamines log log ■ Amine Code *C < f M - 1 •'l M - 1 %C MagNCCH,),- 20 9.240 0.175 48.9 1.69 1.98 0.297 ..

30 9.901 0.001 102.4 2.01 1.02 0.009

MegNCCHg),- 40 10.104 -0.042 101.2 2.00 2.22 0.345

MegNCCHg)^- 50 10.411 -0.126 98.1 1.99 1.75 0.243

PK^M2p log a MÏÎ= % CPE

MegHNtCHg)^- 21 6.866 0.830 40.6 1.61 13.7 1.14 25.4

MegHNCCHg)^- 31 8.536 0.384 68.7 1.84 2.56 0.408 20.0

Me^NCCH^)^- 41 9.330 0.170 51.2 1.71 1.97 0.294 —

MegHNCCHg)^- 51 10.201 0.069 132.4 2.12 3.00 0.477 — 170

The diamines employed in this investigation contain a tertiary Ph

PhMCHsCHsNHPh + RC=0 (90) sr k nitrogen which prevents the formation of a stable imidazolidine, eq.. 91. If this unstable zwitterion were formed, it would

MegUCCHalnMs + RC=0 + % 0 (91) H rapidly abstract a proton from a water molecule and cause the aqueous solution to became much more basic as the reaction pro­ ceeds to equilibrium. Experimentally the aqueous solutions be­ come slightly (ApH »= O.l) more acidic as the reaction approaches equilibrium conditions. Imidazolidinium ion formation from iso­ butyraldéhyde and monoprotonated ' -dimethylethylenediamine has 102 been observed, and this would indicate that monoprotonated u)- dimethylamino-n-alkylamines can also form imidazolidinium ions, eq. 92. Inspection of Figure 56, a plot of log K]- vs pKa, reveals

Me Me V/ + Ry. Me2M(CH2)/H2 + R-C=0 (92) H H

101. A. Eibner and G. P. Purucker, Ber., 33 a 3658 (19OO).

102. K. ¥. Narducy, unpublished observations, Ohio State Univer­ sity. 171 a large positive deviation for monoprotonated lT,N-dimethylethyl- enediamine and monoprotonated N,E-dimethyl-l,$-propanediamine from the linear function obtained for the monoamines. Since it is not chemically possible for the monoamines to form a hetero­

cyclic compound this positive deviation is attributed to imida­

zolidinium ion formation and the equilibrium constant, for these diamines is redefined, eq. 93. The equilibrium constant for imidazolidinium ion formation. Kg, can be estimated by

^ _ [i h ] + [CH] + R /Q.\ ■ A [ E M l P t ]

subtracting the theoretical equilibrium constant, the linear

function, (Pig. $6) from the observed value.

observed theoretical

KjeE % E E %

n M"^ M ”^ M ”^

• 2 40.6 15.2 25.4

5 68.7 43.7 25.0

Deviations from the linear function (Fig. 36) for monoprotonated

N,N-dimethyl-1,4-butanediamine and monoprotonated K,N-dimethyl-

1,5-pentanediamine are considered small and attributed to experi­ mental uncertainty. 172

Rate data evaliiation

The rate-limiting step for imine formation from the diamine and isobutyraldéhyde is assumed identical to that for the mono­ amine, the dehydration of the carbinolamine intermediate. The rate-limiting step in this reaction for monoprotonated N,1T- dimethylethylenediamine, which forms the imidazolidinium ion, is also the dehydration of carbinolamine. The results of kinetic 103 studies of imidazolidine formation by Jencks and coworkers, 104 and more recently by Benkovic and coworkers, indicate that for the reactants (basic amines) and reaction conditions (basic pH

ranges) of this investigation, the rate-determining step is not

the ring closure reaction to form the cyclic cation, but dehy­

dration of the carbinolamine.

Data collected for H,N-dimethylethylenediamine indicate that

dehydration of the unprotonated carbinolamine, C, is subject to

general acid catalysis and dehydration of the protonated carbin­

olamine, CH, is subject to only specific acid catalysis. A two­

fold increase in concentration of the general acid in pH ranges

where the carbinolamine is in the protonated form ( Table runs HO HO • . . I + iPr-C-H(CH^)^me2 iPr-C-H(CH2)^e2 . H H H H H C CH

103. R. G. Kallen and W. P. Jencks, J. Biol. Chem., 24l , $843 (1966).

10*4. S., J. Benkovic, P. A. Benkovic, and D. R. Comfort, J. Amer. Chem. Soc., 91, 5270 (1970): 92, 523 (1970). 175

210-4 9 -6i) does not significantly affect the observed rate

constant.

The observed rate of dehydration of the unprotonated car­ binolamine can be described, eq. 97, by the sum of the uncata­

lyzed mechanism, eq. 9k, the general acid-catalyzed mechanism,

eq. 95, and the specific acid-catalyzed mechanism, eq. 96.

HO ^ o i-PrCmH(CH2)nMe2 > i-PrCH=]Ë(CH2)^me2 + OH" (9%)

HO I + i-PrCHNH(CH2)jMe2 + BH — i-PrCH=m(CH2)+ H2O + B (95)

HO i-PrCHim(C3l2)n™e2 + ^ ■> i-PrCH=&(( % ) ^ meg + H2O (9&)

k2obs [CT] = + KugC [BH] + k^C [H+] (97)

Equation 97 can be simplified by using the ionization constant for

the carbinolamine of form CH.

= # r (58)

k2obs [CT] = k^oC + ku^cCBH] + k^i^KaciCcH] (99)

The observed rate of dehydration of the monoprotonated car­

binolamine, CH,. can be described, eq. 103, by the sum of the un­

catalyzed mechanism, eq. 100; the intramolecular catalyzed

mechanism, eq. 101; and the specific acid-catalyzed mechanism,

eq. 102. 174 HO i-PrCinm(CH2)j^mMe2 i-PrCH=&( + H0“ (lOO)

HO 1-PrCH]m( CHg) -5e_> i-PrCH=&(CH2)^e2 + HgO (lOl)

Hj) l-PrCHHH(CH2)ij!me2 + H"^ i-PrCH=HH( GH2)nHHMe2 + H2O (102)

ksobs^CT] = kpoCCH] + k^iCcH] + k^^^[CH][H+] (IO5)

The observed rate of dehydration of carbinolamine, of all

forms, under the conditions of this investigation can be described by eq. 104, the sum of eq. 103, and eq. 99-

ksobs^CT] = + k^[BH])c + k^Kagg^CcH] + kpolCH]

(104) + kpiCCH] + k^^[CH][H+]

From eq. 104 it is now visually apparent that the mechanisms of eq.

96, eq. 100, eq. 101 are kinetically indistinguishable, that is, the contribution from each mechanism is defined as a constant mul­ tiplied by the concentration of protonated carbinolamine, Ich]. If kg^- is defined as in eq. 105, and eq. 104 is divided by Cct] , we

get eq. I06.

^Sobs “ [ct T ^ '*■ C106)

The ratio c/Cc t ] is the fraction of free carbinolamine, fu, and

[ch]/[ct ] is the fraction of monoprotonated carbinolamine, fp. 175

ksobs = fil \io + fu + fp k g ^ + fp (107)

To estimate the fraction free carbinolamine, fu, and the fraction carbinolamine protonated at the tertiary nitrogen, fp, the first

(eq.. 98) and second (eq. I08) ionization constant for the carbin­ olamine are required, as shown in eq. 109 and eq. 110. The first ionization constants can be calculated from the ratio of the car-

fu = 1.0/[l.0 + ] (109) KaQ]_ (Kaç2^)(Kaç2)

fp = fuCH+3/Ka^^ (110) binolamine equilibrium constants, eq. 112, and are listed in Table hJ .

KÇ ^ C _ A[EMlFt] ^ c [h +] ^ [EMlPt] Kggg a LEMUP] ’ CH CH ' [mu?][H+]

^C 1 Kcee ■ tomt

The second ionization constants for carbinolamine, eq. 108, are estimated from the ionization constants of the diamines, eq. 115, 105 and the pKa values of Clark and Perrin, which estimate the

affects of structural changes on the pKa of amines. An estimate TABLE 47 Calculation of the Ionization Constants for Monoprotonated Carbinolamine, Ka Cl

pKaci n ^CPE ^ ^ ^ I P t ^®MlPt ^«Cl X 10 X 10

2 1.98 13.7 9 .031 9.30 1.34 9.873

3 1.02 2.56 9.514 3.05 1.20 9.921

4 2.22 1.97 9.824 1.50 1.69 9.772

5 1.75 3.00 9.830 1.48 0.863 10.064

o\ 177 of the first ionization constants for carhinolamines, eq. 108, is also obtained by this method; however, these values are not used to evaluate the kinetic data. The discrepancies between these estimated values of pKa^^, Table 48, and those calculated. Table

4-7 , result from the limitation of the method of estimation or an experimental deviation in determining .

The rate data are evaluated with eq. 107 by employing a least squares method of minimizing the fractional standard de­ viation. The results of this calculation are listed in Table 4-9 .

The maximum percent contribution of each rate term to the observed rate is listed in Table 50 .

Evaluation of the pH-rate data for IT ,N-dimet hy let hylene­ diamine yields positive values for the four rate constants. A pH-rate plot. Figures 4-0 and 4-1 , illustrates that the contribu­ tion from each rate term is mutually exclusive of the contribu­ tion of the other terms with the partial exception of

The evaluation of the pH-rate data calculated for this diamine with an estimated K^p, Table 34 continued, is also listed.

This treatment was employed to check the validity of the general- acid catalysis rate constant: for it was initially considered that the increase in the calculated value of kg^^g around the pH of 8 .4-10.0 may be minimized or eliminated if a reasonable es­ timate of Kqp were employed in the rate expression, eq. 67, rather

105. J. Clark and D. D. Perrin, Quart. Rev. (London), l8, 295 (1964). TABLE 48 Estimation of Ionization Constants for Protonated Carbinolamines of the Diamines

Changing £ H N Changing H to OH to HNR ^ Amine n (n+2) carbon n-C atoms away atom away P^^Cl

N , N-Dimethyl ethyl­ ene diamine 2 9.031 —0.1 -0.28 8.65

N,N-Dlmethyl-1,3- propanedlamine 3 0.514 -0.05 -0.14 9.32

N,N-Dimethyl-1,4— but ane diamine 4 9.824 -0.025 -0.07 9.73

N,N-D imethyl-1,5- pentanediamine 5 9.830 -0.013 -0.04 9.78

Changing® Adding OH ^ n PKaM2o 1° to 2° one C atom away P ^ C 2

2 6.866 +0*38 —i «88 5.36

3 8.536 +0.38 —1 « 88 7iOO

4 9.330 +0.38 *-l .88 7.83

5 10.20 +0.38 —1.88 8.70

— pKa values obtained from Clark and D. D. Perrin, Quarterly Reviews, 18, 296 ^ (1964). & Value from ref. Sa. TABLE 49 Calculated Rate Constants for Imine Formation From Isobutyraldéhyde and Four Diamines

Amine R-NH. k k . -1 STD Status R ^ sec ■ M-ÏL c -1 M - C - l M sec

4.01 3.46x10^ 25.4 5 .54x10 ^ 0.112 real

4.05 3.36X10^ 26.2 0.120 discarded 0 MejNCCH^)^- 21.1 - 3 -39x1 0 ^ 4 4 .5 2.42x10 0.077 discarded

19.3 41.8 4 .34x1 0 ® 0.086 real

14.2 -67.9 9.83 3 .43x10^ 0.065 discarded

13.6 —— 8.24 4.30x10^ 0.067 real

20.6 1 .93x1 0 ^ 10.2 2.39x10^° 0.053 discarded

21.3 —— 15.4 1.76x10^° 0.059 real

Me^NCCHj)^- 3.44 3 .72x1 0 ^ 498 1 .57x1 0 ® — discarded

— Values obtained from evaluation of pH-rate data with an estimated K CP Table 34* TABLE 50 Greatest Percent Contribution of Each Rate Term to Observed Rate Constant

k Diamine R—NH uo at pH ^uB at pH at pH at pH R % % % %

Me^NCCH^^- 100. 11.8 55. 9.8 100. 7.2 14. 6.2

Me^N(CH^)3- 100. 11.0 ———— 94. 8.4 15. 7.7

100. 11.0 — — 55. 9.0 35. 8,9

11.0 17.2 9.8 9.6 “®2»(C«2>S- 91. 37. 44. 9.2

S) o l8l than the observed value. With respect to the second order rate

constants (Kgk or Kgpk) for imine formation from amine and al­

dehyde j this alternate method of data treatment did not yield

results ■which were significantly different from those illustrated

in Figure 4l and is not given further consideration.

Evaluation of the pH-rate data for -dimethyl-1 -propane-

diamine and W-dimethyl-1,4 -butanediamine yields positive -values

for only three of the four rate constants. It appears that these

diamines are not subject to general acid-catalysis and the pH-

rate data must be evaluated with eq.. 111)-. The resulting pH-rate

plot is illustrated in Figures 42 and 4-3 for W,N-dimethyl-l,3 -

propane diamine and Figures 44 and 45 for W,N-dimethyl-1 ,4 -

butanediamine.

ksobs = ^ o + ^ I M + ^ph

Evaluation of the pH-rate data for -dimethyl-1 ,5 -pentane-

diamine yields positive values for the four rate constants. A

pH-rate plot. Figure 47 reveals that the contribution from each

rate term is not exclusive of the contribution of the other terms.

Because of the extreme overlap of the contribution from and

kgQ^ with k^-g it cannot be definitely determined if is real

or not. Since N,W-dimethyl-1,3 -propanediamine and N,N-dimethyl-

1 ,4 -butanediamine are not subject to general acid-catalysis,

k^g = 0.0, a reasonable extension of this behavior would be to

assume that W,W-dimethyl-l,5 -pentanediamine is also not subject 4.5001 01 •

0000 - @ 0 0 (I 0 00 3.600L Cl * * $ 00 * » - • 0 0 0 0 @ 0 0 00 - • uo0 0 OO 0 # 00 0 - 000 2.7001 ; i - * @ 0000 • « 0 4 000000000000 OOOOUOOOOOOOOOOOJOOO • 0 - • • • 0 ksobs ; , - O > - 0 l.BOOC 01 - « » « * 0 0 - ' ? 0 0@ o 0 0 - 0 9.0001 00 - ' ♦ . ♦ f 00 0 0@ 000 00000 - 0000

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Figure 4o. A Plot of pH vs. ksobs Me2N(CH2)2NH2. For This Plot the Average % Line Was Generated Assuming that [EMUP] + [EMIP] = O.CA M. ^ 4.bOO^ 01 *

UOOO OU Ü U 00 }«60Uê t*| -

IIU UCiO 2.TOOL ot - * I'UOJ ♦ ouuuuooonoiio uucuonociiUuuunûuùUüdUüUüûiiûnuuuOüiKiuoo OUUOQOO OUOUOlU CLOU onoo n - ÜÜÜ on -ou ou kg.obs u OJ ou oûuon u ou ÜO 0 I.BOOE : I - 0 0 ^ 0 ♦ ♦ 0 0 0 n 0 uo ü o n 0 0 0*0 0 0 ü 0 0 0 0 0 0 0 ^ n” 0 0 n IJ 0 0 0 i.oooi: 00 - 0 # 0 0 « 00 » 0 0 0 0 0 0 ou ou un ÜO fJÜÜ 0 u ouooo u 00 00 ÜO OOÜG -IIOÜ ^ uo iiuuuiluooüuôouuüuniiütiuüoüoo ou ou B» oouoooo UDO no ocoo :ooo onoonooonoonoo onjo unuo iioiincüouo GuoQUüuoü onunuoiiouo onuujooououou 0.0 Oouoo(juuuuuo«iùuni}ouùut)utiuuuojuu0t}0:jüiiuououu .iJonauoüüuüuuoJouunurKiQouüuuoiiucoiiooiJfiaou(;o3uncoour:oüü(tuoi NMK" 0 O.oooe 00 7./00E 00 •i«4ooe ou 9.6oum oo 1.20UC 01 PH Figure 4l. A pH-Rate Profile for Mo2N(CHg)aNHg with the Calculated Contribution from Each Kinetically Distinguishable Rate Term Assuming CemUP] + & [EMIP] = 0 .0^ M. JOj 'SA Hd Jo q.oid[ v '3ij sJtiSTJ

Hd Iw- 13'.'

- 00 3000*6

10 3008*1 000j0li0f#000 puono 000 sqos^ 03 • 00 09 9 0 00 00 - 10 3003*2

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ccixntt ctcccciu'T om-unutcooct. .ucrn { Û (111 -LC ÜLÛtlJLUlûCf. Uü'.t - Il LilU UOu 00 i! no (in 0 3.4006 yi - -.1 ♦ U » O Ou* ♦ ♦ - 0 Ü (I II - I, n 0 on - i: *-0 0 0 - , II 0 0 0 II o n - I’ 0 c o g - G 0 u on ; ; ■ u - g g o 2.T00E 01 - U «0 » #0 U ♦ ♦ nij c. i: 11 Ü Ü Ü 0 uo 00 kg , - O c Ü ito ODS - o 0 CÜC cuuin - 0 c 0 crnro iccLCJ - 0 0 1«800£ VI - li* * ♦ ♦ CCOC^J ♦ r 0 - o i:rno no Ü u 0 uo . c C 0 IÎCJ 0 nc - Ü J n 0 Cl CU L li c , a 0 " Ü 8, 0 C 9.000k 00 - Ù ♦ Ü ♦ ♦ no 0 ♦ ♦ - li * c on u n no n u 00 Û n - r n cn o -oc ou 00 on u n uu ou cou iico cuu cure 0 1 , 0 0 i. o u ulcclucccu inu':*/:r« i»' 0.0 i:rci: iCuri.i»;c» cnocfir.iifjLCLcnroL’OfinrcccrMiuuccir occmiujûrmoonafiniainuuniv.innnoiincn.iMiMuoci'rntîi^iiUjM girr»*.. (%;oi 0 fi.OO'iL 0(1 7.200k UO H.4UÜC CO f.hOl k GU i.rijD^ C| I ./* %'••. .1 pH Figure 43. A pH-Rate Profile for MeaN(CH2)3NH2 with the Calculated Contribution S for Each Kinetically Distinguishable Rate Term. 2 . 5 0 0 t 01

2.0001: 01 -

l.bOOfc 01 - -

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NPNP» 74 6.000E 00 7.Z00E 00 8.400E 00 9.600E 00 1.080b 01 t.7U0E 01 pH o\ Figure 44. A Plot of pH vs. for Mezlf{cUs) 2.iOCC 01 0 U

0

2.000k 01 - U U

a u 0 J

l.tOOF 01 - U 0 0 D ouounounoijnoadDou O O aoucunjouoooouo ooou ouoo 0 0000 00 2obg n 00 G 000 0(1 ooojojounoo 00 “d . i.oooc ;i - a 0 - 0 00 — ll ooou 0 oooco uoboo 0 - n 00 0 ÜOU n OUO U 00 u 5.000C 00 - * ou » ou » * 0 0 0 00 00 0 Q ou 00 u ou 0 00 0 0 00 ou 00 ou • ou UO ou 000 UOU OUO ou cuo oocuo üoo 000 onoo OCOOUOOO UOUOCÜO OOQOJ n.lUOiJOOO 0.0 üUOOOuoonQunünoooüaucnaoüOQooooor:ououuuuocoi:u3couonuoooujuuuouoouououQunuQuuiuoür:joouoüOuoouooi.oiiOüüi NPNP" 71 6.030k 00 7.200E 00 8.400k 00 p jj 9.60Ce 00 1.000k 01 1.2CÛT 31

Figure 45* A pH-Rate Profile for MegN( 0112)41012 with the Calculated Contribution for Each Kinetically Distinguishable Rate Term. § 2.S00C 01 -

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I.OOOE 01 -

5.000E 00 -

0.0 iniitMiiiniiiiiiiiiiiiiiiniiiniiiiiiiiiiiiiniiiiitiiiiiiiiiiiitiiiiiiiiiiiiiiiiiiiiiiiiiiiiiiii NPyP* 6.000E 00 T.ZOOE 00 0.400E 00 9.600E 00 I.OOOE 01 1.200E 01

Figure 46. A Plot of pH vs. for Me^(CHg)sini2 . S 2.500E 01 - a u u e uoooooou oono noouonooQooDOJ 00 00 2.000E 01 - 0 * 00 t 0000000 » 0 000 0000 000 000 00 00 00 00

0 i.sooE r.i - 0 »

I.OOOE 01 -

0 0 oouooaoouo 00 0 00 - c 00 0 00 00 0 0 0 5.0006 00 - ♦ ♦ u ou 00 * t üO 0 t 00 0 0 u D 0 00 0 0 000000000000 00 0 0000 0000 00 0000 o 00 000 00 0000 0 000 000 000 . 0000000 00 00 000 ouoo OOÜUÜO 000 000 0000000000 00000000 000000000 uOUOO 00000 ÜODOOOOOOOOUO 9.0 uuucoooooooooooooooooooooooooooooooooooooooitiiiniiiiiIIiIIIi1IIIIIIIIIIIiouoooooooooooonooonooooooi

NP‘>P« 95 b.OOOE 00 7.200C 00 e.50or 00 9.60UE oo 1.0806 01 1.7006 01 PH

Figure kj. A pH-Rate Profile for MeaN(CH2)5NH2 with the Calculated Contribution & for Each Kinetically Distinguishable Rate Term Assuming General Acid-Base Catalysis. ?.*00k 01 -

000000000000000 0 p 000000 00000 2.000E 01 U ♦ ^ 0000 » 000 • 00 000 000 00 000 0 0000000000 00

l.ïOCb 01

Sobs

0 0000 ono 00 0 I.OOOE 01 - O *00 o 00 0 0 0 000 0 00 0 0 0 0 0 0 0 II 0 0 0 0 - 0 . 0 0 0 k.OOOE 00 * * 0 0 0 * * O 0 0 0 00 0 0 o 0 0 00 00 0 0 0 O 0 0 00 00 00 00 00 OUO 00 00 OOO 0000 000 00 0000 ooouuooo uoooo oooou ooooooou 000000000000000000000000000000000000000000000000000000000000000000000000000000000000000000000000000310.3 NP.XP* a* &.000E 00 T.ZOOE 00 a.aoob 00 9.600E 00 I.OSOE 01 1.200E 01 pH

Figure 48. A pH-Rate Profile for Me2N(CHg)sNHg with the Calculated Contribution o for Each Kinetically Distinguishable Rate Term Assuming No General Acid-Base Catalysis. 191 to general acid-catalysis. The pH-rate data are evaluated with eq_. Il4 and the resulting pH-rate plot is illustrated in Figure

48.

The pH-rate, plots for each diamine. Figures 40-48, illus­ trate the contribution of each mechanism; e is a representation of the observed rate constant ksobs» ^ is the contribution from the uncatalyzed dehydration of unprotonated carbinolamine, fu b is the contribution from the general acid-catalyzed dehydration of unprotonated carbinolamine, fu k^gCsH] ; c is the contribution from the specific acid-catalyzed dehydration of unprotonated

carbinolamine and the uncatalyzed and intramolecularly catalyzed

dehydration of protonated carbinolamine, d is the con­

tribution from specific acid catalyzed dehydration of protonated

carbinolamine, fp k^gCH'*'] •

The uncatalyzed rate term

A good linear relationship is obtained by plotting the second-

order rate constant for the formation of imine from the alde­

hyde and amines via uncatalyzed dehydration of unprotonated car­ binolamine, Kg, against the pKa of the diamines. Figure 49;

the values for the six mono functional amines are also included.

The specific acid-catalysis rate tern

All the observed rate constants for specific acid catalysis

of imine formation are plotted against the pKa of their respective

amine. Figure 50. The slope of this, linear function was determined 2.0

5 1.0 20 bO S 0.0 5S M Slope a 0.8 1.0

2.0

10 pKa

Figure^9 • Plot of log K ^ K q v s . pKa for Monoamines # , for Dl^mlnesI % 10.0

9.0

8.0

7.0 slope = 0.8

10 PK&M2p Figure 50. Plot of pKa^gp vs. log

§ I9h by employing a least sqnares method and excluding points 2 and 6.

As vith the rate constants of the preceeding paragraph, some of the experimental uncertainties can be minimized by plotting the log of the product of and the rate constants for specific acid-catalysis vs pKa for the monoamine and monoprotonated di­ amines, Figure 51-

Evaluation of kgQ%

This rate constant, éq. 105, is the sum of three terms, the rate constant for the uncatalyzed dehydration of protonated car­ binolamine, the rate constant for the intramolecularly catalyzed dehydration of protonated carbinolamine, and the rate term for

specific acid-catalysis of the dehydration of unprotonated car­ binolamine. This last rate term is evaluated by estimating the rate constant, k^, from the linear relationship of pKa and the log of the product of the specific acid-catalysis rate constant

and the equilibrium constant for carbinolamine formation. Figure

51. With the ionization constant for monoprotonated carbinol­

amine, Ka^^, obtained from Table 4%, and eq. 115, the sum of the

^o + ^i ~ “ ^^Cl (115)

rate constants for uncatalyzed and internal catalyzed dehydration

of monoprotonated carbinolamine can be estimated. The significance

of the internal acid-catalyzed mechanism can be realized by com­ paring the, second-order rate constants for imine formation, K__, CxXi ( , for monoprotonated diamines with the amines for which 195

10

31 9 a

21 bO 8 Q

slope = 0 .6l

7

6 7 8 9 10 pKa or p K a j ^

Figure 51. plot of pKa vs. log Kj kg for Specific Acid Catalysis. 1 9 6

internal catalysis is not possible. Figure 52 . From Figure 49 the second-order rate constant for the uncatalyzed dehydration of monoprotonated carbinolamine, can be estimated.

^ E e( ^ o + ^i) ^CEE ^ ^

The ratio, eq. Il6 , reveals that the tertiary monoprotonated

diamines of N,N-dimethylethylenediamine, -dimethyl-1 ,5 - propane diamine, E ,N-dimethyl-1 ,4-butanediamine, and E,E-dimethyl-

1,5 -pentanediamine, respectively, form imines 5,500, 50, 1-5 and

1 .Ô times as fast as estimated ( Table 45). The second-order rate

constant, eq. 117 and the first-order rate constant, for

KcPE ^i = ^FE^^o + ^i) " kpo (H 7)

internally catalyzed imine formation are listed in Table 51. 197

5.0

2.0

1.0 i I bO H

- 1.0

- 2.0

10 pKa

Figure 52 . Plot of pKa vs. log Kppg (kg^ + kpj_) with the Linear Function from Figure 49. TABLE 51 Evaluation of k SUM

1 0

•H M A ft. «1 AI 01 kP 00 + > -d| «1 1 «Ml w 0 0 0 •H •H 0) H bo •H + 0 .s •d X V»' H AS®* . o 0 M M + M 1 Ü joi 0 A ft. ft. 1 0 % 3 . u P uP '— /

e k:P I II Ill IV V VI VIIVIII IX X

21 6.886 3.6 0.05 25.4 348.0 2.54 0.10 3,500 348. 25.4

Me^NH(CH^)^- 31 8.536 6.3 0.84 40.9 105.0 2.02 2.14 50. 103. 41.0

41 9.330 14.0 2.4 5.84 12.6 1.10 8.00 1.5 4.6 4.0

Me^NH(CHg)^- 51 10.200 16.6 1.3 14.7 42.3 1.63 44.0 1.0 — 0.6

— Determined from Fig. 52} — from Table 47} ” from eq. 15} ” from Fig. 50. - V-VII — use K:^pgfrom Table 45 % CHAPTER V

DISCUSSION

A. Hydration of Isobutyraldéhyde

The hydration of isobutyraldéhyde and dehydration of the aldehyde hydrate in aqueous solution at 55° is subject to general acid-base catalysis. A crude Br^^nsted plot, with three points, for the acid-catalyzed hydration of isobutyraldéhyde yields a

"alpha" of 0 .44 , which is within the range of results previously 106 reported by Bell et al. for acetaldehyde in aqueous solution at 25°( a = 0 .54) and formaldehyde in aqueous solution at 25° (a =

0.23). A Br^^nsted plot for the base-catalyzed hydration of iso­ butyraldéhyde constructed from only two points, N-methylmorpholine and water yields a p of 0 .44 , which is in agreement with that reported for isobutyraldéhyde in aqueous solution at 0° (p = 0 .46),

In this plot, Figure 54, the point for the hydroxide ion catalysis lies above the line. This positive deviation may occur because hydroxide ion, like no other catalyst, can behave as a nucleophi- lic catalysis, adding to the carbonyl bond, as well as a Br^nsted base catalyst. Because water is amphoteric, the observed rate constant, Table 52 , for the uncatalyzed hydrolysis of

106. R. p. Bellj Advan. Phys. Org. Chem., 4, 1 ( I966) .

199 TABLE 52 Rate Data for Acid-base Catalyzed Hydration of Isobutyraldéhyde

Cpd. Ka pKa ^hBH if “1 M sec ‘‘hBH hB

^0-l3.6aa 0.0023.' —4 • 7 9600 3.98 “2° 55.5 X 2b 15.7 2bx55.5

440S 0.0023 55.5 —1 .26 2.17 «4*4 55.5 3 ^

N-methylmor— pholine 7 .27“ 0.030 -1.52 0.22 —0.66

— Determined ref. 95 P* 28. — Statistical factor. — Determined ref. 81. “ Determined ref. 2d.

ro o N-methyljnorpholinitmi ion

IfeO

0 2 h 6 8 10 12 14 16 pKa

Figure 5 5 * A Br^nsted Plot for Acid Catalyzed Hydration of Isobutyraldéhyde. fo OH

N -methylmorpholine -2

0 2 4 6 8 10 12 Ih 16

Figure $4 . A Br^nsted Plot for Base Catalyzed Iftrdration of Isobutyraldéhyde. 8 205 isobutyraldéhyde is a sum of water's ability to behave as an acid and a base in catalyzing the reaction. The point for "water lies above the Br^nsted line for acid catalysis and lies on the

Br^nsted line for base catalysis, indicating that -water's ability to catalyze the i^dration of isobutyraldéhyde is greater acting as a base than as an acid.

In his discussion of similar results -with acetaldehyde, Bell lists four possible stepwise mechanisms that are consistent with experimental observations. He then eliminates all of them on the grounds that observed rate constants demand that a step in each mechanism have a rate constant larger than the diffusion-controlled value or so high as to cause considerable curvature to the Br^nsted 106 plots, -which is not observed. Bell concludes by considering the possibility of a cyclic transition state. 82 The discussion of Pocker and Dickerson for the hydration- dehydration studies of propionaldéhyde, isobutyraldéhyde, and pivaldehyde is similar to that of Bell. The purpose of this in­ vestigation is not to elucidate a mechanism but to enable a reasonable estimate of the rate of hydration-dehydration of iso­ butyraldéhyde under the conditions, pH and amine concentration, for imine formation. Ho-wever, since the results herein reported are not dissimilar from those of other aldehydes or at other temperatures, it is reasonable to assume that the mechanisms

currently in vogue -would be applicable. The cyclic mechanism 107 involving one or more -water molecules suggested by Eigen was 20k 106 illustrated, by Bell for general acid catalysis, eq.. Il8 , and 82 by Pocker for general base catalysis.

H .♦ H ^

BH + )]=0 + 2HaO ---^ *"b '0— H— 0--H' *- *- H (118) /OH + BH + HeO OH

B + yC=0 + SHgO 7 ~~^ ')0-— H— a ''0-— H''" H (119)

^OH fg g+ ^OH - + OH" + BH+ , + B + IfeO OH ^ ^OH

The small arrows have been added to illustrate the proton move­ ment as the reaction proceeds along the reaction coordinate to

products.

B. Determination of the Micro-pKa Values for the m-Dimethyl- amino-alkylamine

The relative basicities of a series of primary, secondary,

and tertiary methylamines in aqueous solutions are determined by

the thermodynamic stability of the conjugate acids. Consideration

107. M. Eigen, Disc. Faraday Soc., 3 9 > 7 (19^5). 205 of inductive effects ■would predict the order 3° > 2° > 1°; whereas, consideration of the energetics of solvation, the probability of forming a hydrogen bond, would yield the order 1° > 2° > 3°*

108 Everett and Wynne-Jones reported the observed basicities, which 109 o Hine and Mulders extrapolated to 55 and zero ionic strength; dimethylamine, pKb = 3 •19 is more basic, than methylamine, pKb =

3 -3 T which is more basic than trimethylamine, pKb = 4 .1 0 . In a

solution one molar in a primary amine, one molar in a tertiary amine, and one molar in hydronium ions, the primary amine is ex­ pected to be protonated to a larger extent than the tertiary amine.

The micro pKa values for the diamines (XVl) indicate that in an equimolar solution of XVI and hydronium ions the primary nitrogen

MegN( CHg) nWHg

XVI would be protonated 1.6 to 4 times as much as the tertiary nitro­ gen as n changes from two to five (see Figure 55) • As the number of methylene groups between the two nitrogen atoms of the diamine, n, increases (from 2 to 5) the relative basicities of the primary nitrogens to that of the tertiary nitrogens increases. The pKa's of the primary nitrogens are more sensitive to changes in n than are those of the tertiary nitrogens; as the change in the pKa for

108. D. H. Everett and ¥. F. K. Wynne-Jones, Proc. Roy. Soc. (London), A17T, 499 (19^1)-

109. J. Hine and J. Mulders, J. Org. Chem., 32, 2200 (196T). 0.4

0.5

[EMlPt]

[e m i p ] 0.2

0.1

2 3 4 5 CO n ro Figure 55 • Plot of [EMlPb]/[EMlP] vs n. 207 thé primary nitrogens, ApKa^^g = 1.17 ApKa^^ = 5 -5 5 , is larger, than that for the tertiary nitrogen, ApKaj^^ = 0 .77, ^PKa^^gt ~

2.97, over the range of n. The largest change in the pKa values and in the ratio [EMlPt]/[EMlP] occurs as n changes from two to three. Substituents of Group I should be stronger electron with- drawers than substituents of Group II; for the accepted concept

Group I Group II

H^CHgCHs- MegNCHgCHs- HsNCHsCHsCHs- MegNCHgCHsCHg- of hyper conjugation considers the N-methyls as electron donators relative to hydrogen. However, the dimethylaminoethyl substituent decreases the basicity of a primary amine more than an aminoethyl 83 substituent, pKa^^ of HsN(CHg)gNHg (9*350 ) is greater than pKa^^ of Me2N(CH2)2M2 ( 9 *2Ao) ; whereas, the effects of the dimethyl- aminopropyl and aminopropyl substituents are approximately equal, 83 pKaMip of H2N(CH2)3EH2 is 9.870 and the pKa^^^ of Me2N(CH2)2UH2 is 9 *901. The larger decrease in the basicity of the primary ni­ trogen than the tertiary nitrogen in N,U-dimethylethylenediamine, as compared to the model monoamine compounds, may result from a steric interaction of a dimethylamine group with the water molecules hydrogen-bound to the three protons of the protonated primary ni­ trogen. A hypothetical solution one molar in dimethylethylamine, one molar in ethylamine, and one molar in hydronium ions enq)loyed as a model for the monoprotonated diamine (XVl) when n is infinitely large yields for the ratio CEMlPt]/[EMlP] a value of 0 .225* It 208 appears tiiat as n increases the ratio of basicities of the two diamine nitrogens approaches that value estiznated for very large n, see Figure 55* Plots of the micro pKa values against n, for protonation at the primary position. Figure $6, and for protona­ tion at the tertiary position. Figure 57, indicate that as n in­ creases the micro pKa approaches the pKa for ethylamine and dimethylethylamine respectively. The micro pKa values calculated for monoprotonated E,N-dimethyl-l,5 -pentanediamine are within ex­ perimental error of the pKa of the model compounds for infinitely large n. For an uj-dimethylamine-alhylamine with five or more methylene groups, n ^ 5j the nitrogen atoms are too distant to produce mutual inductive effects; and, the primary and tertiary nitrogen atoms of these diamines would exhibit properties that are independent of each other.

C. Thermodynamics and Kinetics of Imine Formation

1 . Equilibrium constants for carbinolamine-imine formation 110 a. MonofUnctional amines. In his discussion, C. Y. Yeh correlated the free energy of carbinolamine-imine formation from isobutyraldéhyde and methyl-, ethyl-, ^-propyl- and t-butylamine 111 with the Taft equation for a linear steric energy relationship.

110 . Ç. Y. Yeh, Effects of Substituents on the Formation and Con­ formation of Imine8 Derived from Isobutyraldéhyde and Sat­ urated Primary Alhyl^ines. Ph.D. dissertation. The Ohio State University, 196B.

111 . R. W. Taft, Jr., J. Amer. Chem. Soc., 7 5 , ^538 (1955). 209

10

9

pKa

8

7

oo

Figure 56, Plot of pKaj^^p and pKa^gp vs. n 210

10 .

pKa

oo n

Figure 57. Plot of pKa^^^t P^®M2t ® 211 112 eq.. 120. A plot of log K vs Taft's steric E_ values gave a “““ O

log|^ = 5Eg (120)

110 good straight line with 6 = 0 .$2 . Since n-propylamine and n- hutylamine, whose equilihrium constants are not significantly 110 different for the value reported for ethylamine (Alog K = O.l), did not correlate with eq. 120, it appears that the steric effects of imine formation of these three amines are similar. Thus, the amines used in this study, whose carbon analogs are ethyl-, n- propyl-, and n-butylamine are apparently not sterically dissimilar with respect to imine formation. The free energy of carbinol- 7 amine-imine formation for n-propyl-, methyl-, 5-methoxypropyl-,

2-methpxyethyl-, 2,2-dimethoxyethyl -, and 2,2,2-trifluoroethyl- 113 amine correlate with Taft's equation for polar effects, eq. 121.

log ÎL. = p* G* ( 121) Ko „ n o For the plot of log K vs or* with 15 amines Yeh reports a p* of

-1 .1 1 . For the six monofunctional amines of this study a plot of log K vs pKa demonstrates linearity. Figure 2 1 . The results indi­ cate that electron-withdrawing substituents discourage carbinol­ amine-imine formation. Of the carbinolamine-imine equilibrium

112. R. W. Taft, Jr. and M. M- Kreevoy, J. Amer. Chem. Soc., 7 9 > 4011 (195T).

115 • R . W . Taft, Jr., in Steric Effects in Organic Chemistry, M. S. Newman (ed.), John Wiley and Sons, Inc., New York, N.Y., 1956, Chap. 15. 212 constants reported only atout 7^ of the value results from car- binolamine formation. "Electron--withdrawing substituents would be e:^ected to discourage the transformation of the amino group, in which the nitrogen is probably approximately sp®-hybridized, to an imino group, in which the nitrogen is sp^-hybridized, and hence more electron withdrawing. The situation is somewhat similar to that found in the case of olefins, aldehydes, and ke­ tones, whose enthalpies and free energies of hydrogenation are //5b?ais made more negative by electron-withdrawing substituents. 110 The equilibrium constants, Table 2 9 , determined by Yeh and 5b reported by Hine et al. were employed in the free energy rela­ tionships and the kinetic expression for imine formation. This action results from the experimental limitations of this inves­ tigation. Solutions were not designed for equilibrium determina­ tions but for kinetic expediency; for example, a ratio of free amine concentration to aldehyde concentration of 2 to 1 was main­ tained in order that the excess amine may act as a buffer. For each equilibrium constant determined by Yeh approximately eight different ratios of amine and aldehyde concentration were employed; and also, two experimental methods (pH and spectrophotometry) were used.

b. Monoamines and unprotonated diamines. The equilibrium constants for carbinolamine-imine formation from isobutyraldéhyde and unprotonated H,H-dimethylethylenediamine, N,N-dimethyl-1,3- propane diamine, N ,N - dimethyl-1,4 -butane diamine, and 213

N ,N-dimethyl-1,5-pentanediamine were determined and correlated

with the Taft free energy relationship as were the monoamines,

Figure 3 7 • A plot of log K vs pKh yields a reasonably good

linear function, the small deviation of N,N-dimethyl-1,5-pentane-

diamine, log K of -0 .1, is considered experimental, Figure ^ 6 .

These free energy relationships indicate that for carbinolamine-

imine formation the monoamines and unprotonated diamines are not

chemically dissimilar.

c. Monoprotonated diamines. The equilibrium constants for

carbinolamine-imine formation from isobutyraldéhyde and the four

.monoprotonated diamines were determined. Since N,N-dimethyl-1 ,4 -

butanediamine and E-dimethyl-1 ,3 -pentanediamine correlate,

within experimental error with the linear free energy relation­

ships, Figures 36 and 57, these two monoprotonated diamines appear

to react in a manner similar to the monoamines and unprotonated

diamines. The large positive deviation for N,H-dimethylethylene-

diamirie and N,N-dimethyl-l,5 -propanediamine from this linear free

energy relationship is attributed to imidazolidinium ion and hexa- 102 hydropyr imi dinium ion formation respectively. K- W. îîarducy

has determined the equilibrium constant for imidazolidinium ion

formation from isobutyraldéhyde and monoprotonated N,N*-dimethyl-

ethylenediamine to be 25 M"^. Thé equilibrium constant calculated

for imidazolidinium ion formation from monoprotonated N,W-dimethyl-

ethylenediamine is in agreement with the value observed for its

structural isomer, '-dimethylethylenediamine; whereas, the 214 calculated value for K,N - dimet hyl-1,3-propane âiamine is slightly smaller.

2 . Equilibrium constants for carbinolamine formation

Equilibrium constants for the addition to the carbonyl group of isobutyraldéhyde of six monoamines and four unprotonated and monoprotonated diamines were determined. The values for the six monofunctional amines appear to be correlatable with their re­ spective basicities. Figure 2 2 . However, the plot of log vs pKa for all fourteen determinations of this investigation does not illustrate a correlation. Figure 3 8 . Nine points, within the dotted rectangle, appear randomly distributed. These results confirm previous conclusions that in this series of primary amines there is little sensitivity of the equilibrium constant to the basicity of the nucleophile^ i.e., polar substituents have similar effects on the stability of -H and -CHROH and conversely, -H and 3.14-116 -CHEOH do not have a large difference in polarity. Kallen 103 and Jencks report that steric requirements of the amine pri­ marily control the magnitude of the carbinolamine equilibrium

constants for the addition of formaldehyde to the amine. As

expected the equilibrium constant for addition of methylamine to

isobutyraldéhyde is larger than that for propylamine. Results for

ll4 . J. Hine and R. D. Weimar, Jr., J. Auer. Chem. Soc., 87, 5587(1965).

115.. W. p. Jencks, ibid., 81, 475 (1959).

116.. E. G. Sander and W. P. Jencks^ ibid., 90, 6l54 (1968). 215

2,2,2-trifluoroethylaiiiine and monoprotonated N,H-climethylethyl- enediamine significantly deviate from a mean value. The positive deviation for the monoprotonated diamine is believed to result from the experimental limitations of the stopped-flovj spectro­ photometer. The rate of change of aldehyde concentration with respect to change of time, dA/dt, for dehydration of the carbinol­ amine of monoprotonated N,N-dimethylethylenediamine was the largest observed in this study. During mixing time, the time required for the two feed solutions to traverse the cuvette’s optical path length, as much as 1C6^ of the aldehyde concentration could be converted into imine and thus render questionable the accuracy of extrapolation to zero time of the transmittance vs time plot.

For a qualitative description see Figure 69;. for a quantitative

evaluation of the sensitivity of the calculated values of Kg to

changes in the extrapolated value for the percent transmittance at zero time, tg, see Table 60 in the Appendix. The small car­ binolamine equilibrium constant for 2,2,2-trifluoroethylamine is believed to result from an experimental design and/or a small

polar effect. Because this amine has a small carbinolamine-imine ■

equilibrium constant a large excess of amine realtive to aldehyde

(10/1) was used to produce a readily observable change in aldehyde

concentration during the imine formation reaction. Thus, the

expression for calculation of Kq has a relatively large number in

the denominator. Experimental limitation will not permit solutions

of large but equal concentrations of amine and aldehyde, A > 0 .05, 216 to be measured accurately; because of the low percent trans­ mittance of the solutions, the increased ionic strength, and longer■ mixing time. The small equilibrium constant, Kq , could result from a polar effect which would not be observable for the other monofunctional amines of this study. Although the difference in polarity for -H and -CHROH is small, -CHROH should 113 be slightly more electron-withdrawing, and the Taft p for the reaction would have a small absolute value but would be less than zero (p = - 0.1 to -0 .4 ). Therefore, the strong electron- withdrawing ability of F3CCH2-, d* = -O.92, would cause the car­ binolamine equilibrium constant to be smaller than the values observed for the other amines, which possess only moderate electron-withdrawing substituents, o* = -0.4 to -K).l.

5. Rate constants for imine formation

a. The uncatalyzed,dehydration of carbinolamine. The ob­ served rate constants for the uncatalyzed dehydration of carbinol­ amine with the six monoamines and isobutyraldéhyde correlated with

Taft's equation for polar effects. Figure 25, yields a p = -2 .5 7 *

If the point for 2 ,2 ,2-trifluoroethylamine is eliminated a p =

-I.79 is obtained. In either case the p is negative and is in agreement with the mechanism previously proposed ( see Chapter II) eq. 122 . A negative p dictates that an electron-withdrawing OH I slow ^ i-ETv^ fast . i-Pr-C-N-R + OH" imine + HgO (122) U « 217 substituent "will destabilize the transition state and thus increase the activation energy for imine formation.

The slope of a plot of log ko pKa, Figure 2h, serves as a comparison of the difference in polarity between reactant and transition state for imine formation to the difference in polarity between reactant and product for protonation of the amine, eq. 123.

slope = polarity difference ( C-transition state) (125) polarity difference (BÎIH2-EIIB3+}

This slope can be used to describe the geometry of the transition state. The difference ' in polarity of carbinolamine and amine was discussed in the preceding paragraph and is considered small. If the transition state for the imine formation reaction is a com­ pletely protonated imine, the slope, eq. 123, should approximate the sum of the slopes of similar plots for the two reactions eq.

124 and 1 25 . The protonation reaction, eq. 125, should have a

ENHs + i-PrCHO i-PrCH=KR (124)

i-PrCH=M + Ifeo'*' ;-- ^ i-PrCH=NHR‘*‘ + HsO (125) slope near that of the analogous protonation of KNHg, which is 1.0 by definition. Since the slope for the imine formation reaction sb is 0.256, the slope in eq. 123 would approximate 1.26 if the transition state is an iminium ion. A more rigorous estimate (of the polarity of the iminium ion) considering the effect of the partial distribution of the charge over the carbon-nitrogen double bond, although not quantitative, would result in a slightly smaller 2 1 8 estimate for the slope in eq.. 125 than 1 .2 6 . The observed slope of 0.51 qualitatively indicates that the transition state is located on the reaction coordinate approximately mid-way between the reactants and products. The transition state (XVII) can be described as possessing about one-half of a positive charge on the nitrogen atom and one-half of a negative charge on the oxygen atom « -0.5 HO i-Pr^i y H ^ 6 — H 40.5 H/

XVII

The equilibrium constants, Kq, are used in the calculation of the rate constants; and because of the experimental limitations small uncertainties in the determination of result in small uncertainties in the calculated value of the rate constants. If the observed Kq is smaller than the actual Kq , then the calculated value of kg is larger than the real kg. Consideration of the pro­ duct, K^kg, can minimize this uncertainty. A plot of log K^ko vs pKa yields a very good linear correlation for the six monoamines,

Figure 2J. The results with the unprotonated bifunctional amines can be similarly correlated. A plot of log KqK v s pKa with the ten amines for which the uncatalyzed rate constant is experimentally observable yields a good linear relationship, which is used to cal­ culate the value of K^pgk^^ for the monoprotonated diamines. Figure k$. 219

b. Specific acid-catalysis of imine formation. The specific acid-catalysis rate constants determined for the dehydration of the 97 carbinolamine S of isobutyraldéhyde and methyl^, 2-methoxyethyl-,

2,2 -dimethoxyethyl-, 3 -methoxypropyl-, and 2,2,2 -trifluoroethyl- amine are not correlatable -with the pKa of the respective amine.

Figure 2 6 . A similar plot including the observed rate constants for the monoprotonated diamines demonstrates linearity with a slope of about 0.8, Figure 50. This slope, utilized as described in the preceding paragraphs, indicates that the transition state is positioned on the reaction coordinate much closer to products than reactants, and that the geometry of the transition state can be approximated by XVIII. The points for methylamine and 2 ,2 ,2 -tri- fluoroethy lamine deviate significantly from this linear function. + 0.2

i - Pr^i Æ + 0.8

XVIII

Part of this deviation may result from the uncertainties in deter­ mining the equilibrium constant for carbinolamine formation and can be partially eliminated by plotting log vs pKa, Figure 5 1 *

The remaining positive deviation for 2 ,2,2-trifluoroethylamine may result from a contribution from general acid-catalysis. Since

general acid-catalysis -was not expected to occur, the experiments

of this investigation were not designed to test or evaluate this

catalysis, i.e., -varying total amine concentration at constant pH; 220 and it is possible that a small (k < 10^) general acid-catalysis tenu has remained hidden within the large (k > 10^) specific acid-catalysis rate constant. General acid-catalysis will be given further consideration in the next paragraph. Estimates of the specific acid-catalysis rate constants for dehydration of carbinolamines of the unprotonated diamines can be obtained from the linear function of Figure 5 1 * The estimated rate constants are used to evaluate Since the contribution of specific acid-catalysis to kg^ is very small (see Table 5l), the deviation for methylamine and 2 ,2,2 -tr if luor oet hylamine from linearity in

Figure 51 is considered of little relevance and these points are not used in the calculation of the slope.

c. . General acid-base catalysis. The observed catalysis for the dehydration of the carbinolamine of B ,K-dimethylethylenediamine was experimentally evaluated as general acid catalysis, which is a function of the product of the genral-acid [bh] and carbinolamine

[c] concentrations. This observed catalysis is kinetically indis­ tinguishable frcm a form of general base catalysis which is a function of the product of the general-base [b] and protonated carbinolamine Cch] concentrations. Both possibilities are con­ sidered. 117 . Several researchers have reported that the dehydration of carbinolamine proceeds almost entirely through hydronium ion

117. K. Koehler, ¥. Sandstrom, and E. H. Cordes, J. Amer. Chem. Soc., 8 6 , 24-15 (19^) and references therein. 221 catalysis (Br^nsted Of = 0 .7 -0 .9 )- A small contribution, from general acid-catalysis may occur for the dehydration of car- hinolamine8 of weakly basic amines, pKa < 5 -0 . As the basicity of the amine increases the dehydration reaction becomes less susceptible to general acid-catalysis. Similar results are re­ ported for Schiff base formation under reaction conditions in 118 which carbinolamine formation is rate-determining. From work conducted in our laboratory on the hydrolysis of i s obutyll dene - methylamine, it is concluded that this carbinolamine is probably not subject to a significant contribution from general acid- 97 catalysis, i.e., none was observed. It was therefore antici­ pated that general acid-catalysis would not be important in the dehydration of the carb inolamine s of this investigation. Perhaps with 2,2,2-tr if luoroethy lamine and monoprotonated N,N-dimethyl- ethylenediamine, whose pKa values are 5*22 and 5-52 respectively, a small contribution from general acid-catalysis may occur. How­ ever, experiments designed to test this catalysis for monoprotonated

N,N-dimethylethylenediamine revealed none. A positive deviation for the specific acid-catalysis rate constant for 2,2,2-tr if luoroethyl- amine from the linear function of Figure 51 gives some indication of a small contribution from general acid catalysis.

Il8. C. G. Swain and J. C. Worosz, Tetrahedron Lett., $6, 3199 (1965). 222

Unexpectedly, general acid-catalysis was observed for the dehydration of the carbinolamine of isobutyraldéhyde and N,N- dimethylethylenediamine. This carbinolamine is composed of a rather basic amine (pKa = $.24) and the rationalization of the observed catalysis does not appear obvious. General acid- catalysis dictates that the transition state contain a molecule of the monoprotonated amine with the carbinolamine. The other diamines studied in this investigation do not function as general acid cata­ lysts; yet, their basicities are not sufficiently larger than that of N,N-dimethylethylenediamine as to render them inefficient acid catalysts. The only significant distinction appears to be the distance between the two basic nitrogen atoms of each diamine.

Perhaps a molecule of K,IT-dimethylethylenediamine remains with the carbinolamine in the transition state as a result of a type of 119 hydrogen bonding. Jencks has stated that "although monofunc­ tional hydrogen bonds have no significant stability in aqueous solution, polyfunctional hydrogen bonding may contribute signifi­ cantly to the driving force for intermolecular interaction in water."

For example the large increase in the solubility of acetyltetra- glycine ethyl ester XIX upon addition of urea is attributed to

p Cj 0 ClfeCNHCHaC CNHCIfe) ‘3COC2H5

XIX

119 • W, p. Jencks, Catalysis in Chemistry and Enzymology, bp. cit., pp. 525-550. 225 polyfunctional hydrogen bonding of the possible structure XX.

For the interaction of monoprotonated If-dimethylethylenediamine

! I H H

and the corresponding unprotonated carbinolamine, several possible polyfunctional hydrogen bonding combinations can be envisioned, of

•which t'wo are sho'wn. To simplify these combinations the dimethylethylenediamine is considered protonated at the primary nitrogen. Structure. XXI. does not mechanistically aid in iminium

\

/ H-Q

XXI XXII

ion formation but could hold the general acid near the carbinol­

amine so that it may convert to XXII. However, it is not apparent

why a seven-member polyfunctional hydrogen-bonded ring with H,N-

dimethylethylenediamine ( for structure XXII) is significantly more

stable than the possible eight-membered ring with N,H-dimethyl-l,5- propanediamine. 224 26,27,28 Several researchers have observed general-base catalysis for the dehydration of carbinolamine of weakly basic amines, pKa < 5*0 and have proposed a mechanism involving the removal of the nitrogen-boimd proton with expulsion of "OH as the sp^ hybridized bond is formed. The basicity of tertiary mono­ protonated H-dimethylethylenediamine is much greater (pKa =

6.86) than other amines for which base catalysis is observed for the dehydration reaction. It appears that this proposed mechanism can not account for the observed catalysis (IF it is base catalysis)

However, since the other diamines did not exhibit this catalysis a mechanism Involving intramolecular acid and intermolecular general- base catalysis (XHIl) is proposed.

H. 0“**H— (CHs)2

''B XXIII

In Chapter VI, the conclusion, experiments designed to eluci­ date this apparently observed catalysis are suggested for future investigations ; however the current rationalization of the observed catalysis with monoprotonated H-dimethylethylenediamine remains uncertain.

_d. Internal acid-catalysis. The results of this investiga­ tion demonstrate that the rates of dehydration of carbinolamines ^ 225 of monoprotonated n-dimethylethylenediamine and monoprotonated

IT,IT-dimethyl-1 ,5 -propanediamine are significantly faster than predicted with Taft’s equation, Figure 5 2 . This increase is at­ tributed to internal acid-catalysis. The relative rate ratios, i.e., the observed rate constants for imine formation, KQ(kpQ + kpj^), divided by the theoretical value, Egko, can be compared to rates of ring formation and relative acidities of the acid cata­ lysts, Table 5 3 - The observed rate acceleration does not correlate with the relative acidities of the amines. The relative rates of ring formation were obtained from rate of formation of cyclic sulfonium salts by chloride displacement from the compounds, 120 RS(CHg)^C1 . These relative rates are similar to those reported 121 for lactonization of the anions of bromo-acids. The observed relative rate ratio for internal catalysis of imine formation cor­ relates reasonably well with those reported rates of cyclization.

A plot of the relative rate ratios vs rate of cyclization yields a line of slope 1.0 for the carbinolamine of monoprotonated. W,II- dimethylethylenediamine and monoprotonated IT ,IT-dimethyl-1 ,5 - propanediamine. The point for IT,IT-dimethyl-1 ,4 -butanediamine deviates significantly in that its relative rate ratio is located

120. G. M. Bennett, et al., J. Chem. Soc., 1958 (1938); 6l5 (1958).

121. G. Salomon, Helv. Chim. Acta, l6 , 156l (1953); 17, 851 (1934). TABLE 53

Correlation of Internal Acid-Catalysis

Ring Relative Relative Rela-tive ^ size rate ratio rate of ring acidities formation

5 — — 19,600 —

6 3,500 260 4-0

7 50 3.57 2.0

8 1.5 1 1.0

9 1.0 —— 1.-

a Determined from 22T

1.1 powers of ten below the line. Data for comparison with rate of dehydration of the carbinolamine of IT,IT-dimethyl-1,5-pentane­ diamine is not available for this proton transfer would require that the "intermediate" or transition state resemble a nine- membered ring. Rings of size between 8-l4 members, mostly carbon IP? atoms, are very difficult to form. These results indicate that internal catalysis does not significantly accelerate the rate of dehydration of carbinolamines of monoprotonated -dimethyl-1 ,4- butane diamine and monoprotonated N,IT-dimethyl-1,5 -pentanediamine.

It appears that the energetics of ring formation are similar to those for the internal catalysis of imine formation; and thus a transition state involving internal proton transfer can be con­

sidered, XXIV.

‘ 0"“ H RMeg ' I , £-Br-Cv (0112)2-3

I H

XXIV

122. V. Prelog, J. Chem. Soc., 420 (1950). 2 2 8

The observed rate term kg ^ contains a term for specific acid catalysis of the unprotonated carbinolamine with the term for intramolecular acid catalysis of the protonated carbinolamine, eq. 105. The term for specific acid catalysis was evaluated by assuming that the rate of this catalysis for carbinolamines of monoamines and monoprotonated diamines is linearly correlatable to that for the carbinolamines of the unprotonated diamines; how­ ever, this correlation is invalid if specific acid catalysis occurs via a transition state (XXV) in which intramolecular base catalysis is also involved. Thus the observed rate increase for -dimethyl-

^ 0 ' '

3QC7 ethylene diamine and N -dimethyl-1 -propanediamine may result from the concerted mechanism of intramolecular base catalysis and bimole cular specific acid catalysis. This and previous in- 6 vestigations indicate that the rate-determining step for the dehydration of carbinolamines of amines of similar basicity is the formation of an iminium ion and not the removal of the nitrogen bound proton as illustrated for mechanism XXV. Also for XXV the relative rate ratio observed for E,N-dimethyl-l,3-propanediamine is not expected to be substantially smaller than that for 229 n -dimethylethylene diamines because a five and four-membered

ring, respectively, must be formed in the transition state.

Therefore, the observed rate increase is attributed to intra­ molecular acid catalysis. The estimated ring size for the

transition state does not include the. hydrogen bond which would

form as a result of the proton transfer. These bonds usually prefer to be linear and probably cause the ring in which hydrogen bonding occurs to have an elongated side with a N-H-N angle of

about 180°. Thus the thermodynamics of the ring of XXV with

N ,11 -dimethylethylenediamine would approximate the thermodynamics

of a four membered ring more so than that of a five membered ring. CHAPTER 71

CONCLUSION

The kinetics and themodynamics of the addition of a series of primary alhylamines to isobutyraldéhyde -was studied in neutral and basic solutions. The two step reaction involves the rapid formation of a tetrahedral carbinolamine intermediate and the sub­

sequent, rate-determining dehydration of this intermediate to form an imine. A stopped-flow spectrophotometer was employed to deter­ mine the equilibrium constants for carbinolamine (Kg) and imine formation and the rate constants for carbinolamine dehydra­ tion.

The values of the carbinolamine equilibrium constants deter­ mined for six primary alkylamine and four unprotonated and mono­ protonated. u)-dimethylamino-alkylamines are not correlatable with the polarity or basicity of the amines, i.e. Taft's equation is not applicable. The size of Kg appears to be moderately dependent of steric factors.

The values of the carbinolamine-imine (Kg + Kg^ = Kj) equili­ brium constants determined with fourteen primary amines are

correlatable with a Taft free energy relationship (p = -l.ll).

The constant Kg constitutes less than ten percent of the value of

Kj. The observation that electron donating substituents (o* < O.O)

2)0 251 enhance imine formation appears to result because sp^ nitrogen is more electronegative than sp^ nitrogen. The positive deviation from this linear relationship observed with monoprotonated dimethylethylenediamine and monoprotonated R,IT-dimeth yl-1,5~ propane diamine is attributed to imidazolodinium ion and hexahydro- pyr imi dinium ion formation, respectively.

The rate constants for uncatalyzed dehydration of the carbin­ olamines of six primary monoamines and four unprotonated diamines are correlatable with a Taft free energy relationship (p = -1 .8) and indicate that the rate-determining step of this reaction in­ volves the ejection of a hydroxide ion with formation of the iminium ion. The observed specific acid-catalysis rate constants for imine formation are also correlated with a linear free energy relationship. Intramolecular acid catalysis was observed for the dehydration of carbinolamine s of monoprotonated IT ,IT-dimethyl­ ethylenediamine and monoprotonated IT,IT-dimethyl-1 ,5 -propane­ diamine (monoprotonation at the tertiary nitrogen). The relative ra,tio of the rate constants, that is the observed rate constant divided by the theoretical rate constant determined from a linear free energy relationship^ for tertiary monoprotonated 1T,1T- dimetlylethylenediamine, R,IT-dimethyl-1,3 -propanediamine, R,R- dimethyl-l,4-butanediamine, and R,R-dimethyl-l,5 -pentanediamine are 5 ,500, 50, 1 .5 , and 1.0 respectively. The magnitude of these relative rate constants are compared to the rates of ring forma­ tion. 2 5 2

The micro“pKa values for the œ-dimethylamino-alkylamines,

Me2Îî(CH2)jj]ra2, were determined by observing the downfield shift of the nmr absorption peak for the N-methyl protons with increasing concentration of protonated amine. It was determined that the monoprotonated diamines contained the proton on the tertiary amine site 58, 29, 5^, and 21 percent of the time as n increased from

2, 3, and 5 respectively.

The hydration of isobutyraldéhyde at 35° in aqueous solution was studied spectrophotometrically at pH 6.5-7 *8 and found subject to general acid-general base catalysis.

The results of this investigation should give impetus to future research in the specific area of the intramolecular catalysis and the general acid-base catalysis of the dehydration of carbinol­ amines of the diamines.

■ Intramolecular Catalysis. Experiments designed to distinguish between intramolecular acid catalysis (XXIV) and intramolecular base-specific acid catalysis (3QCV) should be of interest. If ca­ talysis results via mechanism XXIV both cis and trans (n,N- dimethyl-l,2-cyclopentanediamine should exhibit intramolecular catalysis with the cis compound being the more efficient catalysis.

If catalysis results via mechanism XXV the cis-diamine should ex­ hibit intramolecular catalysis and the trans compound would be inefficient as an intramolecular catalysis. 255

. Greneral Acid-Base Catalysis. Experiments designed to confirm and identify this catalysis should be considered. If the dehy­

dration of the carbinolamine of isobutyraldéhyde and dimethyl­

ethylene diamine exhibits this catalysis with (Me2E)s( when n is 2 and perhaps 5 but not when n is 4 or 5, perhaps the tran­

sition state resembles XXI or XXII • However if catalysis is ob­

served with all four of the tetramethyldiamines and with a series

of monofunctional tertiary amines (EIîR'R*') it would appear that

catalysis occurs via mechanism XXIII and further experimentation with the carbinolamines of É -dimethyl-1,5-propanediamine, N,N-

dimethyl-1 ,4 -butanediamine and H,N-dimethyl-1,5-pentanediamine would be necessary to verify the role of the carbinolamine as an

internal acid catalyst. APPENDIX

23^ 235 Symbols Frequently TJsed

General

A = real concentration of isobutyraldéhyde calcu­ lated spectrophotometrically using £ R = initial concentration of isobutyraldéhyde before a kinetic run

A_ — total concentration of isobutyraldéhyde V *o + “o B = concentration of general base

[BH ] = [b h ] = concentration of general acid

C = concentration of unprotonated carbinolamine

[CH ] = [CH] = concentration of monoprotonated carbinolamine

[CTI — total carbinolamine concentration, C + [c h ]

[EMT] = total amine concentration

[eMUP] = concentration of unprotonated amine

[EMP] = concentration of protonated monofunctional amine

[eMIP] = concentration of monoprotonated diamine

[EMlPp] = concentration of monoprotonated unsymmetrical dimethylalkyl diamine, protonated on the primary amine

[EMlPt] = concentration of monoprotonated unsymmetrical dimethyl alkyl diamine, protonated on the tertiary amine

[eM2P] = concentration of diprotonated diamine

£ = apparent extinction coefficient of iso­ butyraldéhyde, 14*5 M cm 2 3 6 € — real extinction coefficient of isobutyr­ aldéhyde, 20.86 cm""^

H = concentration of the hydrate of isobutyraldéhyde

[h ] = hydrogen ion concentration

I = concentration of imine

= concentration of monoprotonated imine

M = same as EMUP

= [EMT] at t^ for a kinetic run

[m H ] = same as [eMP] or [e MIP]

R = concentration of imidazolidinium ion

^ = number of observations recorded

run 7^ = the run number is the internal accounting number used to relate each photo and solution. The number left of the decimal refers to the page in the laboratory notebook on which the solution for that run is described. The number to the right of the decimal is for photo identi­ fication.

Vj = rate of dehydration of the hydrate of iso­ butyraldéhyde calculated assuming that the reaction is not reversible

Vj = rate of formation of imine calculated assuming that the reaction is not reversible

For equilibrium constants;

Kç = real equilibrium constant for carbinol­ amine formation, C/AM

K = apparent equilibrium constant for carbinol- ^ amine formation, C/(A+H)M

K = equilibrium constant for imine formation from carbinolamine, l/C or (I + [lH"^]/(C + [CH+]) Cobs = experimentally observed equilibrium constant for carbinolamine formation, from monoamines same as K^, from diamines (C + Cc h ’^])/a [eMT]

Kccal " average

%lcal = *Iobs Kçp = equilibrium constant for monoprotonated carbinolamine from monoprotonated diamines, c h +/a (m h +)

K = same as K _ only the apparent value where ^ A + H is used instead of A

K_ = equilibrium for isobutyraldéhyde and its hydrate, A/H

Kj. = the sum of the real equilibrium constants for carbinolamine and imine, + K^, (l + C)/AM

K = the apparent counterpart of K , (I + C)/(A + H)M la = the equilibrium constant for formation of imine from aldehyde and amine, 1/a[EMUP]1

K = the apparent counterpart of K „ , replace ^ A with A + H

K ■ = the equilibrium constants for imine formation for each kinetic run (I + Ci h ’^])/a [EMT]

Kimp “ the equilibrium constant for formation of monoprotonated imine from aldehyde and monoprotonated diamine, [IH’^]/a Em H’*‘D

^IMPE ~ equilibrium constants for formation of monoprotonated imine from aldehyde and tertiary protonated, monoprotonated diamine, [lH+]/A[EMlPt]

Kjobs ” experimentally observed equilibrium constant for imine plus carbinolamine formation, from monoamines same as K_, from diamines, (I + C)/AM + ([IH+] 4 [CH+])/A(MH+) 238 K^p — equilibrium constant for formation of mono­ protonated imine and carbinolamine from monoprotonated diamine, + [CH'*'3 )/A(MH'*‘)

K _ = the apparent counterpart of K ([IH+] + [CH+])/(A + H)(MH+) ^

K^PE — equilibrium constants for formation of monoprotonated imine and carbinolamine from tertiary protonated, monoprotonated diamine, ([IH+] + [CH+])/A[EMlPt]

K^p e ~ equilibrium constants for formation of monoprotonated carbinolamine from the alde­ hyde and tertiary protonated, monoprotonated diamines, [CH"^]/ A[EMlPt]

= equilibrium constant for formation of imidazolidinium ion from tertiary protonated, monoprotonated diamines, R/A[EMlPt]

Acidity constants

K = the dissociation constant for protonated monoamines

Ka^^ = the dissociation constant for monoprotonated, at the tertiary amine, carbinolamine formed from unsymmetrical dimethylalkyldiamines and isobutyraldéhyde

Ka^2 “ the dissociation constant for diprotonated, from the secondary amine, carbinolamine formed from unsymmetrical dimethylalkyl­ diamines and isobutyraldéhyde.

Ka„- = the dissociation constant for monoprotonated Ml j - _• diamine

. = the dissociation constant for monoprotonated unsymmetrical dimethylalkyldiamines with the acidic proton located on the tertiary amine 239 “ the dissociation constant for monoprotonated unsymmetrical diamethylalkyldiamines with the acidic proton located on the primary amine

Ka^g — the dissociation constant for the loss of one proton from diprotonated diamine

Kaij^gt “ the dissociation constant for the loss of the proton bound to the tertiary amine of a diprotonated unsymmetrical dimethylalkyl— diamine

Ka^. — the dissociation constant for the loss of the proton bound to the primary amine of a diprotonated unsymmetrical dimethylalkyl- diamine

Rate constants

k^ =7 the second-order rate constant for the addi­ tion of a primary amine to isobutyraldéhyde to form a carbinolamine

k ^ = the first-order rate constant for the collapse of a carbinolamine into its primary amine and isobutyraldéhyde

kg = the first-order rate constant for the dehydra­ tion of a carbinolamine into an imine

k g = the pseudo first-order rate constant for the hydration of imine into carbinolamine

^2obs ~ experimentally observed pseudo first- order rate constant kg 24o k —a first-order rate constant for the uncata­ lyzed contribution to , for monoamines 2 obs kjj = a second-order rate constant for the specific acid-catalyzed contribution to k„ . for 2 obs monoamines

kgjj = a second-order rate constant for the general acid-catalyzed contribution to k_ , for , 2obs monoamines

k = a first-order rate constant for the un­ catalyzed contribution to k_ . for an 2 0 0 & unprotonated carbinolamine of a diamine

k^g = a second-order rate constant for the gener­ al acid-base -catalyzed contribution to ^2obs unprotonated carbinolamine of a diamine

k ^ = a second-order rate constant for the speci­ fic acid-catalyzed contribution to k_ , 2 obs for an unprotonated carbinolamine of a diamine

k = a first-order rate constant for the un- catalyzed contribution to ^2obs &

monoprotonated carbinolamine of a diamine

k . = a first-order rate constant for the intra- molecularly catalyzed contribution to ^2obs for a monoprotonated carbinolamine of a diamine

k , = a second-order rate constant for the specific acid-catalyzed contribution to for a monoprotonated carbinolamine of a diamine k_T^ ~ ® first-order rate constant which is a sum of + kuh"acl

k = k, - + k - for hydration-dehydration of AO TiO dO isobutyraldéhyde 241 k.„ — le „ + k for hydration-déhydration of isobutyraldéhyde k = k, + k _ for hydration-dehydration of AUn iiUfi aUn . . . - _ . isobutyraldéhyde

k._ = k, _ + kj_ for hydration-dehydration of A^w fit) Cl iH. • « , .# « « , isobutyraldéhyde k = k + k for hydration-dehydration of ABH hBH dBH isobutyraldéhyde

kj^Q — a first-order rate constant for the uncatalyzed hydration of isobutyraldéhyde

kj^ = second-order rate constant for the specific acid-catalyzed hydration of isobutyraldéhyde kj^Qjj = second-order rate constant for the specific base-catalyzed hydration of isobutyraldéhyde

kj^g = second-order r^te constant for the general base-catalyzed hydration of isobutyraldéhyde khBH ~ second-order rate constant for the general acid-catalyzed hydration of isobutyraldéhyde

k ^ = where X is 0, H, OH, B, or BH, similar to a rate constant defined as k, , above, except it is for the dehydration or the hydrate of iso­ butyraldéhyde. 242

TABLE 54 Major Infrared Bands of the Diamines

2 3 4 5 3360(m) 3360(m) 3370(m) 3360(m) 3280(m) 3280(m) 3290(m) 3280(m)

2975 2970 2950 2940 2940 2940 2870 f . 2860 , . 2860 (s) 2860 (s) 2825 2820 2825 2820 2780 2720 2775 2770

1580(m) IS80(m) l600(m) 159 0(m) 1450(s) I450(s) I450(s) 1450(b) I340(m) l360(m) I 370(m) l360(m) 126o(m) 1255(m) I260(m) 126o(m) llOO(m) ll60(m) lllO(m) ll60(m) 1050(s) llOO(m) 1065(m) 1040(m) 850(m) 1040(s) 1050(s) 840(m) 770(m) 880(m) 850(3) 820(m) 850(s) 790(m) 770(m) 820(s) 730(m) 750(a) 245 TABLE 55 Example of Kinetic Run for Hydration of Isobutyraldéhyde

g Optical density Time, sec. Function —

1.664 10.0 0.195

1.616 20.0 0.312

1.570 30.0 0.439

1.532 40.0 0.557

1.498 50.0 0.682

1.467 60.0 0.799

1.440 70.0 0.921

1.418 80.0 1.031

— The function is InCAg/K^CA—A^) ] from eq.28 where “ 2-333 and optical density measured at equilibrium is 1.230. 2kh

1.0

1 <

Q < 3

0.5

slope — 0.0120 0.3

3010 50 70 'time, sec

Figure 58. Plot of time vs. In Ag/Kp (A-A^) for Hydration of Isobutyraldéhyde 245 Examples of nmr datai

A copy of five nmr spectra vrith a sweep width of

100 Hz of the 0,-methyl protons of trimethyl amine illus­ trate the linear downfield shift with respect to degree of protonation. Figure 59 •

TABLE 56

Chemical Shift of Protonated Trimethyl amine

Peak Peak % Free number location— trimethylamine &

1 2.25 100

2 2.42 75

3 2.59 50

4 2.76 25

5 i .93 20

— External T.M.S. as standard for this determination 4 f

Y V i 1 I j I 5.0 4.0 3.0 2.0 1.0

Figure 59* NMR Spectra of Protonated Trimethylamine in Aqueous Acid Solutions Described in Table $6. ON 2k7

A copy of nine nmr spectra with a sweep width of 100 Hz of the a-methyl protons of N-dimethylethylene- diamine, each with a different acid concentration illustrates the relationship between chemical shift and degree of protonation. Figure 60.

TABLE 57

Chemical Shift of Protonated N,N-Dimethylenediamine

Peak Peak % Free number location — Diamine S

a 2.15 100

b 2.22 87.5

c 2.30 75.0

d 2.38 62.5

e 2,46 50.0

f 2.58 37.5

g 2.70 25.0

h 2.82 12.5

i 2.92 2.0

— External T.M.S. as standard for this determination 6.0 4.05.0 2.0 1.0 Figure 60. NMR Spectra of Protonated W,N-Dimethylethylenediamine in Aqueous Acid Solutions Described in Table 57. 24g Figure 6l

Example of kinetic runs made with the stoppedr flow spectrophotometer. These two examples illustrate the change in optical density and reactant concentration with respect to time. The reaction was followed nearly to completion.

The reaction of N,N-dimethylethylenediamine and isobutyraldéhyde - run number 210.803 is used for this example.

Kimc^I =46.23 m"^ [e MT] = 0.04095 M

^Ccal = 2.08 = 0.00988 M pH = 11.23+0.03 k ^, = 5.89 sec"^ 2 obs. + H = 0.02663 M 0 0 “ Figure 62 - The observed transmittance vs time plot> obtained from stopped-flow photo

Figure 63 - Plot of f(A) vs time where ^2obs^ ” f(A)» that is the right side of expression

Figure 64 — Plot of calculated concentration of aldehyde * hydrate, carb nolamine and imine vs time. 250

Fig. 61.— A copy of the photo for run 310.801-310,803. 1) Run .801 time 0.1 sec/unit* 2) Run .8 02 time 0.2 sec/unit; 3) Run .803 time 0 .5. sec/unit.

TABLE 58 Data for Run Number 210.803 c d Time^ Aldehyde Hydrate^ Carbinolaminer- Iminer- sec M M M M

0.18 0.0170 0.0073 0.00137 0.00094 0.44 0.0159 0.0068 0.00122 0.0027 0.96 0.0139 0.0060 0.00099 0.0057 1.47 0.0127 0.0055 0.00086 0.0076 1.99 0.0120 0.0051 0.00078 0.0088 2.50 0.0114 0.0049 0.00073 0.0096 3.02 0.0110 0.0047 0.00069 0.0102 3.53 0.0108 0.0046 0.00066 0.0106 3 b — Corrected for photographic parallax. — Calculated assuming IC is .established. — Calculated assuming K is established. — Calculated as I = Aj - H - C — A. 9.T3«t-0l -

S.JÏ7E-0I -

2 . 9 a i E - 0 t transmittance

2.604E-01

2.22TE*01 -

I . 8 S I E - 0 E llllllIllilllltlllllllllllllllltlllUnitlltttllllllllltlUlllillllllllllltltlltlllltttllltlliniKtl

NPXPs 0 0.0 l.422e-0l I.404E :0 2.?7aE Ct) 03 l.TIU 30 time

Figure 62. The observed transmittance vs time plot, obtained from stopped-flow H photo. / 2.CbJ( Cl

I.6ME 01

t.?3TE 01 - f(A)

S.243E 30

4.I11E CO -

-l.aiTE-02 IIIIIFII111111111111111111IIIII1111111111II111111111111III 11111111IIII11111MII111II111111111111111II HPNP* 0 0.0 T.422E-CI l.4d*E vO 2.226C 00 2.0«.'IE 00 1.21IÇ >0 time

Figure 65. Plot of f(A) vs time where kaobs ^ = f( A), that is the right side fO of ejqiression. . I.T8K-02 -

Aldehyde

t.0a6t-Qi

Imine concentration

Hydrate

/o Carbinolamine 3.32IE-0* IIIIMIIIIinillllllllillinillllMIIIIIIIIUIIIIItllltllllUMMIMIIIinilllllllllllllllilinilllll hPNPa 0 0.0 I.48«E CO 2.226C 00 2.<)blr 00 3.TIIC 10 time g Figure 64. Plot of calculated concentration of aldehyde, hydrate, carbinolamine and imine vs time. 254 Figure 6j?

Example of kinetic runs made with the stopped- flow spectrophotometer. The reaction of N, N—dimethyl— ethylenediamine and isobutyraldéhyde.

K ^ c a l m ”^ a + « = 0.02663

Kçcal 4.94 [EMT] = 0.0851

pH = 7-83^:0.07 = 0.0092

H — 0.0080 M ^2obs ~ 29-3 sec ^

Figure 66 - observed transmittance vs time plot, obtained from stopped-flow photo

Figure 67 - plot of f(A) vS time where k , t = f(A), that is the right side of equation

Figure 68 — plot of calculated concentrations of aldehyde, hydrate, carbinolamine, and imine vs time. 255

o> 0 a - *T CO 1I % u H

Time Fig. 65.— A copy of the photo for runs 210.661-210.664. 1 ) Run .661-.663 time 0.01 sec/units 2) Run .664 time 0.02 sec/units

TABLE 59 Data for Run Number 210.664

Time^ Aldehyde Hydrate^ Carbinolamine^ ünine^ sec M M M M

0.018 0.0131 0.0080 0.0051 0.00047 0.028 0.0124 0.0080 0.0049 0.0013 0.038 0.0118 0.0080 0.0046 0.0022 0.049 0.0114 0.0080 0.0044 0 .00^9 0.059 0.0110 0.0080 0.0042 0.0035 0.079 0.0103 0.0080 0.0039 0.0044 0.100 0.0099 0.0080 0.0037 0.0050 0.121 0.0096 0.0080 0.0036 0.0055 3L b — Corrected for photographic paralax. ~ Calculated assuming dH/dt = 0 .0 . — Calculated assuming K is estab­ lished. "" Calculated as I = A„ — H — C - A. . «.ISIE-Ol -

3.asr£-ci -

3.»93t-31 - transmittance

3.30ilE-0l -

3.00&E-U1 -

Z.TUE'Ul lllllllinilllllllllllllltllltlllllllllllillllllllllllllllllllllllllltlllllltlltllllttlllllltlllltlll NPNP- 0 (i.O P.53St-0£ S.0TU.-C2 7.60bP-02 1.01*;-0I 1.2r.6E-)I time Figure 6 6 . Observed transmittance vs time plot, obtained from stopped-flov photo. ON J.S94E 00 -

2.304E CO - f(A)

1.659É CO -

t.O U c 00

3.6B »t-0l lltlllllllllltFIIIIIIIIIIIIIIIIIIIIIIIIiltillltltlllllllllilllllllttlllllllllllltlKllllillltlltllltl hPNP- 0 0.0 2.S3SE-0? 0.07U-C2 T.600E-32 I.QI*r-OI I.2WL-3I time Figure 67. Plot of f(A) vs time where kgobs^ = f(A), that is the right side of equation. I.ITOt-07

1.IJ2E-P2 - Aldehyde

a . W l k - 0 3 - ♦ ♦ ♦ ♦ Hydrate . e . ■ . .0—--- 0---- e— 0 ■ concentration

6.S»9k-CJ -

Carbinolamine Imine

I. G Ü 0 E - 0 I 11111 111111 1111111111i11111 111111111111111II11111111111111111111111111111II11 It 1111111111M1111111111 NPNP" 2 0.0 2 . 0 3 0 E - C ? S.CTIE-C2 r.60bE-C? I.Qlbr-Ol l.2bftE-31 time

Figure 68. Plot of calculated concentrations of aldehyde, hydrate, carbinolamine, CO and imine vs time. 259

An illustration of the experimental observations for determination of is given in Table 60.

Symbolism:

%T. — initial percent transmittance for aldehyde solution before reaction. This was deter­ mined by mixing the aldehyde solution with water (1/1) in the stopped-flow.

%T = extrapolated value of percent transmittance at t f i.e., after K has been established*

D = distance in cm between %T(, - % T . on the photo. Photo measures about 6.^ cm from top to bottom.

Scale “ percent transmittance range of the scope when photo was taken

Kg = the value of Kg determined from the given observation.

t , /_ = the approximate half time of the reaction to form iminè. 26o TABLE 60 An Illustration of the Sensitivity of to Change in Percent Transmittance

Amine Code ^ T i %'S. D Scale C cm ■'c V 2

cm ... % M - 1 sec

3~Methoxy- 3^ 22.69 24.94 3.08 5 2.35 1.5 propyl 22.69 24.39 2.34 5 1.76 1.5

3-Methoxy- 4^ 16.4 18.9 3.39 5 2.30 2.5 ethyl 16.4 18.6 3.10 5 2.10 2.5

2•2—Di— 5-â 22.69 23.94 0.86 10 0.897 10.0 methoxyethyl 22.69 23.69 1.38 5 0.760 10.0

2,2,2-Tri- 6-â 13*5 14.1 2.03 2 0.151 408. fluoroethyl 13.5 14.9 0.97 10 0.174 408.

Diamine 20 16.4 17.4 1.43 5 1.40 2.5 21 16.4 21.5 3.52 10 4.73 0.17

Diamine. 30 11.1 11.8 0.50 10 0.644 0.94 31 11.1 11.8 0.51 10 0.552 0.34

Diamine 40 17.5 19.2 1.49 10 1.82 0.90 41 17.5 19.1 1.13 10 1.00 0.90

— Conditions identical for both runs given. 26l

A qualitative description of the mixing phenomena of the stopped-flow spectrophotometer and kow it can cause an incorrect determination of the carbinolamine equilibrium constant is illus­ trated, Figure 69. Line a is the line for the flushing out of the cuvette and filling it with the two unreacted solutions and would be observable if the reaction did not occur. Line b is the transmittance vs time trace that would be observ­ able if there were not mixing (line a). Line c is the experimentally observable transmittance vs time trace. Point d is the transmittance at t o after establishment of K^. Point e is the incorrect transmittance at t^ determined from extrapolation of line c. 262

t t o mix Time

Fig. 69.— An exaggerated illustration of the mixing for the Stopped-Flow Spectrophotometer; a) mixing - the flushing-out of the cuvette and filling it with unreacted solution, required time ® time . ; b) transmittance vs time trace; c) observable transmittance vs time trace; d) transmittance at t ; e) transmittance at t^ determined from extrapolation of line c. 263 Program for determination of least square slope

for linear equation with two unknowns*

Symbolism

AIN = intercept of plot; Y = felope)X + intercept

N = number of points per plot

NJ — number of determinations

5 = N floated

SLOP = the slope of the plot

STD = the standard deviation

SÜM = term used to indicate a summation process

U = square of the deviation

X — a. variable

Y = a variable

Program line Very brief description

10-11 input

12-15 summing process

18-20 calculation of SLOP and AIN

22-26 determining standard deviation

27-28 output 26k

FORTRAN IV G LEVEL 1, MOD 4 MAIN

0001 DIMENSION X(IOO), Y(100),U(100) 0002 DO 50 NJ=l, 4 0003 N= 6 0004 IF(NJ .EQ. 4) N=4 0005 SUMX=0.0 0006 SUMY=0.0 0007 SUMXX=0.0 0008 SUMXY=0.0 0009 DO 20 1 = 1,N 0010 READ (5,3) X(n,Y(I) 0011 FORMAT (2*F10oQ) 0012 SUMX= SUMX + X(I) 0013 *SUMV=SUMV + Y(I) 0014 SUMXX = SUMXX +X(I)**2 "00*15 20 SUMXY=SUMXY +X(I)*Y(I) 0016 CONTINUE "0017 S= FLOAT fN) 0018 DENOM= S ’«'SUMXX - SUMX**2 0019 ■Ta n ^ ( " s u m ÿ '*s u m x X ^ sW x v s ù m x VT/ d*e n ^ 0020 33 SLOP= (S «SUMXY - SUMX’î'SUMY)/DENOM "0021 SUMU= 0.0 0022 DO 22 1=1,N 0023 U(I)=(Y{I) -(X(I)*SLOP + AIN))**2 0024 22 SUMU= SUMU + U(I) 0025 CONTINUE 0026 STD=SQRTtSUMU/(S'-l. ) ) "ooYf FORMAT1/3Ê20.6) 0028 WRITE(6,4) SLOP, AIN,STD 0029 50 CONTINUE* 0030 STOP 0031 END 265 Program for calculation, of an observed rate con­ stant from experimental data for hydration-dehydration of isobutyraldéhyde.

Symbolism: o A observed optical density at 2 850 A recorded at time

AE = observed optical density at t , after establishment of

AO = observed optical density at t ^ calculated from AE °

AIN = intercept of least squares function of f(A) vs t for the first-order expression, kt = f(A)

FA = the function, f ( A) in the first-order rate expression kt = f(A)

HE = AO - AE

KT = first-order rate constant calculated from k = f(A)/t

SLOP = the least squares slope of the function f(A) vs t; the observed rate constant

T = the time at which the observed optical density was recorded.

Program line Brief description

28-33 input

34-54 calculation of least squares slope, SLOP

62-63 output FORTRAN IV G LEVEL 1, MOD 4 MAIN DATE = 70023 1Ï/55/42

0001 DIMENSION B(10Oi,BH(l00) 0002 DIMENSION KHB(100),KDBI100),KHBH{100),KDBH1100) ‘' 0003 ■ ■...... DIMENSION RK( 1001,BTl lOOT,Ü 1FITOCTI,C%L(TOm 0004 DIMENSION A(6)tT(8)tFA(8)tKT(8)»SLO(I00)tOS(100)tU(8) 0005 DIMENSION KBI100),KBH|100) 0006 REAL KHB,KOB,KHBH,KOBH 0007 REAL KB.KBH 0008 REAL KT.KM.KO 0009 DO 400 M=“l#l5 0010 IFIM .EQ. 1) NN=8 0011 IFIM .EQ. 2) NN=5 0012 IFIM .EQ. 3) NN=4 0013 IFIM .EQ. 4) NN=10 0014 IFIM .EQ. 5) NN=2 0015 IFIM .EQ. 6) NN=5 0016 IFIM .EQ. 7) NN=9 0017 IFIM .EQ. 8) NN=5 0018 IFIM .EQ. 9) NN=5 0019 IFIM .EQ. 10) NN=5 . 0020 IFIM .EQ. 11) NN=5 0021 IFIM .EQ. 12) NN=4 0022 IFIM .EQ. 13) NN=5 0023 IFIM .EQ. 14) NN=5 0024 IFIM .EQ. 151 NN-5 0025 nKllciotXJ^I 002b 1330F0RMAT(IH0,4X,3HRUN|6X,7HKT AVE f7X,3HSTO,IIXtSHKT 1 ,5X,5HKT 2 , -fO 15X,5HKT 3 ,5X,5HKT 4 ,5X,5HKT 5 ,5X,5HKT 6 ,5X,5HKT 7 I 0027 DO 30 N=1,NN 8^ 0028 1 REA0(5t2l A OÎO o o OIO 0 |0 0 (0 o o OlO O'O OIO o l o OIO 0 (0 o o 0 : 0 o o c o o l o OIO o o o o 0 0 , 00:0 o: 0 0 / 0 0 0 o l o o l o 0 :0 O o O' O 0 ,0 0 o o O' O' O' O' VI W vn VI ui ui; ui V) VI V : 4> 4> -P- ^ •P'^U w L i t a U i w V i u > tjj w ; w N ^ o

t— w W N V 4 N o N N w o O' VI W tVi o ■n X C o C/5 < n ( n t/l O t/l t n t/l o X o o * n o z > • n 7 3 - n y a ■n o O 7 i r H O c — o C r - 1-4 m a C; c cic o c c c c c - 4 o o > o m o o m : o m o z T a — o o Z z m o z o z z z z Z Z!Z z 3 z z z ta z Il II > 7 3 J> 7> Z -4 » » Il - 4 C « - r v c T> Il o - 4 X X -< X N i X X - < X - 4 t a O' -4;ta VI z o z O > m z «— Il Il M Il il z >-4 • < X II Il O -c X Il II ta V ta > t a 3 > t a > z - 4 — ' i n Z t n O « ta VI m < n 7 3 m z * n il ô • 3 • m d c z Z II • • o o m -n m > 1 N i ta 4» C D <4 03 I-* »— r -4 c > * -< # z Z ■< X •- o o > ta • p* ta 0 N J t a -n üJ *n -n w o — t o t/l X X « 4 ■ o ^ > 0 ■ n t a t a t— N i • 0 * n + w • . c t/l C < X 4* W t a -4 o 0 m * t a > o o O — C ~ CD ■ Z C z > O m # - 4 # . z c X z X + + * n -4 4> z m • ' « o o wi c 1 •< X X > m o t a • c : ■ X i i m X H “4 t a t a t a c * - z - 4 1 . 1 — * «ta t a M» z •n # • - 4 1 ta» t a > »-* i n trt t/l 1 r C t/> c # # t a • O Z i C z *n * • n 0> T3 * x ! z X > N i > • «* t/3 * ! X * 1 »-«A r* t/3! » * #« t a n - 4 O c N ta > t- o • o Z:C m O* •< Z t a • + X Ul - 4 « -i > O t a Vfl *-< m X . X w z Z o <• w o m + - 4 * . 3 z ■n * o t a M'W * 3 O II r v • r— » V I "" ■ 4

t a t a

'

V

i , 9 S 268

Program for evaluation of the observed rate data for hydration-dehydration of isobutyraldéhyde, using the SIMQ subroutine.

Symbolism;

A = matrix coefficient for the 4 hy 4 matrix

B = matrix constants

D = concentration of hydroxide ion

E ~ concentration of N-methylmorpholine

F — concentration of the N-methylmorpholinium ion

K “ observed rate constant for hydration-dehydration

N = size of matrix, 4 equations and 4 unknowns

PH = pH

POH = pOH

VI = coefficient for k._ AO V2 = concentration of hydroxide ion

V 3 = E

V4 = F

VS = K Program line Brief description 6-12 input 31-36 calculation of solution conditions 47-81 construction of least squares matrix 84 subroutine call 85 solutions of matrix, output 87-92 output FORTRAN IV G LEVEL 1$ MOO 4 MAIN DATE = 70023 11/55/25

'0001.. DIMENSION At4,4l»8l4l 0002 DIMENSION DI15),E(15)fF(15)tKC15) •PH(15),P0HC15) - UÜÜ3 REAL K 0004 REAL K1I20I 0005 N= 15 0006 READ 15,3» (PHin, 1 = 1,N) 0007 READ (5,3) (Eli), 1=1,N» 0008 READ (5,41 (Fill, 1=1,N) 0009 READ (5,5) (K(I»ff 1=1,M) 0010 3 FORMAT (8F10.0) ■ 0011 4 FORMAT (8F1O.0I 0012 5 FORMAT I8F10.0I • 0013 SUMl = 0.0 0014 SUM2 = 0.0 GDIS" ■ SUMS = 0.0 0016 SUM4 = 0.0 0017 SUMll= 0.0 0018 SUM12= 0.0 0019 SUM13= 0.0 0020 SUM14= 0.0 0021 SUM15= 0.0 0022 SUM22= 0.0 0023 SUM23= 0.0 0024 SUM24= 0.0 0025 SUH25= 0.0 . 0026 SUM33= 0.0 0027 SUM 34s 0.0 0028 SUM35: 0.0 o o o o o o o' o o' o olo c> o o o ® ® 2 ° ° o olo oio o |o o o ci O; o 0 ,0 o; o 0| o o o Ol o o o o O 0 ;0 0 ;0 OjO o o o O OiO 0 :0 0 0 0 0 0 : 9 ‘ 0» o o üi ui vr. uj u t ; ut uii VIwi VIVI vnVI ‘ ^ i •»: f- ^ W w Mt'.Mi w! w w, w w w IV; M i r o ^o *0«o 00 -j O' UT; • fV- ■ T— O -iO — ®® ^ -1-'IO'U1-P'Wf\3H“ ~ ^ “ 0 ' 0 C9 >VO'Ul<^WIVr«0'0

o o

OD CD o c/1 o n 00 l/l c/1 UT C/T O o U T U l < < < < < < < < < m n 7 : O "V o C(0 l / l — O C C C d e UT w N) H» U l ^ W IV o o c c IS> ► - 2 Z 3 3 3 il Il II Il II Il II Il II II 2 2 2 WWW î î i I UT ^ UT >{ w VJ M I—' < < < < < 7t TT m Il II « * Il M O Ul i«' H II il Ul ^ W rv H — o Il II II II Il II Il II II V. 's. o O T- II > (/) o> m i/j ' < < < < < < *- o 7 i O o O . C C < c < w Ul Ul Ul UT UT — * a — * I - Il • • . z 3 w * S 2 S 03 O w o o • IV * * ' * < W — * • f- I VI Ul < ^ 5 5 S * u: * Ul i^lv Kj m I ■**

0 * * •o1 %

ois 0 0-0 oj o o_ o o o o|oo o CÎO 0!0 o O 0 ; 0 O O O; O 0 , 0 0 o o o o; o o: o o! o o oo! o o; o o; o o O 0 ; 0 o; o O; o o; o o ; o o o s O «fi' vC 00 00 œ 00 00 OD oo; 00 00; OO -4 N •V "<4: "V mil O'! O Oi O* O'' O' r v 1“ O «fi CD n O' Vi O' VI 4* W IV P- o «O 00 O' VI: f

•V V o o o N ■ V p i w o ro m o z o -n *n X o X ■n > > î> ï» > J> > > > > > ï> > O; CB o 70 o a o 30 o %) O;?0 §i Z p- 30 30 PI z PI 70 w N w N I- ^ V I N I p i ^ w IV p- W *0 •H -H -V 2 z - 4 -4 H Z PI m •V > > m n m > m ^ W w V) W tv N rv rv z — - 4 H — 1 p » Z — VI H < - Il II c O' •» Os c O' Il II II 11 m <• (_ P- pi 1 m * Il II II Il II Il u Il II Il (A N Z X % -4 IV (/» 00 V) (/» C!C O Z O pi O IV C c c (/I l vo w co c 1/1 (/)(/)(/)(/) (/I ■f* ^ .f' VJ W l\l f ^ rv IV pi p— p- p- C. 1 «# 00 w w w w IV IV w IV z Z VI z •n pi 75 VI • w c_ VI II z z p i « II I— 7« * pi ■ e « L. Il

na ■ V I «* z % 00 z p— «# m M II z! I-» Z «

c_ g - 2

Tl3 272 Program for calculation of equilibrium constants^

^Caobs ^laobs All symbolism is explained: F. A. Via> A Kinetic

Investigation of the Reaction of Methyl amine and Iso­ butyraldéhyde, M.Sc. thesis. The Ohio State University,

1968.

Program line Very brief description

9-10 input

11,27f 36 elimination of correction for protonated amine

16 correcting amine concentration

18-25 calculation of initial and final 32-34 concentration of aldehyde

30-31 expression for calculation of K_ la 38-39 expression for calculation of K

40-41 changing estimated value to calcu­ lated value before recalculation

47-48 output

50-53 method of preventing undesirable de- 55—60 terminations from being averaged

66-86 calculating the average and standard deviation

87— 88 output

95-96 format for output 1 ■ *n c o c O C O cl O Ü o o o o o o o d o C C) o c o c o c> C' o o o O O O Cj o o O C O O o c o c o o o o o o c c o 73 l\JN> IS.IV IV N N IV IV ^ h- ►“ 1— h- f— 1- r-| o o c o c o c c c —1 05 -g O' V* W IV ►- 0 O 03 -4 (7 Ul 4' W IV nC 03 O' U 4' U IV u- 73 > Z . M» ' < c r m 1 IV IV < c o m u> l\! f- (—• r K- o c a c m > c r> t> - — 1 — c nr T1 C C/1 C3 73 T s: l-m z c o o 1— c Q a «“> IT II n —I % 73 XJ ■H z 3 IT G II 7^- C m Q PV O G n '-i II II Z II II n H rn II II 73 %> Z II >:i > u 73 — II O -< z 2 S II P o II II II m Z C 4' % C Ü G -H "'-1 rr m rr. 3 o CP o r > c • II o o o Z rr > O > X > rr o u II z z z d h- - 4 H •H >s IV • IS Z —1 “— 3 • —IV. c ■H u !y- C/1 c/3 o rr H- Q IV C -4 n '*> V II o — — O' o — z 1 1 > -«4 • h- 73 73 IV •V * V IV 1—' <• (_ o Q O 4> *- c: h- o -V cr s* V • Ti U H- — Z IV II z z I > 4> »- 1— V O' O' c »-• ~ <• ■< o o II m t» # 41' O' • • o z . o s> •» t— 1— O 3 > • Ul • a 4> — m — O -H 1— ÜJ O. m3 c JS o •< X c < I— ►— o C C — —i n w 4' 1— Ol . o o + > IV & oi h-» <* f” 1— « * h— •- c m X O |\J o * Cl > r* IV o — V o z Q —4 •s o P> m \0 - C/v — > V CD X a c/v m. » C ■» m -H if -1 + (/) IV < r- (T c G IV O Z — c ÿC 3 X — o •>. « m o 1— — 3 m c IS <• — o - > A* c z 00 O -t c 7- X — G h- . 03 ■tt u — c c I IV :< — m c 3 — p rr -4 > 3 — < 1 Ul m X X 73 c c • c • z t—' fv > h- O' —1 C X ■ ## nr iTi Ln o O It ■ V c c ro -4 —1 r- W O MM o 03 < IV 73 IV H4 Mm c z z ro D 4# -1 - < > ■ z •9 4> w X V Ul X IV V 4' w V. h- «0 Q z 0029 B=(UO-AT- 1.0/Y4 + SW4T((B0-4T-1.0/YA)*<2 + 4.0*Ü0/YA))/2.0 JOJÛ.. ELEFr= AI/(Al»EA '+'bO*ÜB - 0) ' 0031 VI= l.Ü/(b*(ELEFT*SUMü - 1.0)1 0ÜV2""" AC= C O . 2I V1 1 .0/ 1 XC ) ) / E A 0033 • AC*EA 0Ô34 CBA= AT -AC • 0035 i.C0=llT-l3CH 0034.. " HCU= 8r 0037 15C=(BCO-AT-1,0/YCA + SQRT((0CO-AT-1.0/YCA)**2 + 4.0*BC0/YCA))/2.0 00 3b ULLFl= AI/(A)*LA + BCU*bd - C) 0039 YC= 1.0/(BC*(CLEFT*SUNE - 1.0)1 U04(T YA= YI 0041 YCA=YC 0042 401 CONTINUE 0043 81= 1.0/8 0044 BC1= 1.Û/BC 0045 X (M ,1)= YI 0046 x (M,2)= VC 0047 WRITE (6,10) YI,YC,87,80,81,ELEFr,BCl,CLEFT,RUN 0043 10 FORMAT ( 1F10.2,1 F 10.3,6F10.4,IF 10.3) 0049 4 CONTINUE ÜÛ50 DO 101 JJ=1,N 0051 IF(J .EO. 1 .AND. JJ .EQ. 5) GO TO 110 0052 IF(J .EQ. 2 .AND. JJ .EO. 5) GO TO I10 0053 101 CONTINUE , 0054 GO TU 44 0055 110 X(JJ,1)=0.0 0056 X|JJ,2)= 0.0 0057 WRITE(6,111) JJ 00)8 111 FCRMATdriO, I3»22HTH ENTRY NOT TABULATED) ' ' ' 0059 IC= IC + I 0060 GO TO 101 0061 44 CONTINUE 0002 Nl= N - IC 0063 Sl= Ni UÜb4 SUM1= 0.0 .. • ÛObl> • SUM2= 0.0 ■ CObb UC) 301 1= 1 ,M U0ü7 5UK1= SUMI + X(I,1) 0063 5 UN 2= SUriZ + XI 1,2» 00o9 301 CONTINUE 0070 rniALt1)=SUMl 0071 TOT AL (21= SUI-12 "OÙ 72" AVEK(I)= SUMl/SI 0073 AVËR12)= SUM2/SI 00 74 SUMU1= 0.0 0075 SUM02= 0.0 0070 OU 302 1= 1,M 0077 1F(X(I,1) .EQ. 0.0) GO TO 302 0078 DEV1(II= (X(I,l) - AVEatl))**2 0079 l)EV2(I)= (X(I,2) - AVER(2)»**2 0060 5UMül= SUMDl + OEVlli» 0081 SUM02= SUMD2 + DEV2II» 0032 302 CONTINUE 0083 SD(1)= SQRT(SUiMÜl/(Sl - 1. ) ) 0084 SOI 21= S0RT(SUiMU2/lSI - 1.)) 0035 00 00 1= 1,2 0086 rui (i )= iülAL

3 Ü096 TF(NCYI VKOr Jl oO rO T55 0097 7 7 CONTINUE 0Ü9Ü 5b COiTITTTOt 0099 00 TO 99 OlOO" viK 1 TE t 6 f 1561 0101 156 FORMAT!IH11 0TÛ2' ■ NCYT= 0 0103 GÜ TO 77 0104 ... .■... 09"xUiSi'irrTUb ■ ■ ■■ " 0105 STOP

a 277 Program for calculation of from stopped— flow photo.

All symbolism iis explained: F. A. Via, A Kinetic

Investigation of the Reaction of Methyl amine and Iso­ butyraldéhyde, M.Sc. thesis. The Ohio State University,

1968.

Program line Very brief description

7-8, 16-22 input

23 concentration of amine in cuvette is one-half that of the stock solution

24-44 calculation of initial and final aldehyde concentration

48-65 calculation of kinetic expression

71-90 least square calculation for each run

91-92 output

94-105 averaging runs of same conditions

106—114 output and format for output FORTRAN IV G LEVEL I, MOO A MAIN DATE = 70027 19/38/22

C FORTRAN 4 PROGRAM FOR DETERMINING X2 FROM STOP-FLOW PHOTO 0001 DIMENSION DMtO),TR(8),T10),X2{7),WX2(8),A(8).0D(B),0X2(7),018) 0002 blMENSiON TA(8),DMA(8),DS(100J,SLO(100),AA(100)»CP(100),ZI(100) 0003 DIMENSION ZZE(IOO),771(100) _ ’ 0004 DATA EEIH/0.01/ 0005 . _.NFPO= 0 ...... 0006 DO 400 11=1,13 0007 _____ RE AD ( 5 , 500 ) Y I , YC , NN , DDMER______0003 500 FORMAT(2F10.0,I3,F10.0) 0009 YI= YI - YC 0010 8= 1.4286 0011 WRITE (6,133) YI,YC ...... 0012 133 FORMAT (IH0,26HTHE FOLLOWING RUNS USE YI ,1F6.2,7HAND YC ,1F6.3) 0013 WRI TE (6, 131)______0014 131 FORMAT(1H0,7X,3HRUN,lix,4HSC6p,l6x,3HSfD) 0015 DO 30 N= 1, NN ...... 0016 1 READ (5,2) DMA 0017 2 FORMAT (8F10.0) _ . ______...... 0018 READ (5,3) TA 0019 FORMAT(8F10.0) ______0020 READ (5,4) DMTR,DM0A,DMER,70RB,DT,EH,RUN 0021 IF (DDMER .GT. .0.0) DMER= DDMER...... 0022 FORMAT(7F10.0) 0023 EM= EM/2.0 : 0024 DMO= O.0313«OMOA - 0.108+ DMOA + ( 0.157*70RB + 0.002)*DT 0025 DME= ( 0.157«DMER + 0.002)«DT ______■ ...... 0026 DMT= ( 0.157»DMTR + 0.002)*DT 0027 100 TRO=DMO/DT 0020 TRE=DME/DT 0029 TRT=DHT/DT 279

0 m4 o m *

o ' CM OÎ K: O O 0 3 0

#"4 S i 0 CO 0 QC 0 o •■ "si 0 # N # in K 3 ' i •

A U i N N 0 CO + ' < 1 * 0 CO , I— CM * CO! ■ k- <' i J - » CM = + 0 N» IJU _J —» N | : CO CD I Ui <1 _ i ; O - — LU "^1 + < . ■ NI in • * LU •; fO H M < H" «—I —4 * * • • •«• N O U i , 1 co 0 # II < m% f*. (S. o O O "d" NI ^ ~ 3 > . h ~ < O f + Î » 0 • NJ 3 U < m# —I r-i r\i f\i r\i «-t > »-« — < 0 m4 • 0 E #. 3 > #"4 *— •H 0 fM -N) (NJ *s. -s. II + m LU 4- 0 0 0 CD II 0 0 0 — m • » • o m 00 :Ni ( «—1 II II II II II II + 0 0 : + o o o o a Q o - 3 -3 CD O N I— I 0 0 * . ! < II m> m* a 0 3 — I— It II II O O Cj a < a i n *—• *—• ml m l 3 3 m l «S ml 1— >- E U ~i — O LU «- II II II I! LU %lf* ; II II UJ II < > ' II Q C O C UJ I- ^ II *U u _ j II 3 II II Z 0 II < CL

• • : ' : i . I I : • . 0 '-

I I I I ■ ! : ! o o o 0:0 ocooooooooocoooooo ooooo oo ooooooooooooooooooooooo ooooo oo 'O'C'O'OOPCüOa O3CO0C 0O 0:0CC3N«<— -u o o o 0» c-o WN»—0 f'WNi-»O'CC0 NOV'^wr\*»— O'ÛCD-'JO' u*. J' I I

10.' : w w j PO Iw PO m o O PO PJ :w o i t» ' Ul ^ i 1 b- o n •n y t/l'l/l O l/l C t/l t/l > O O t/l t/l t/l l/l w t/l l / l C/^ t/l n X C C 0 — . i: I» o o » r* :-i Q c (3 c i— m o c c C C O C c c c o PO X o c t» X — Z p o o 3 w O 3 0 z z z 3 3 3 3 3 3 PO 2 r* PO •— H 3 H « « Il - 4 C « c ,x II c - 4 XX - < X PO X X -< X —t * > m 3,1/1 II II PO II II 3 -< X II U o •< X Il II w t—t M c Il Z H *-.o Z t/l PO o t/l II z II II t/l t/l II II o o Z * — C o Il 3 -' c c »— • C3CCSCt/it/iCC O o • • a II il c > Il > m W « i/i.-i m 3 z o • 3 • m c c 3 Z II # ■ o o l/l m X z I— m —« > X r-,— C X II * < 3 3 ■< X b- o o PO X - W o t/i PO b - t/l ■» t/l X X « X — Po —01 — PO r* 1- PO X c + —• — c t/l c -< X . + + 0 3 73 i • •n w |3 w 00 c 3 a ■ + # + *— : le C X 3 X + + a: "4 o j ■ < • 73 . < X X :x 7 3 ' ! o I— w c l - J 1 1 X -H -4 PO <-• 1 > — Z ;• , 1 1 » »— w >• —» 1 bX ' r I X v> t/l *n ■ + . + c O' I- I I > M o 2 s 70 ■ <— m •n ■o X X X PJ o o * i* * PO n . S ^ V» t/1 t/l * > ■ PO a o • H r- C t/) N VI o o C R: -ÎJ- « -o 3 C • > 7» X r- r- + i a a j O ,m X, o a :ü C t» ~ O m i; 1 o J> :3 3 * if -* as !: 5 i : * PO m; — ip t/l; ^ a C' + Tl; o Z": ;

îo ; : Ci': t-4 : < m

< ; > ; r- ' c . m V» O "n,

083 C THE FOLLOWING STATEMENTS CALCULATE AVERAGE YC, VI, ANO SLOP 0094 _ SUMSL= 0.0 ______0095 ÜO 201" N = 1 ,NN ' - 0096 SUMSL= SUMSL + SL0(N)**2 0097 201 CONTINUE 0090 WN= NN 0099 AVeSL= SORT(SUMSL/WN) 0100 SUMDS= 0.0 _ Olül DO 202" N=1,NN 0102 D 5 ( N » = (SLQIN) - AVESL)**2 0103 SUMDS= SUMDS + DS(N) 0104 202 CONTINUE 0105 STDSL= SQRTISUMDS/IWN - 1 .) ) 0106 WRITE(6,212)AVESL,STÜSL 0107 212 FORMAT I1HÜ,9X,3HX2 fF14.5,2X,1H$,2X,FI4.5) 0103 NFPO= NFPO + 1 0109 IFINFPO .GT. 3) GO TO 401 0110 GO ro 400 0111 401 CONTINUE 0112 WRITEI6,402» 0113 402 FORMAT(IHIÏ 0114 NFPO= 0 0115 400 CONTINUE 0116 STOP 0117 END

% H 282 The program for evaluation, of the pH-rate data of the monofunctional amines

This program as shown on the following pages is designed to evaluate the equation

^ o b s = ’'o +

By changing line #45 to read V3=0.0 this program can construct a 2x2 matrix for the solution of the equation

However the SIMQ subroutine will not properly solve a

3x3 matrix with one coefficient equal to zero. Thus, the 2x2 matrix was not solved by the program. The average concentration of the general acid (BH) During the reaction is assumed to be that at t^y^. The average reaction was observed for two half-lives and this assump"- tion appears reasonable.

Symbolism:

A = matrix coefficient ( 3x3)

AT = total initial aldehyde concentration

AE = aldehyde concentration after establishment of Kj

AHAF = aldehyde concentration at t^y^

B = matrix constants

DMER = measurement used to determine TRE 283 DMTR = measurement used to determine TRT

EMHAF = [e MT] at t^y^

EMT = total amine concentration

G AC ID = average concentration of general acid

H = hydronium ion concentration

K 2 0 B S =

ODE optical density after establishment of K_, 2850 Î

ODT = initial optical density before establish­ ment of 2850 S, i.e., the optical density before reaction

PH = pH

PKA = pKa of the amine

SD = individual deviations

STD = standard deviation

TRE = observed transmittance after establishment of Kj at = 2850 X o TRT — observed transmittance at 2 850 A before reaction

VI = coefficient for ko V 2 = H

V 3 = GACID

- k j o b s XI = mole fraction protonated amine 284 Program line Brief description

25-29 input

30 adjust concentration to stopped-flow cell

31-44 calculation of [BH] at t^^y^

47-74 constructing the matrix

76 call the subroutine to solve matrix

77 output of matrix solution

80-107 input and calculation of a standard deviation

108 output

IIO-II3 this entry is used to convert the 3x3 matrix to a 2x2 matrix 285

0001 DIMtNSIÜiV A (3,3),8(3) 0002 REAL K208S 0003 REAL KCU8S 0004 REAL IVP,IHAF 0003 rJ= 3 0006 DO 300 NJ= 1,3 0007 SUMl = 0.0 0008 SUM2 = 0.0 0009 SUN3 •= 0.0 • 0010 SUM4 = 0 . 0 0011 SUM11= 0.0 0012 SUM12= 0.0 0013 SUM13= 0.0 0014 SUM 14= 0.0 0013 SUM 15= 0.0 0016 SUM22= 0.0 0017 SUM23= 0.0 0018 SUM24= 0.0 0019 SUM25= 0.0 0020 SUM33= 0.0 0021 SUM34= 0.0 0022 SUM35= 0.0 0023 . SUH44= 0.0 0024 SUK43= 0.0 0025 REAL)(5,2) PKA,NN 0026 2 EÜRMAÏ(FIO.O,13) 002/ 0Ü ICO JJ= i,NN 0028 READ (3, 3) K2ÜBS, EM r,l>U, DMTR, DMER 0029 3 rURMAT(3F10.0) 0030 EMT= EMT/2.0 0031 IF( IMJ . LU. 3) EMT= 2.»EMT 00 32 TRE= DMER/6.33 0033 T R 1 = IJnTK/6.33 00,4 0DE=0.2171*ALUG(l./TRE) 0035 UUr=0.2171*AL(10( l./TRT) 0036 /.T = 00r/14.6 003/ A5= ÜDÉ/14.6 0038 H= 10.0*»(-PH) 0039 V 1= 1.ÜOÛOQ 0040 V2= 10.0»*(-PH) 0041 AHAF= (At - AD/2. + AE 0042 EMHAF= EMT - (AT - AHAF) . 0043 XL= !./(!. + 10.0**(-PKA)/H) 0044 UACIU= X1*EMHAF 0045 Vj= ÜAC1D 0046 V4= K2ÜDS Vl= V1/V4 U048 V2= V2/V4 286

0049 V3= V3/V4 ■■■ ■■ , OObO V4= V4/V4 UU3i isUi'Ur = Vl*»2 + SOMIi ' ' ...... OObZ SUMI2 = V1*V2 + SUM12 üùi>3 5ÜM13 = V1*V3 + SUM13 "" ~ 0034 ; SUM 14 = V1*V4 + SUM14 UUbb 6UM22 = V2**2 + 60M22'. ' ■ ..... 0036 SUM23= V2*V3 + SUM23 ÜU3 / SUM24= V2*V4 + SUM24 ...... "" 0058 SUM33 = V3**2 + SUM33 00^9 SUM34= V3*V4 + SUM34 ... " “ " ...... 0060 100 CONTINUE 0061 on 503 jbi)= 1,2 ■■ ...... 00*62 A (1,1) = SUMIl 0063 A(Z,1) = SUM18 ■■ ...... ■ 0064 A(3,l) = SUM13 0063 A I 1,2) = 8UML2 ■ ...... 0066 A(2,2) = SUM22 OOo { A13,2) = 6UM23 0068 A (1,3) = SUMl3 UÜ69 A I 2,3) = SUr.23 ...... 0070 A (3,3) = SUM33 00 a 8(1)= SUM14 ...... 0072 8(2)= SUM24 Ü0 n 8(3)= SUM34 0074 WRIT£{6,22)■ ( (A(I,J),I = l,N),J = l,N ),(B(I) 00 73 22 hUHMAI l/3t:l6.3) ' ...... 0076 CALL SI.MQ(A,8,N,0) 00 7 7 w:

1 _ _

CD * m > O o + # o r-4 o ‘ • > I/Î V. '0 3 Q t IJJ .-C 3 )— N M • ••• *-4 co > > > > • V c CO LU OC V. o 3 O O LU LU O CD 'V »s *S. V* r4 o c —- 3 3 < N -t OJ rr\ ^ 3 > II Z to 1— II II II II Z Z O D£ > > > > CO 3 w •XI < •o CO •0 1—1 —) CO t - II L— ro ■n CM r - t L- 1— X II II II II II II — II nc z 3 M oc % s : Z Z O 3 m -f r-l N m -j- u . o 3 C or o 3 3 3 5 3 1 - > -•> > > > > •—« / ) CD ^0 K X CO CO cO cO t_) 3 to IX) fM ro m O O CM o Q rH n CO

r- I» O ' O r-* fO •4' JD O h- CO O' O -1 (M ro <- in vo O ' O ' O ' O 3 CD o o 3 O O o O — i rH r ~ t r—1 —4 #-H o o 3 ,-1 U l r-t p"H fH *~4 1—4 r-l r-4 f - i rH rH p—4 r—4 CD O 3 OiO O o o o o o o o o 3 O 3 o 3 O 288

Program for evaluation of ^Cobs PH data, for diamines.

Symbolism:

EM = EMI = total amine concentration, all forms

EMUP = concentration of unprotonated amine

EMlP = concentration of monoprotonated amine

EM2P = concentration of diprotonated amine

H = hydrogen ion concentration

KCOB = K calculated from the average values ^ KU(2) and Kl(2) and the particular solution conditions

"(COB = KIOB = K _ , calculated from the average values KU(1 ) and Kl(l) and the particular solution conditions

KIOBS = K_ . lobs KOBS = used to store values of K_ _ and K_ _ _ . ., . _ Ccal leal for plotting purposes

KTHEO = used to temporarily store values of and ^ to compare them to and ^ respectively, for determining the standard deviation

KTJ(l) = average value cf the equilibrium constants for carbinolamine formation with unprotonated diamines

Kl(l) = average value of the equilibrium constants for imine formation with monoprotonated diamines

Kl(2) = Average value of the equilibrium constants for carbinolamine formation with monoprotonated diamines 289

PH = pH

PKAl = pKa^

PKA2 = p K a j ^

SUM — summation function

XX and W = real variable used to store pH values

XU = [ EMUP ]/[EMI ]

XI = [e m i p ]/[e m t ]

X 2 = [EM2P]/1eMT]

Program line Very brief description 15-16, 90-91 input

19-27 calculation of solution composition

28-29 output

30-47 least squares determination of average equilibrium constants

50-58 determination standard deviation

S3-54> 59-60 output

6I-68 assignment of plotting coordinates

69-71 plotting subroutine for observed points

74-84 calculation and plotting subroutene of average curve, for a pH vs K plot

92—96 calculation of experimental deviation for equilibrium constants TT/a'N DAfF. = 70022 21/45/00 QOÙT i3n^E^3TD:Tnrrrmi~,ik7rmT

0028 WRITE 16#777) PH.XUlI),X11 I)#X2I I) o Ù0Z9 777 FORMAT* 4F10.6)... ' 0030 50 CONTINUE 0031 ■ ■ SÜÎ>ri= 0.0 0032 S U M 2 “ 0 » ü 003 3 bUM3= 0.0 0034 SUM4= 0.0 003 5 6UI:5= 0.0 0036 DO 102 1= 1»N ■ 0037 ■■■...... 1F xi<= i3 r r 0066 YU= 0.0 OOo/ IFfj .LU. l) YL= 120. 0068 IFtJ .ÉQ. 2) YL= 2.20 UÜ69 CALL PLOrA{K,XLiXr<,YU,YL,U 0070 CALL PLÜTO(XX,KIÜBS,4H0G00,N) 00 71 CALL PLÜFG (AAA,NAA.,AAt),NAB, AAC,NAC) 0072 CONTINUE 0 (T73 l'H= 3.0 0074 00 110 Jj=l,lOO UO'/S" Pu 0076 H= 10.0»x=(-PH) ü07 r T73UP= ÉMf/(1.0 + H/10.Ü*»(-PKA1 1 + H**2/lO.0**(-PKAl - PKA2)T 0073 EMIP= (EMUP*H)/10.0**(-PKA1 } 0079 K 0 B'St JJ)= (E,\lP/EHT)*Ki( J) + (EMUP/EMT) 101 CUNTTmjk 0066 WRITE(6,4) UU87 0 = 8 0 0 8 8 Nl= M ITObO' : Q,j 301 1 = 1,Ni ' : 0090 KEAD(3t302) EMT,PH "ÜITÿT' '302 FDRMAr {2F10.O) ------0092 H= 10.0**(-PH) 0 0 9 3 b,NUP= tMf/d.Û + 0/10.0»»(-PKAl ) + H**2/lü.0C*(-PKAl - PKA21 ) 0 0 9 4 EMIP= (EMUP*H)/1D.0**(-PKA1) 0Ù95 K[D fî=TE MUP/1M T)*KU 11) + (ÉMiP/èMrRiaTiT 0 0 9 6 KC0B=(EMU?/EMT)«KÜ(21 + (EM1P/EMT)*KI(2) OOÿ 1------l Tf {'6 ."BCT3 r~K roB7KCûBI PH itM f “ OOia 303 FORMAT(m0,4F10.6) T-09-;------îDTnjTïTmimE- : ~ ülOO WKITE{6,303) KU(1),KU(2ï,Kl(1),Kl(2 » "OTüT S T û T 0102 EîMO 294 The program for evaluation of the pH-rate data for imine formation with -dimethyl- amino-alkyl amine s

Two programs with identical symbolism are employed.

The first uses a 4x4 matrix to solve the equation,

tC-*««^2obs“‘'uo tC ]+ k ^ [BH ] [C ]+kg^[c«+]+kp^[CH+] [H+]

The second uses a 3x3 matrix to solve the equation,

[c-K:H+]k2^^^=k^^[c]+kg^[CH+]+kp^[CH+][H+]

The composition of these programs is illustrated on the following pages.

The concentration of the general acid [BH] is assumed to be equal to the initial concentration and mono- and diprotonated diamine. N,N-Dimethylethylenediamine is the only diamine in which general acid catalysis appears to be significant. In the pH range of general acid catalysis, 10.4-8.4) the amine composition varies 90-9^

LeMUP], 10-90^ [EMlP], and 0-1% [EM2P], respectively and since monoprotonated amine can form monoprotonated carbinol— amine and monoprotonated imine all of which can behave as acids the assumption that [BH] = [EMIP] + [eM2P] is a reasonable approximation. Although EM2P is expected to be a better acid catalyst than EMlP the concentratic# of EM2P is always very small relative to EMlP and this catalytic difference is not considered significant. The estimated 295 pKa values of the diamine, carbinolamine, and imine differ pKa S» 0.5 but a two-fold excess of diamine to aldehyde will minimize the effect caused by the different acidities of these acids.

Symbolism:

A = matrix coefficients

B = matrix constants

EMT = total amine concentration

G AC ID = concentration of general acid

H = hydronium ion concentration

K20BS =

N = matrix dimension, N=4 4x4 matrix, N—3 3x3 matrix

NN = number of data entries

PH = pH

PKACl = pKa^ for carbinolamine

PKAC2 = pKa_ for carbinolamine 2 PKAMl = pKa^^ for diamine

PKAM2 — pKa. for diamine I t SD = square of the individual deviations

STD = standard deviation XU -* mole fraction unprotonated carbinol- 296 amine

XI = mole fraction monoprotonated carbinol­ amine

X2 = mole fraction diprotonated carbinol­ amine

XMU = mole fraction [eMUP]

XMl = mole fraction [eMIP]

XM2 = mole fraction [eM2P] 297 Program line Brief descrimlbiom

3-30 input

37-40 calculation of [SH] concentration

41-85 construction of matrix

88 subroutine to solve matrix

9I-II2 input and calculation of standard deviation

31^ 86, 116 output I G> Ci o Oj o c O O o o o o o o o Ql o o cl Q o o o c s s O Q O C O O C o o c o Oj 0 0 0 0, 0 0 0 0, 0 0 1 0 0 o o N NJ ry fV N; t\i —h - I— I— »— r — f—I ^ K—I r— o o c o c o o o o O» ^ e* vTt -f» w i\» c 00 -J O ' 0.5 .i' oc- l\J f-; o o CO S' OTl -> ro 1 75 >■ 2

<

m c m

»';e X'-C TJ "C2 (T. 05 05 0 5 0 5 0 5 ( / ) 0 5 0 5 : 0 5 0 5 0 5 0/5 05 0 5 05 2 m a y: 7: 7 ; 2 C C cz c c C cri c: Cj cr c jc c C c c II rr ►—I > > II jC 2 2 ? 2 2 Z| -ï: ■< — > 2 a- O'#" n o ^ U rc !v r\J 5 J: r-* w rv — m wm @ ÎVJ 1- N.' M -P> c 00 l\)' U" z \ r . o ti II Il III O' Il II II Il II ; Il lii II II Il I II II II 7: 05; IV •—'j w 6. 0 5 h-" I-, c o c Oi o o: o o: o o o|o o o o o: 4»• o: . Oj • I « « ! • e o 2' II • IU1 • o o o;• !• c Cl o 0 0 0 o5 m o o! - u- 2 1- ■■ %C5 r- ■0 2 i 7: i\j S 3: tyt >

C >

o o en V,

I

o j ! -) SUM. 24= V2*V4 + SU M2 4 UÜ39 SUM25 = V2+V5 + SÜM25 0060 SUM33 = V3**2 + SUM33 0061 Su m 34= V3*V4 + SUM34 0062 SUM35 = V3*V5 + SUM35 0063 SUM44= SUM44 + V4»*2 0064 SUM45 = V4*V5 + SUM45 o o c O o oi o c c o o ci o o o ac} Oo ci O Q O O i O ci O ci o c] o O oc. o o O o o o o o o o c; o G O O O 0 o o o Oj o o q o c o 0 O Cf. c < i «s c o OI'

IVJ o rs> o XX o 5" > > > > > > > > 3> > > > > > G (T O M m > c —^ — O il II i?" r- X w (V w (V h- J> uj rv r— J s O J (VI; M ^ fvJ f- Z ■ M I— X XX «M c > ■s- j> w o w w fvji (\j (V; t— M C c/1 —( II Il II Il z « •» p o Il III II II c 3 Z ! Il Il II II Il II Il II Il II < / l ooi oo m >*> — » * M 1/1 c/5 (/I M — c z ci c c c 3: cr c c c c c c IVJ I-- • • Z. M :< u- .-V! r- 3; 3: Zj Z fvj z z ir ^ ■î'. on u1 on 0 o :?S — j> f" bJ bJ (VJ h- .f" rvj ivj « # U) KjJOJ Ijj fv»fv .f w IVJ « % 1 I ■c % X 35 X NJ 3> » c d n.c •i;- C3 MM oo I •o £> O II M* 2 X » * M G V.

o * il * M T z ■o X > o

I ■o X c I

OOÇ • OOV f X;'U = l./( 1.0 f H/10.0«*(-PKAM1 ) + Hi:*2/10.0*«(-PKAMl - PKAM2 )) C0-;:3 XfU= XN'U*H/10.0**(-PKAM1I 009 V XiV2= X M l /10 . 0* :>( - P K AP 2 ) OlOC ÜACID= EMT*XK1 + EM7*XM2 ülOi ■■ Vl= XU 0102 V2= XU^-GACID Ü103----- V3= XI C104 V4= XISH ülOb V5= K20BS Cl 06 Vl= VI/V5 (jro/"" V2= V2/V5 cioa V3= V3/V5 0 1 0 9 ..... V4= V4/V5 01 10 V5= V5/Vb Û IL l S0= (V1

VJJ P 302 Program line Brief description

4-29 input

34-67 construction of matrix

70 call of subroutine to solve matrix

73-91 input and calculation of standard deviation

32,68,92 output 305

m 3: m

N «f so r cv _ >- 5. > - :L Zi s; Il < < < < — ül U l M Z3 ^ O 3 3 3 o 3 3 3 3 3 3 3 Z? V •ü£ O \A

r\i .-o \T CÜ c '-I f\j m ,\j- y\ o f cO'ô' O r-'-iM p >r ir. \c r*- p 3 O O O l3 O lO O o iM r v fM ;.NI CM irsi (M !\ J C M O O o 0 J0 -'0 “'0 ^ O o p O !0 o o o iO O o O io o o o :o o o o OsO P 'O lor o io o o o p o p o p o io o o o p o o o jO o o o 002') ReA0(5,3) ÉH1 ,Pll,K2UBS 0030 EMT= (:MT/2. CÜ31 3 FÜUf'ATOFlO.Oi 0032 w m T E ( 6 , 4 ) C.U ,PH,K2Ü8S 0033 4 FOKf-iAT { lH0,:if-l6.4) 0034 i\= 10.C-=!=(-Pll) 0033 X'U= i.0/(i.0+ H/i:j.0^.-*(-PKACl) + H»*2/10.0**(-PKACl - PKAC2 )) 0036 Xl= XU»H/10.0»*(-PKAC1) 003r X2= XliH/lO.O*»(-PKAC2) 0033 Vl= XU 0039 V2= XlvH 004 0 V3= XI 0041 V4= K20BS 0042 Vl= V1/V4 004 3 V2= V2/V4 0044 93= V3/V4 C045 V4= V4/V4 0046 SUM 11 = Vl**2 + SU.Mll 0047 SUM 12 = VI =5=72 + SU.‘^12 0048 SUM 13 = Vl*V3 + SUM 13 004 9 SUM 14 = V 1<:V4 + SUM 14 " ... ' 1 0050 SUM22 = V2*»2 + SUM22 0051 SUM23= V2»V3 + SUM23 0052 SUM24= V2*V4 + SUM24 0Ü53 SUN 3 3 = 7 3**2 + SU.M3 3 0054 S U N 34= 73*74 + SUM34 0055 ioc CONTINUE 0056 A U f1) = S U M l 1 005 7 A 12, I) = SUM12 ...... 0058 A (3,1) = SUM13 0 0 5 9 A (1,2) = SU M 12 0060 A (2,2) = SUM22 0061 A (3,2) = SUM23 0062 A (1,3) = SUM.13 g. 0 0 6 3 A(2,3* = SUM23 0064 A(3,3) = SUM33 00àl> 0(11= SUM14 OObt) ü(2)= SUN?4 006/ ■ ■ ti(3)= SUM 34 1 0060 w.RI TE(6,?2) M A ( I, J) , I = l,N) , j=l,N) 1 (B{ I ) , I = liN) 0 0 6 9 21 H(lKMAr(/3[:i6.5J 0 0 7 0 CALL SIM.Q{A,B,N,0) OÙ 71 SUMD= 0.0 007;? ÜG 201 JJ= 1,NN 0073 KtADl5,3) EM.T f PHfK20l3S 0074 L-r:T= E N T/ 2 .0 0 0 7 ü il= iO.O*»(-PH) 0 0 7 6 XU= 1.07(1.0+ H/10.0*>M-PKAC1 ) + H**2/10.0»*(-PKACl - PKAC2 )) 007 7 ■ ■ Xl= XU'^=H/10.C^*(-PKACI) 0 0 7 0 X2= X1*H/10.0*»(-PKAC2)

" CÙ79 Vl= XU OOÔO V2= X 1 *11 ooül V3= Xl 0 0 0 2 V4= K2CBS 0003 Vl= V1/V4 0084 V2= V2/V4 0085 V3= V3/V4 0 0 6 6 VA= V4/V4 008 7 SfJ= ( Vi*l3( 1 1 + V2*B(2) + V3*B( 3) - V 4 1**2 0080 SUM.C= SUMD + SD 0089 2ül CONTINUE • ■ ...... 0 0 9 0 SON= N.\ ■ 0091... siL= Sur T ( sum.d /( suw — i.)) 0 0 9 ? WRITE(6,23) (B(I),l=l,N),STD 0 0 9 3 2 3 FORMAT(/4E16.5) 0094 5CÜ CONTINUE 0095 7/7 CG\TINUE 0096 STOP

VI « 506 SIMQ, the subroutine for solving an nxn matrix for n unknowns•

The program is part of the IBM-360 subroutine library and was used frequently. This copy of the program was obtained from the Systems Engineering Department. SI MO 001 c 002 c SI MO 003 c SUBKCUTINe SIHi; SIMQ 004 c SI MO 005 c PIJ^PÛSF SIMQ 006 c OBTAIN SOLUTION UF A SET ÜF SIMÜLTÂNFOUS LINEAR EQUATIONS, SI MG 007 c AX U SIMQ 008 c SI MO 009 c USAGE SI MO 010 c CALL SIKQ A,B,i\’,KS SIMQ Oil c SIMQ 012 c ÜESC«lPTION OF PARAMtreRS SIMQ 013 c A - f-ATRIX ÜF CCFFFICIENTS STORED CGLUMN’rtISE. THESE ARE SIMQ 014 c DESTROYED IN THE CONPUTATION. THE SIZE OF MATRIX A IS S I MU 015 c N HY N. SI MO 016 L b - VSCrCR OF ORIGINAL CONSTANTS LENGTH N . THESE ARE S I MO 017 c REPLACED BY FINAL SOLUTION VALUES, VECTOR X. SIMQ 010 c "k - NUMBEK 'OF OUATIONS' AND VAR I able S.'IJ MUST BE" .GT. ONE". ■“SI MO M o r c k S - OUTPUT DIGIT SI MO 020 c 0 FOR A NORMAL SOLUflUN Si MO 021 c L FOR A SINGULAR SET OF EQUATIONS SIMQ 022 c SIMQ 023 c REMARKS SI MO 024 c FAIR IX A MUST BE GENERAL. SIMQ 025 c IF MATRIX IS SINGULAR , SOLUTION VALUES ARE MEANINGLESS. SIMQ 026 c AN ALTERNATIVE SOLUTION MAY BE OBTAINED BY USING MATRIX SI MO 027 c INVERSION XTNV AND MATRIX PRODUC T GMPRD . SI MO 028 c SI MO 029 c SUBROUTINES AND FUNCTION SUBPROGRAMS REQUIRED SIMQ 030 c NONE S I MO 031 c SIMQ 032 NETHCn SI'/Q 033 C iVEniGO OF SOLUTION IS HY EL IMINATION USING LARGE SI PIVOTAL SIMO 034 C ü lVÏT(JRV'TÂcrri"'S^rÂGE 'OF fcLÏH ÏIMT I UN C OF I NTt RCHÀNGI NGS I KQ 0 35 c R O V i S WtiFN NECESSARY TO AVOID DIVISION BY ZERU UR SMALL______SIMQ 036 c VLbMENlS.' SIMQ 037" c THE FORWARD SULUIION TO OHTAIN VARIABLE N IS DOME IN S I M.Q 030 c N STAGE'S. THE BACK SOLUTION FOR THE OTHER VARIABLES IS SIMQ 039" c CALCULATED BY SUCCESSIVE SUBSTITUTIONS. FINAL SOLUTION SIMQ 040 "c VAŒËni<Ë"'DEVECOPED IN VÊCTOR' Br W ITH' VARIA IN 8 L V SI MO 041 c VARIABLE 2 IN B 2 ...... , VARIABLE iN IN B N . SIMQ 042 c IF NO PIVOT CAN BE FOUND Erc¥ED TNG "A TULER'ANC E OF' OTO SIMQ' 043" c THE MATKU IS CONSIDERED SINGULAR AMD KS IS SET TO 1. THIS SIMQ 044 c TOLERANCE CAN BE MODIFIED BY REPLACING THE. FIRST STATEMENT. SIMQ 045" c SIMQ 046 SIMQ 0 4 7 ■ SIMQ 046 TUBruutTNl 51 MG ( a'TbTN",K"ST “SI M Q" 049 dimension All),B( I) SI MO OSO ~ c "Ti'MQ'051' c FORWARD SOLUTION SIMQ 052 "SŒQ- 05T rUL=0.0 SI MO 054 KS = C ~~smî~urf5~ J J=-N SIMQ 056 UÜ 63 J =I » N 3 T M Q D 3 7 " JY=J+1 SIMQ 058 T 3'STr3 + N+"I ■"S1'FQ“ 0'5'9" b IGA = 0 SIMQ 060 T r = T J - ' j STM C T W r DO 30 I=J,N SIMQ 062 -\n 8 L bl MO U6 3 C StAfsCH FÜK MAXIMUM CCEFF IC I tfN I IN COLUMN S I MO 064 C SI MO 1)65 IJ= l T+I SIMQ 066 rFr?iTïS{f^riTûi-ATST'rrm 7 "i“ ? 0T 30V 3o- 'SIMQ' "067 2C BIGA=A(IJ) SIMQ 060 l(/AÂ=I ‘SÏKQ" TÔT 3C CCN7IN0E SI MQ 070 "C" SI MO 071” C TEST FOR PIVOT LESS THAN lOLERANCE SINGULAR MATRIX SIMQ 072 -STMGT T T 3 ” IFlABS(BIüA)-TCL) 35,35,40 SI MO 0 74 33 \S= I “si'Mir T73" RET URN S I MQ 076 L ' S I W 077" C INTERCHANGE ROWS IF NECESSARY SIMQ 070 X' "STMQ 070 ' 40 Il=J+N*(J-2) SIMQ 000 I“TFTT"Ain--- " S I W T o r 0 0 50 :< = J,N SIMQ 002 ---- SIMQ "003 12=11+1r SIMQ 004 S

, . 65 B{IX)=B(IXJ-(5(JJ«A(IXJ)) SIMQ 107 c SIMQ 103 ‘ c BACK SGLUriCN SI MO 109 c SIMQ HO ire NY=N-1 SIMQ 111 IT=N*N SIMQ 112 UU eo J=i,\' Y SIMQ 113 IA=1T-J SIMQ 114 IO = iN-J SI HO 115 IC = N SIMQ 116 DU 60 K=l,J , SIMQ 117 0 ( II3)=LM IO)-A( IA)=fB( IC) SIMQ 118 1A=IA-N SIMQ 119 00 IC=IC-1 SIMQ 120 RET0B.\! SIMQ 121 END SIMQ 122

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