THE THEORY OF ACIDS AND BASES

By F. M. HALL, )1.SC. Wollongong University College, N.S. W. , Australia

The theory of acids and bases, like many and the term hydrolysis relates to the inter­ other chemical theories, has undergone action of the ions of the salt with water to numerous changes in recent times. Always give (a) a weak acid or a weak base, or (b) a the changes have been such as to make the weak acid and a weak base. theory more general. The three main Applying the Law of l\llass Action to such theories in use today are: a system, the hych·olysis constant, at equili­ (l ) the Water or Arrhenius Theory; brium, may be written as (2) the Proton or Br0nsted- Lowr.v l [Base] [Acid] [OH-) [HA) Theory; (" = (Unhydrolysed salt] = [A- ] (3) the Electronic or Lewis Theory. and if x is the extent to which hydrolysis

WATEl~ OR AlmHENIUS 'l'HEORY occurs and c is the molar concentration of salt The Water or Arrhenius Theory was in solution widely accepted up to the early years of this xcx xc x 2c century. It defines an acid as a hydrogen K" = (1 - x)c = (1 - x) compound ionizing in water to give hydrogen The equilibria involved for the salt of a ions, and a base as a hysolutions we may replace]( and Kh' by two strong bases ; new constants, Kw and J(h respectively, (iii) those derived from strong acids and defined by weak bases; Xw = [H +) [OR-] (iv) those derived from weak acids and _ [.HA] [OR-] } 7 weak bases. and '-h - [A- ) When dissolved in water these various salts Kw is called the ionic product of water and do not necessarily give neutral solutions. 14 2 2 Interaction between the salt and water has a value of ,...._, 10- mole litre- at 25°C. (hydrolysis) accounts for the acidity, alka­ It follows that linity or neutrality of the solution. For K,. _ [H+] [OR-] [HA] _ [OH-] [HA) _ K example, a salt BA derived from a weak acid Ea - [H+) [A-] - [A ) - 11 HA and a strong base BOH gives an alkaline If two assumptions are now made: solution in water because the acid derived by hydrolysis gives few hydrogen ions and the (i) that there is only a pure aqueous base derived by hydrolysis gives many solution of the salt of the acid HA (no hydroxide ions. This may be shown as added hydrogen or hydroxide ions) ;

B+ + A -+ H 20 ""' IIA + B + + OR- and 91 92 EDUCATION IN CHEl\flSTRY

(ii) that the concentration of hydroxide A solution then is not necessarily neutral at ion obtained from the ionization of pH 7 but rather when the hydrogen ion con­ water is negligible compared with that centration is equal to the hydroxide ion resulting from the hydrolysis of the concentration. However, at 25°C a neutral salt, solution has a pH of 7 and all subsequent then [OH- ] = [HA] during hydrolysis of the salt, calculations refer to this temperature. [OH- F K w PROTO~ OR BR0NSTED- LO\VRY TTIEORY ancl [A-] = !{11 = K;; The Arrhenius Theory makes use of and [OH- ] = J [A -]. Kw hydroxide ions, which may not exist in non­ ] (,. aqueous solvents, and does not cover weak J(w bases. In 1923 Bnmsted and Lowry put but [H +] = [OH - J forward a more general theory of acids and bases which incorporates all protonic solvents, hence [H +] = w • JK [A-IC] and not just water. They defined an acid as a substance ·which yields a proton and a base or pH = ?-pKw + t piCa + -~ log c as a substance which can combine with a where c = [A-], the stoichiometric molar proton. concentration of the salt.* This is only Thus an acid HB dissociates to give a justified if the hydrolysis of the salt is very proton and its conjugate base. Alternately a small. base B can combine with a proton to yield the This equation then allows calculation of the conjugate acid BH+ of the base. In general pH of an aqueous soh1tion of a salt or the pH terms, this may be written as at the equi valence point of a weak acid- strong Acid 1 + Base 2 "" Acid 2 + J:3ase1 base titra.tion. Similarly for a strong acid-weak base in which the proton is pat•titioned between system the pH is given by two bases and the equilibrium constant is determined by the relat ive affinities for the pH = t pi(,,. - ·~· pKb - -~ log c proton. where J(b is the ionization constant of the A Br0nsted acid may be an electrically base and c is the molar concentration of the neutral molecule, e.g. HCI, a cation , e.g. salt. C6H5~TH3+ , or an anion, e.g. H S04- , whilst a It should be noted that ICv , the ionic Br0nsted base may be a neutral molecule, product of water, like any other equilibrium e.g . aniline C6H 5NH2 , or an anion, e.g. Cl-. constant, varies with temperature. Table I One important result of this definition of an illustrates this point. acid is that the strength of an acid depends upon its environment. The acidic strength TABLE I of a weak acid is etihanced by its solution in a strongly basic solvent and the basic strength pH of neutral solutiort of a weak base is enhanced by its solution in Temp. ' C Kw x 101'1 [H+] = [OH- ] ----·-1----- 1------a strongly acidic solvent . In fact, all acids 0 0·1139 7·47 tend to become indistinguishable in strength 10 0·2920 7·27 in strongly basic solvents. This is known as 20 0·6809 7·085 25 1·008 6·998 the levelling effect of the solvent. 30 1 · ~69 6·915 Solvents may be protophilic, protogenic or 40 2·919 6·77 aprotic. If a solvent exhibits both proto­ 50 5·474 6·63 60 9·614 6·51 philic and protogenic properties it is termed a.mphiprotic. Examples of each form are *pH is here defined as pH = - log[H +], i.e. in protophilic solvents-ethers, ammines terms of concentration rather t han activity of the (basic substances); hydrogen ion. This is done for the sake of sim ­ plicity in t his and in all subsequent calculations in protogenic solvents- sulphuric acid this article. (acidic substances); THE THEORY OF ACIDS A.1'

amphiprotic solvents-water, acetic acid, strengths of what, in water, are normally alcohols; termed strong acids are to be determined, aprotic solvents-benzene, chloroform then solvents other than water have to be ('inert' substances). used. This necessitates the use of other Actually, these definitions involve a modern acidity sca.les, the most familiar of which is extension of the Bn'insted- Lowry theory, probably the Hammett acidity scale. which might be termed the 'auto-protolysis For example, the acids perchloric, hydro­ theory,' viz. that in a solventS in which the chloric and nitric appear equally strong in equilibrium water where the strongest acid that can exist 2S .= A++ B- is the hydroxonium ion, H 30 +, but in other solvents jheir strengths differ because of occurs, an acid is any substance whose dis­ variation in the ease of formation of the solution increases the concentration of A+ and solvated proton in that particular solvent. a base is any substance which increases the The basicity and dielectric constant* appear concentration of B- . to be th<: principal factors in the manifesta­ Thus when water ionizes, the equation may tion of acidity. be written as Since the extent of acidic and basic dis­

H 20 + H 20 <=' H 3 0 + + OR- sociation is influenced by the ability of the where one molecule of water is behaving as solvent to accept or donate protons, the same solute may be dissociated to widely different an acid and another as a base (amphiprotic). degrees in different solvents. For example, Similarly for acetic acid ammonia is not highly protonatccl in water, CH3COOH + CH 3COOH <=' CH 3COOH2 + + but glacial acetic acid, a solvent more proto­ CJ:I. 3COO- genic than water, induces extensive protolysis. one acetic acid is acting as a proton donor .::\TH3 -! C H.3COOH .= ~ H, - + CH3COO­ (acid) and the other as a proton acceptor Again, sulphuric acid is over four hundred (base). times as acidic in acetic acid as in water at Again, ammonia ionizes as follows: the same concentration. Such systems are ~1-J3 + NH 3 <=' NH,,+ + .::\TH2 - often referred to as 'super-acid systems.' and acids arc those substances which, in P erchloric acid may be dissolved in acetic liquid ammonia, increase the concentration of acid to give a solution containing CH3COOH2+ ions, and, as this ion can readily give up a NH4+. It follows then that a substance which proton to react with a base, the solution is ftmctions as an acid in one solvent docs not strongly acidic. On the other hand acetic necessarily react in this way in another acid may itself donate protons to a suitable solvent. Urea, for example, is a weak base in base. This acidic property will exert a water, a stronger base in acetic acid and yet levelling e:ffect on a weak base, which will is an acid in liquid ammonia. thus have its basic properties enhanced. Thus the titration of perchloric acid with H 0 + H K·CO·NH <=' H ~ · CO·NH + ..L OH- 2 2 2 2 3 pyridine in water fails to give a satisfactory acid, -r base2 acid2 end-point, but the same titration in glacial CH3COOH + H 2N ·CO·NH2 <=' acetic acid is quite successful. The reactions

acid1 + base2 involved are

H 2N ·CO·l\TH + + CH COO - 3 3 HC10,1 + CH3COOH <=' CH3COOH2 + + Cl04 - aci d -+- base 2 1 C5H5N + CH3COOH <=' C5H 5NH + 7 CH 3COO­ ).1H + H ~ · CO · NH .=NJ:I,,+ H i\'"·CO ·NJ:-I - 3 2 2 + 2 CH3COOH2+ , CH3coo- <=' 2CH3COOH base1 acid acid base + 2 1 + 2 Adding, H C10 4 - C5H 5::\ <=' C5H 5NH+ -:- CL0 4 -

Strictly the pH unit is limited to dilute * Tho Dob~·e-Hu ckd equation allows t.he eiTecv aqueous solutions. The useful range of the of changes of dielectric con stant on t he fw1ctional relationship between the activity coefficien t of tho pH scale is 0 to ] 4 and is fixed by the auto­ ionic species and t.lw ionic strcngth t o h e taken into protolysis constant for water. If the relative .:u•count. 04 EDUCATIO~- IN CHEl\U STRY

The intrinsic strength of an acid HA in a To explain further point (iii), consider a particular solvent, then, is formally expressed weak acid, e.g. acetic acid, and a weak base, by its acidity constant in that solvent, and e.g. ammonia. the relationship between a base and its According to the Arrhenius theory conjugate acid makes it unnecessary to deal . . [H+) [CH3COO-] with basic dissociation constants when the K. (acecw ac1d) = CH COOH solvent is amphiprotic. 3 The Bnilnsted theory evidently differs from or pKa = 4· 76 the Arrhenius theory in the followi11g ways: [~H +) [OH-] Kb (ammonia) 4 NH40H (i) proton transfer processes only need be considered (protonation and d.eproto­ = 4 ·75 nation); However, in the Br0nsted sense there is no need to invoke the presence of molecular (ii) the term hydrolysis is no longer NH40H (which probably does not exist) to necessary; give pKb. Rather, using the conjugate acid of the base, i.e. NH4+ ~ H + + NH3, a pKa (iii) the value Kn can be neglected and the value of 9·25 can be obtained for the ammo­ strengths of all acids and bases can be nium ion, since pKa + p]{b = pKw = 14. given by the J(J. value alone by con­ The major weakness of the Br0nsted­ sidedng the conjugate acid of the base Lo"rry Theory is that it makes the definition and the value of the solvent ionization of an acid and a base as rigidly dependent constant (e.g. Kw). upon the solvent as does the Arrhenius To illustrate (iii) more clearly, a ca.lculation Theory. H owever, it does take into account involving the same system referred to under that there are solvents other than water t he Arrhenius Theory (p. 91) is more simply which exhibit typical basic properties, even derived from the Br0nsted approach. though it does not recognize the comple­ In the reaction mentary data with regard to acids.

HA + OH- ~ A- + H.O acid1 + base2 base1 + acfd2 ELECTRONIC OR LEWIS THEORY [HA] = [OH-] at the equivalence point for The Lewis Theory first introduced in 1923 there is only the salt of HA in the solvent. is a more embracing theory still than those of Arrhenius and Br0nsted. K = [HA] [OR-) = IC. It is an electronic theory and the definitions b [A- ] K ,. of acids and bases do not depend upon the presence of any particular solvent. since [H 0] is effectively constant as before. 2 An acid is defined as an electron-deficient species or one that seeks a molecular species containing available pairs of electrons. For

example, H +, N02+, BF3 and AlCl3 are acids. Bases are species which contain a pair of electrons capable of being donated to another .. pOH = t piC.- t pK, - t log [A- ) species. For example, 01- , H 20, OH-, :t\TH3, But pOH = piC. - pH (pH + pOH = pKw ethers, esters and ketones are bases. = 14 ac 25°C). Searching for a property common to all Hence acids, or that common to all bases, Lewis pKw - pH = tplCv - tpKa- tlog [A-) concluded that acids and bases correspond or pH = i pKw + t pK. + t log c respectively to what Sidgwick later called 'acceptor' and 'donor' molecules. Neutraliza­ where c is the molar stoichiometric concentra­ tion is the formation of a co-ordinate covalent tion of the salt. bond between the acid and the base. THE THEORY OF ACIDS A~D BASES 95 For example, more usual neutralization of pyridine by a proton-donating acid: H ++ :O: H - -+ H:O:H ac id base H +1 + ' -;/'~ -'; [n' -q--']+ 1 Cl H Cl H (froma '=/ ~ I I I I proton donor) Cl- H + :~-H -+ Cl- B:N-H I I I I Cl Cl Cl H Cl H I I~ -L, • , .:?-,. ""' CI- A! -+ Cl- AI: ~· ) The base donates a share in a lone pair of I . -''==/ I "=/' electrons to the acid to form the co-ordinate Cl Cl covalent bond bct,Yeen the two. Formation Agai n solutions of boron trifluoride or of this bond is always to be considered the sulphur trioxide in inert solvents bring about first step, even though ionization may subse­ colour changes in indicators very similar to quently take place. In the above case of those produced by protonic acids. These boron trichloride and ammonia, ionization changes can be reversed by adding bases so does not occur after neutralization. In othct· that a titration is possible; yet no proton is cases, however, the electrical 'strain' produced involved. by the formation of the co-ordinate covalent The major di ·advantage of the Lewis bond is sufficient to result in ionization of the system appears when its quantitative aspect neutralization product, as when aluminium is considered. The protonic acids make up a bromide reacts with pyridine. group which show greater uniformity than do

~~· #~ [ ~.r -~ Br ]+ Br:~ + :K~= Br :~.l: <=> l = :~r: ~~=~J Br Hr _j ~J3r Lewis was concerned with broadening the the non-protonic acids of the Lewis definition, basis of the acid- base definition from both when related to any simple system of acid­ the experimental and theoretical standpoints. base strengths. He chose four familiar experimental criteria, A ftu·ther disadvantage is that certain sub­ viz. neutralization, titration with indicators, stances which experimentally behave like

displacement and catalysis, and defined as acids, e.g. HCl and C02, have electronic acids and bases all substances which exhibit formulae which, as usually written, do not the ability to take part in these 'typical' show the possibility of their acting as electron­ functions. On the theoretical side he related pair acceptors. Such acids and bases are these properties to the acceptance and dona­ called 'secondary' by Lewis as distinct from tion of electron pairs irrespective of whether his 'primary' acids and bases which involve the transfer of protons was involved. electron-pair sharing. This introduces a The scope of the electronic theory is suffi­ cumbrous name fo1' common substances and ciently broad to include the proton-transfer raises the question of the value of the term definition of an acid, and, since electron­ 'acid' as commonly used. donor molecules are able to combine with protons, the Lewis concept of a base embraces CALCULATIONS ni' ACID- BASE TITRATIOXS the 13r!1nsted-Lowry definition. On the other Each of the above theories presents certain hand the Lewis definition of an acid embraces difficulties, and a practical point at issue is many substances which do not contain a which theory is able to clearly and quantita­ hydrogen atom, and consequently radically tively assist in calculations involving acid­ increases the number of acids over those as base systems. The Arrhenius Theory is defined by the Br0nsted concept. restricted to water as the solvent. The Lewis For example, it is evident that precisely the Theory covers more completely substances same principles are involved in the reaction of that show the qualitative attributes normally aluminium chloride with pyridine as in the associated with acids. The Br0nsted-Lowry 96 EDUCA'l'lON IN CHEMIS1'RY acids, on the other hand, form a more uniform protogenic character of the conjugate acid of group and obey the quantitative relationships the base. Beyond t he equivalence point confined to this group. (pH < 7) the pH change is due only to the For this reason, and because a solvent is d il ution effect of the strong acid. normally used in simple acid- base systems, The pH values can be calculated from the the Br~nsted-Lowry approach is generally following equations: used in the following calculations of acid- base Before addition of any acid: t itrations. pH '--' ! pJ\",.. -t t pE~ + } log c (base) l. St1·ong Acicl - Strong Ba.se 'l'itmtion Cp to the equi,·alenco point: The pH change is due only to the clilut ion H _ -r c (salt) 1 effect and the neutralization of some of the P p1 \.. w - P1 ~,, - ogc(baso) acid or base. l3oth acid and base are fu lly c (salt) dissociated, and calculations involve a sess­ - pJ(,.b - log c (base) ing the actual acid or base concentration A t the cquivttlenco point: before and after the equivalence point. The cqui,,alcnce point in water is at pH 7 at 25°0. pH = t pi\:,. - { pKb - } log c (salt) = ~ pKab - } log c (salt) Example: If to 50 ml of O·lN HOI, 20 ml of O·lN NaOH are added, the total acid con­ where p Ko.b refers to the conjugate acid of the cent ration is 30 ml of O·lN in a volume of base and is equal to pJ(,.- pKb. 70 ml. 4. W ealc Acid- W ealc Base 'l'itmtion [H +) = ~ X 0·1 or pH = 1·37 70 The pH change tru·oughout is due to If to 50 mlofO·lN H OI, 51 mlofO·l N NaOH clilution, neutralization and the proton are added, the total hyru·oxide ion concentt·a- affinities of the conjugate acid and conjugate 1 base. This t,vpe of titration is rarely u ed, tion is X 0·1, or pOH = 4·0, or pH = 1 0·0. 101 but calculations may be made from the equations under (2) and (3) above, except 2. Weal.; .d cicl-Strong Base Titmtion that at the equh·alence point itself the pH is The pH change, up to the region of the given by equivalence point, is due to the dilution pH = !-pl(,. + 4 pi{.--}plCb = -,tpK,. - ~ pKab effect, neutralization of the acid and the protophilic character of the conjugate base. where pKo.b again refers to the conjugate acid Beyond the equivalence point the pH change of the base. is due only to the dilution effect of the strong base. The equivalence point occurs at a pH 5. Polyprotic rl cid8 wul Bc£ses greater than 7 ·0 at 25°0. In a polyprotic system, where transfer of 'l'he pH values can be calculated from the more than one proton is involved, e.g. in following equations: H 3X., there are three dissociation constants: Before any base is added: - [H +) [H X -) pH = t pK. - ·i log c (acid) 1 2 )..1 = [H 3X) Titration up to equivalence point: ? (H+) (H...'\:.Z-) c (sal~ ) l"z [HzX ) pH = pJC., , log c (acid) = , [H+] (X3-] A~ ~he equh·alence point: il..3 = [HX2 ] pH = t pK. + ~ p i(,,. + t log c (salt). Provided pK1, pK2 and pK3 differ by at least 3. Strong .Acid- lVeak Base Titmtion four units, then each step, im·olving one The pH change, up to the region of the proton transfer, proceeds ,-irtually to com­ equivalence point, is due to the dilution pletion (99·99 per cent for four units) before effect, neutralization of the base and the the next neutraLi zation ~tep commence:>. THE THEORY OF ACIDS AND BASE S 97 Such a titration will give two or three inflec­ and arises because the ratio in the concentra­ tion points, and the pH at the first inflection t ions of the two coloured forms of the indi­ point is given by cator will vary continuously as the hydrogen­ ion concentration is changed. pH = -} (pK + pJ(2) 1 For ease in observing colour change, which corresponds to the pH of the salt screened or mixed indicators may be used. NaH2X. A screened indicator contains an indifferent The pH at the second inflection point is dye which allows the colour change to be given by more easily seen. A mixed indicator con­ sists of a main indicator and an auxiliary pH = ~· (pK2 + pK ) 3 indicator which indicates the approach of the and corresponds to the pH of the salt Na2HX. change point, or two indicators with over­ The appearance of a third inflection point lapping pH change_ The advantage of this in the titration of H 3X with alkali depends latter type is that a sharp colour change upon t he pH of the solution of the salt Na3X occurs over a more limited pH range. in relation to the pH of the solution when a For a substance or system of subst ances to slight excess of the alkali titrant has been function satisfactorily as a pH indicator, the added. change observed should satisfy the following conditions: DETECTION OF E:SD-POI:STS TI ACID- B ASE (i) the change should be a distinct one TITRATIO~S and should occur over the shortest In all titrations some means of detecting possible pH range; the end-point or end -points is necessary. In (ii) the indicator change should be rever­ acid- base systems two distinct techniques sible and the reaction involved should are commonly used: be rapid in both directions ; (i) instrumental, e_g_ pH meter, con­ (iii) the indicator should be sensitive, i .e. ductivity cell ; only a small quantity should be (ii) indicators. required to impart a distinct colour to the solution. Indicators which cover a wide range of the pH scale are available, and it is necessary to Common pH indicators, their colours and select one whose colour change occurs at a pH ranges are given in Table II. pH attained immediately after the equiva­ lence point is reached, that is, at the point 'l'ABLE II wh0re the rat e of change of pH with change of volume of titrant is at a maximum. Naturally Colour pH thiil does not necessarily occur at pH 7 at 25°0. ~arne Acid Alkaline 1·ange Indicators a.re themselves weak acids or bases, and consequently the ratio of ionized Thy mol b lue - - red yellow 1·2-2·8 :i\.fethyl orange . - red orange . 3·1- 4·4 to unionized form depends on the pH. For yellow simplicity they may be considered as acting Bromo cresol bluo. _ yollow b lue 3·8- 5·4 Methyl red _ . -- red yellow 4·2- 6·2 like any ot her weak acid or base, that is B romothy m ol bluo yellow b lue 6-0- 7·6 Cresol red -- 0 0 yellow red 7·2- 8·8 Hin ~ H + + In- Phenolphthalein 0 0 colourless red 8·2-10·0 where H ln and In- are of different colour in Thy mol pht halein 0 0 colourless blue 9·3-10·5 the solution. The chief characterist ic of these indicators is that the change from a predominantly SUMMARY 'acid' colour to a predominantly 'alkaline' Because of the gradual fusion of ideas on colour is not sudden, but takes place over a acid- base t heory from which our present pH range of about two units. This is termed concepts have sprung, the importance of this the colour-change interval of the indicator topic has extended far beyond the reaction of 98 EDUCATIO::< IN CHEl\'IISTRY

an acid and a base to give a salt and water. 'wrong,' forgetting thn,t they are only con­ Indeed calculations of pH in acid- base venient, altogether artificial schemes for systems, as described above, are only a part classifying systems. No one of them is either of its usefulness. true or false and the differences are in degree, The Br0nsted-Lowry definition has not kind. The choice between them should initiated many investigations of acid- base depend solely upon the region of chemistry equilibria and kinetics in different solvent s, in which one is operating. The water theory whilst the Lewis concept has led to much is applicable to aqueous solutions, the proton valuable work on t he reactions of acceptor theory is preferable for dealing with a variety molecules. of solvent s, whilst the electronic theory The following are but a few examples of covers acid behaviour in the absence of studies involving acid-base theories and. protons.

serve t o illustrate the scope of this field. : FURTH ER READING (a) the forma.tion of the so-called Bates, R. G., Elect•·omet?-ic pH Detennination-~ . hydroxides of iron and other metals; Chap man & H all, L ondon, 1954. Bell, H.. P., The P•·oton in Ghemist•·y. Methuen & (b) determination of equilibrium and Co., London, 1959. stability constants of simple and Boll, R P., Qtw.1·t. R ev. Ghem. Soc., 1947, 1,113. Becket, A . H., and Tinley, E. H., Titmlions in complex molecules, both organic and Non-Aqtwous Solvents (bookleL). Third Edition. inorganic; British Drug H ouses Ltd ., 1960. Bryson, A., Background to GhemisM·y (Editor: D. 1'. (c) the role of sulphmic- nitric acid in :iHellor), pp. 54-63. Tho University of No"· nitration processes; South Wales, 1960. Gold, V., pH .M.easunments: Thei•· Theory and (d) the use of lithium aluminium hydride Pmctice. Methuen & Co., London, 1956. in the reduction of organic compounds. Lucier, W. F ., and Zuffanti, S., ElecM·onic :L'heory of Acids ctnd Bases. J ohn Wiley & Sons, New York. There is, of course, a temptation when 1946. Vogel, A. I., T extbook of Quantitative I norganic dealing with rival points of view to assume Analysis. Third Edition. L ongmans, Green & that one of them is 'right' and t he others Co., London, 1961.

COURSES AND SYl\1POSI A FOR TEACHEl~S The following symposia have been arranged Campbell, Director, CHE.M Study P ro­ by Local Sections of the Royal Institute of ject.) Chemistry in collaboration with the Associa­ tion for Science Education: The following information has been received University of Liverpool, 25 April, 1964-­ from Loca1 Education Authorities about 'New Techniques in Practical Chemistry courses for teachers that are being arranged for Schools.' during the Summer Term : The School of P harmacy, University of Cambridgeshire Local Education Authority London, 2 May, 1964-'The Teaching of (in collaboration with the Ministry of Inorganic Chemistry at Pre-University Education), 13- 17 April, 1964--'Science Level.' Teaching in t he Secondary School.' University of Leicester, 2 May, 1964--'The Surrey Local Education Authority (King­ Teaching of Chemistry in Schools.' (A ston College of Technology), July, 1964- follow-up conference arising from the 'New approaches to teaching theoretical symposium on the teaching of organic and practical chemistry in secondary chemistry held on 26 October, 1963.) schools.' Manchester College of Science and Tech­ Somerset Local Education Authority (in nology, 9 .May, 1964--'CHE.M Study-a collaboration with the Ministry of Educa­ new look at the teaching of chemistry.' t ion), 27 April-1 May, 1964--'Science (Guest Speaker, Professor J. Arthur Teaching in the Secondary School'