ACID-BASE EQUILIBRIA
INTRODUCTION
How about the chemical workers - are they unionized or ionized?
What are acids and bases? Svante Arrhenius noticed that acids release hy- drogen ions in solution. He classifi ed acids and bases in this way: acids are compounds that dissociate, or break up, in water to give a proton; bases are compounds that dissociate in water to accept a proton. Acids that break up completely in water; these are strong acids. Acids that do not dissociate completely in water are call weak acids. Can you guess what bases that dissociate completely are called? Do you think weak bases dissociate completely? Things that don’t dissociate completely are most interesting - and often overlooked when it comes to safey! Take vinegar, for example. Did you know that this common kitchen item is actually a weak acid called acetic acid? While it’s not dangerous in the form you buy at the grocery store be- cause it’s very dilute (about 4-5% acetic acid), more concentrated solutions of acetic acid can be quite dangerous even though it’s a weak acid. Take a look at the dissociation of acetic acid in water.
+ - CH3COOH H + CH3COO Consider an acetic acid burn on your skin. Because acetic acid is neutral, it can be absorbed into the skin readily. Now you’ve got acetic acid in your hypodermis - the lower layer of your skin! If it were to remain neutral and not dissociate, this wouldn’t be a problem, but it does dissociate. The ions then attack the layers of skin and burn it. All the while, more weak acid can enter the hypodermis causing further damage, and a very severe burn. ACID-BASE EQUILIBRIA
ELECTROLYTES
A simple conductivity experiment is the following: A plastic screwdriver handle is placed against the electrodes - no light! When the metal part is used - bulb lights brightly! The handle is an insulator, whereas the metal is a conductor! 1.1 1.2
We can do the same type of experiment with solutions. Solutions are homogeneous mixtures of a solute dissolved in a solvent (module #). When substances are dissolved in water, the ►The diagram below represents the apparatus solutions can be tested by seeing if the solu- with electrodes dipped into liquid water. Is liquid tion conducts electricity, just as the metal of the water a conductor? Look at the top left-hand screwdriver does. An electrolyte is a substance corner of the conductivity experiment (the light whose aqueous solutions conduct electricity. bulb) to answer the question. Electrolytes are often identified by experiments testing the conductivity of solutions.
1.3 Bulb Switch
Filament Alternating Current
Solution Electrodes Sample
Sample H2O
In the following pictures, you will be asked to Be aware of the potential danger of alter- determine if the solutions are conductors or non- nating current electricity. Do not attempt conductors. These pictures will help you realize any of these experiments without appropri- the defi nitions of some key chemical terms, and ate supervision. hopefully some chemical concepts!
1 ACID-BASE EQUILIBRIA
ELECTROLYTES
1.4 1.5 A Sodium Cloride Solution. Pure Distilled Water. Solute: NaCl (aq)
Na+
H2O Cl-
Is the solution a conductor? ►Note that the bulb is not lit. State ► Name the solvent. whether the liquid is composed of Are atoms, ions, or molecules dissolved atoms, ions, or molecules. in the solvent? Name them.
1.6 A Sucrose Solution. 1.7 A Hydrochloric Acid Solution. Solute: HCl Solute: C12H22O11 (aq) (aq)
C12H22O11 H+
Cl-
►Notice that the arrangement of water ►Is an aqueous sucrose solution a molecules is different around different ions. conductor? Explain the reasons for the differences. Is the solute covalently bonded or Name the anion and write the symbol of ionically bonded? the anion. What charge does the hydrogen ion have? Name the cation.
Solutes dissolved in water to give solutions that conduct are called ELECTROLYTES. Solutes dissolved in water to give solutions that do not conduct are called NON-ELECTROLYTES.
►List the electrolytes for pictures 4-7. ►List the non-electrolytes for pictures 4-7.
2 ACID-BASE EQUILIBRIA
ELECTROLYTES 1.8 An Ammonium Cloride Solution. 1.9 A Carbon Dioxide Solution. Solute: NH Cl Solute: 4 (aq) CO2 (aq)
H+ + NH4 CO2
- - Cl HCO3
►Name the ions in solution. Which species ►Name the ions in solution. is the anion? Which species is the cation? Name the molecules in solution. Is the solute an electroyte? Is the solute an electrolyte?
1.10 0.7 M Sodium Hydroxide Solution. 1.11 0.7 M Acetic Acid Solution. Solute: CH CO H Solute: NaOH (aq) 3 2 (aq)
H+ Na+
CH3CO2H OH- ______
►Write the formula for the anion. ►Is the solute completely dissociated? Write the formula for the cation. Complete the list above of the chemical spe- Write the formula for the solute. cies in solution. Write the formula for the solvent. Let’s do some ChemLogs in order to further Let’s do some ChemLogs in order to further understand strong electrolytes. understand weak electrolytes.
+ - + - NaOH Na + OH CH3COOH H + CH3COO
NaOH CH3COOH
+ CH3COOH Na + - H OH - CH3COO 10 9 8 7 6 5 4 3 2 1 0 1 2 3 4 5 6 7 8 9 10 10 9 8 7 6 5 4 3 2 1 0 1 2 3 4 5 6 7 8 9 10 (x10-1 mol L-1) (x10-1 mol L-1) Dissociation is the separation of ions that occurs ►Why is the equilibrium arrow used for this when an ionic compound dissolves. weak electrolyte dissociation? Notice the solutes which dissociate completely cause the light to shine brightly. The dim light ►List the weak electrolytes. List the strong is caused by partial dissociation. Solutes that electrolytes. dissociate completely are called STRONG An ACID is a proton (H+ ion) donor. A BASE is ELECTROLYTES. Solutes that dissociate only par- a proton acceptor. This is called the Bronsted- tially are called WEAK ELECTROLYTES. 3 Lowry system. ACID-BASE EQUILIBRIA
ELECTROLYTES ►Fill in the “empty” ChemLogs for each of the ► Complete the chemical reaction and the following experiments: chemLog where necessary.
1.12 0.8 M Sulfuric Acid Solution. 1.13 0.8 M Potassium Chloride Solution. Solute: KCl (aq) Solute: H2SO4 (aq)
+ H+ K
2- - SO4 Cl
+ 2- + Cl- H2SO4 2H + SO4 KCl K +
H2SO4 KCl
10 9 8 7 6 5 4 3 2 1 0 1 2 3 4 5 6 7 8 9 10 10 9 8 7 6 5 4 3 2 1 0 1 2 3 4 5 6 7 8 9 10 (x10-1 mol L-1) (x10-1 mol L-1)
1.14 0.8 M Nitric Acid Solution. 1.15 0.8 M Potassium Hydroxide Solution. Solute: HNO3 (aq) Solute: KOH (aq)
H+ K+
- - NO3 OH
KOH HNO3
HNO3 KOH
10 9 8 7 6 5 4 3 2 1 0 1 2 3 4 5 6 7 8 9 10 10 9 8 7 6 5 4 3 2 1 0 1 2 3 4 5 6 7 8 9 10 (x10-1 mol L-1) (x10-1 mol L-1)
►Strong electrolytes that dissociate completely pH is a measure of the concentration of H+. + to give H ions are called strong acids. Name pH = –log [H+] the strong acids on this page. When pH is between 1 and 7, the solution is ►Strong electrolytes that completely dissociate acidic. When the pH is between 7 and 14, the + to give a H acceptor are called strong bases. solution is basic. The pH of pure water is 7 (neu- Name the strong bases on this page. tral pH). There are other solutes that are strong elec- ► ►For each of the solution on this page, state trolytes. These are called salts. Name the salts whether the ph is less than 7, greater than 7, or on this page. 4 equal to 7. ACID-BASE EQUILIBRIA
ELECTROLYTES
►Fill in the “empty” ChemLogs for each of the following experiments. Include the chemical reaction and the appropriate arrow for the reaction.
1.16 1.17 0.8 M Nitrous Acid Solution 0.8 M Ammonia Solution Solute: HNO2 (aq) Solute: NH3 (aq)
NH3 HNO2
- + NO2 NH4
- H+ OH
HNO2
10 9 8 7 6 5 4 3 2 1 0 1 2 3 4 5 6 7 8 9 10 10 9 8 7 6 5 4 3 2 1 0 1 2 3 4 5 6 7 8 9 10 (x10-1 mol L-1) (x10-1 mol L-1)
1.18 0.8 M Caffeine Solution ►Weak electrolytes that dissociate to give H+ ions are called weak acids. Name the weak Solute: C8H10N4O2 (aq) acids on this page. C8H10N4O2 Weak electrolytes that dissociate to give H+ ion acceptors are called weak bases. Name the weak bases on this page. + C8H10N4O2H ►On this page, which solutions are acidic? Is the pH of these solutions less than 7 or greater than 7? OH- Which solutions are basic? Is the pH of these solutions less than 7 or greater than 7? Are there any solutions on this page that have a pH equal to 7?
►For review, define a strong electrolyte and a weak electrolyte. 10 9 8 7 6 5 4 3 2 1 0 1 2 3 4 5 6 7 8 9 10 Can a weak base be a strong electrolyte? (x10-1 mol L-1)
5 ACID-BASE EQUILIBRIA
CONJUGATE ACIDS AND BASES
What do you notice about the arrangement of ACIDS ► the water molecules around the H+ ion in picture We’ve already seen the following reaction earlier, 2.1? but let’s study it more.
+ - The oxygen atoms in the water molecules have a HNO H + NO2 2 slightly negative charge (negative dipole). The H+ 2.1 Nitrous Acid Solution. ion is actually bonded to several water molecules. As you can see from the picture, a more complete way of writing the reaction of the dissociation of nitrous acid is - HNO2 + NO2 + - + 5 H O H O NO HNO2 2 11 5 2 + + H +
Obviously, this could get complicated (and long!)
very fast. A more common way - to We know that HNO2 is an acid because it gives a + NO2 + H ion; + HNO2 + H2O H3O + NO - HNO2 H + 2 - We call the H O+ ion the hydronium ion. However, at equilibrium, NO2 accepts a proton to 3 - Note: Water must be acting as a base in this reac- become HNO2. That makes NO2 a base; tion (see next section). + NO - H + 2 HNO2 - So, HNO2 is an acid, and NO2 is its conjugate - We chemists usually just write the most simpli- base. On the other hand, NO is a base, and + 2 fi ed version using H . You may notice that some HNO is its conjugate acid. 2 books do use the hydronium ion. As a student, it is important to recognize that these two versions The species formed when an acid loses a proton is indicate the same thing - a proton surrounded by called a CONJUGATE BASE. The species formed water molecules. when a base gains a proton is called a CONJU- GATE ACID. BASES
A simple way to remember how to identify acids When NH3 (g) is dissolved in water, it acts as a and their conjugate bases, and bases and their base. Bases accept protons. + OH-
conjugate acids, is the following table: + NH3 + H2O NH4 2.2 + Notice that H2O is acting as an acid. We cannot
exclude H2O as a reactant because we need to from the acid, conjugate base + H+ have a balanced chemical reaction. substract H+
+
to the base, conjugate acid add H+
►Identify the acid and conjugate base in the fol- lowing reaction: CH COOH H+ + CH COO– 3 3 6 ACID-BASE EQUILIBRIA CONJUGATE ACIDS AND BASES Now that we’ve seen that acids in water give hy- The dissociation of water gives both an acid drogen ions, let’s look at pure water. We’ve stud- (substance that provides a proton) and a base ied the following picture earlier and determined (substance that accepts a proton). We call water that water is a non-electrolyte because the light AMPHIPROTIC. bulb didn’t light. While all of this is true, there’s In the general equilibrium module, we learned to write more to this story. equilibrium constants. For water, ►For a review, state the defi nition of a non- [H+][OH-] electrolyte. K = c [H O] Pure Distilled Water 2 2.3 Since the concentration of H2O in water is 55.6 M, it can be regarded as a constant. A simplified
version of Kc is called the ion-product constant of water. H2O + - Kw = [H ][OH ]
Kw is a constant. At 25°, the concentrations of both H+ and OH– are 1 x 10-7 mol L-1. K = (10-7)(10-7) = 1 x 10-14 When a very sensitive light is used in the conduc- w tivity experiment, the light lights. How can this Note that there are no units associated with equi- be? Well, pure water does form ions in solution; librium constant values, K. This is standard, and it forms protons and hydroxide ions. But, the con- you are not expected to show units. centration of these ions is so low that a sensitive light is needed in the conductivity experiment to Will the value of K change at varying tem- see evidence of the ions. To show this in the form ► w of a reaction, we write peratures? Explain.
+ – Ion Product of Water. H2O + H2O H3O + OH
º C Kw pKw 0 0.12 x 10-14 14.93 + + 10 0.29 x 10-14 14.53 25 1.01 x 10-14 14.00 which can be simplifi ed to + – Remember the defi nition of pH from earlier in this H2O H + OH module. pH = – log [H+] For most purposes, we can simplify pictures of solvent molecules. The neutral water molecules This calculation can be used for more than just the in the following picture are simplifi ed to a blue concentration of H+. The pX of X (where X is any background. The ions in the water are thus shown number) is equal to the –log of X. more clearly. Other examples include: 2.4 pOH = – log [OH-]
pKw = - log Kw H+
- OH ►Determine the pH, pOH, and pKw for the dissociation of water. Check your answers by
using the fact that pH + pOH = pKw.
►What is the pH of a solution in which the pOH = 4? ►Water is an example of a molecule which gives Which is greater for this solution - the value of [H+] or the value of [OH–]? 2 conjugate acid/base pairs. What are they? 7 ACID-BASE EQUILIBRIA
CONJUGATE ACIDS AND BASES
►For each reaction below, identify the acid, ►Write the appropriate Ka or Kb for each of base, conjugate acid, and conjugate base. the reactions involving ammonia at the left.
For each reaction, we must determine whether + – NH4Cl NH4 + Cl the solvent (water) is acting as a base or an acid. NH + NH + H+ Look back at your acid-conjugate base and base- 4 3 conjugate acid designations for a reminder. 2.5 For the dissociation
+ + NH4 NH3 + H
NH3 we write the acid dissociation constant called Ka.
+ [NH ][H+] NH4 3 Ka = + [NH4 ]
►Is this solution acidic or basic? For the reaction
+ - NH3 + H2O NH4 + OH
we write the base dissociation constant called Kb.
+ - + - NH + H O NH + OH [NH4 ][OH ] 3 2 4 K = 2.6 b [NH3]
An important relationship for (and only for) conju- NH 3 gate acids and bases is Kw = Ka x Kb.
+ + - + NH [NH4 ][OH ] [NH ][H ] 4 x 3 Ka x Kb = [NH ] + 3 [NH4 ]
+ - -14 = [H3O ][OH ] = 10 = Kw
►Is this solution acidic or basic?
+ -9 ►Ka for NH4 is 1.0 x 10 . Calculate the Kb for
NH3.
+ ►What is the pKa for NH4 and the pKb for NH3?
►Another important relationship for conjugate
acids and bases is that pKa + pKb = 14. Prove this relationship.
8 ACID-BASE EQUILIBRIA
INDICATORS
One way to determine the pH of a solution is to use an indicator. An indicator exists in different colored forms depending on whether the compound is protonated or unprotonated. The indicator shown is purple cabbage extract. 3.1
H+ +
C H N O + - 15 15 3 2 H + C15H14N3O2 HIn H+ + In–
Rearranging the K expression gives ►Is purple cabbage indicator an acid or a a - base? Ka [In ] = + + - [H ] [HIn] ►Write the Ka expression for HIn H + In . 3.2
Water molecules HIn In- H+
►Which part of the solution causes the color, the solute or the solvent? ►The 1 x 12 array shown below shows the pH where the color changes. Mark which pH values show [HIn] > [In-]. Mark which pH values show [HIn] < [In-]. Mark which pH values show [HIn] = [In-]. 3.3
pH = 4 5 6 7 8 9 - - K [In ] + a = Therefore, when [HIn] = [In ], = 1, so, Ka = [H ]. [H+] [HIn] K [In-] - a 10 Ka [In ] 10 To see “pure” color, = = for blue. For “pure” purple, = = [H+] [HIn] 1 [H+] [HIn] 1 K + a + 10 Ka = [H ] = [H ] 10 Therefore, the pH range of an indicator is pKa 9 +/– 1. ACID-BASE EQUILIBRIA
INDICATORS
Here is a list of several acid/base indicators and the pH range in which they are useful:
pH Ranges of Indicators 3.4 pH 0 1 2 3 4 5 6 7 8 9 10 11 12 13 14 Malachite green Cresol red Metacresol purple Thymol blue Orange IV Methyl yellow Bromophenol blue Congo red Methyl orange Bromocresol green Methyl red Chlorophenol red Litmus para-Nitrophenol Bromocresol purple Bromothymol blue Neutral red Phenol red Cresol red Curcumin α meta-Cresol purple Thymol blue Phenolphthalein Thymolphthalein Alizarin yellow R Curcumin ß Malachite green Clayton yellow Various indicators Thompson-Markow Universal Indicator [H+] 10-1 10-2 10-3 10-4 10-5 10-6 10-7 10-8 10-9 10-10 10-11 10-12 10-13 10-14
A universal indicator, shown at the bottom, is a mixture of several indicators. Universal indica- ►What is the pH range of the Thompson-Mar- tors have a wider range of usefullness because kow universal indicator? + they mix indicators. What is [H ] at pH 0? ►The Thompson-Markow Universal Indica- tor is a mixture of six indicators. Metacresol purple is one of them; name the 5 others. 10 ACID-BASE EQUILIBRIA
THOMPSON-MARKOW UNIVERSAL INDICATORS
3.5 Thompson-Markow Indicator
Individual indicators that make up the Thompson-Markow Universal Indicator:
meta-Cresol purple
Methyl orange
Bromocresol green
Bromothymol blue
Phenolphthalein
Alizarin yellow R
pH 1 2 3 4 5 6 7 8 9 10 11 12
[H+] 10-1 10-2 10-3 10-4 10-5 10-6 10-7 10-8 10-9 10-10 10-11 10-12
►Determine the Ka and pKa of each indicator that makes up the Thompson-Markow Universal Indicator.
3.6 A more accurate way to determine the Ka and pKa of each indicator is to decrease the change in pH each pH interval. For example, instead of measuring 3.0 3.1 3.2 3.3 3.4 3.5 3.6 3.7 3.8 3.9 4.0 4.1 methyl orange from pH 1 through pH 12, we could measure it from pH 3 to pH 4.1. Of course, you have to know the general range of where the pKa lies in order to do this properly. pH 1 2 3 4 5 6 7 8 9 10 11 12
►Which indicator would be used to determine the pH of a 7.6 buffer solution other than the Thompson-Markow Universal Indicator?
11 ACID-BASE EQUILIBRIA
INDICATOR WELL TRAYS
Below are several chemLogs showing the pH and pOH for each well. Each well contains a solution of methyl orange at the differ- ent pH values listed.. 3.7
pH pH 1 1 2 3 4 5 6 7 8 9 10 11 12 13 14 ►What can you determine 14 13 12 11 10 9 8 7 6 5 4 3 2 1 about the relationship between pH pOH 1 2 3 4 5 6 7 8 9 10 11 12 13 14 pH and pOH? 2
14 13 12 11 10 9 8 7 6 5 4 3 2 1 pOH pH 1 2 3 4 5 6 7 8 9 10 11 12 13 14 3 ►Complete the following 14 13 12 11 10 9 8 7 6 5 4 3 2 1 pOH equation: pH 1 2 3 4 5 6 7 8 9 10 11 12 13 14 4 pH + pOH = ______
14 13 12 11 10 9 8 7 6 5 4 3 2 1 pOH pH 5 1 2 3 4 5 6 7 8 9 10 11 12 13 14 ►What is the pKa of methyl orange? 14 13 12 11 10 9 8 7 6 5 4 3 2 1 pOH pH 6 1 2 3 4 5 6 7 8 9 10 11 12 13 14
14 13 12 11 10 9 8 7 6 5 4 3 2 1 pOH pH 1 2 3 4 5 6 7 8 9 10 11 12 13 14 7 ►The approximate pH range of any indicator is the 14 13 12 11 10 9 8 7 6 5 4 3 2 1 pOH range between the pKa – 1 pH 1 2 3 4 5 6 7 8 9 10 11 12 13 14 and the pKa + 1. What is the 8 pH range of methyl orange?
14 13 12 11 10 9 8 7 6 5 4 3 2 1 pOH pH 1 2 3 4 5 6 7 8 9 10 11 1213 14 9 Remember that K x K = 14 13 12 11 10 9 8 7 6 5 4 3 2 1 ► a b pOH Kw for any conjugate acid/ pH 1 2 3 4 5 6 7 8 9 10 11 12 13 14 base pair. What is the pKb for 10 methyl orange? Using pKb, 14 13 12 11 10 9 8 7 6 5 4 3 2 1 determine the pOH range for pOH methyl orange. pH 1 2 3 4 5 6 7 8 9 10 11 1213 14 11
14 13 12 11 10 9 8 7 6 5 4 3 2 1 pOH pH 1 2 3 4 5 6 7 8 9 10 11 1213 14 12
14 13 12 11 10 9 8 7 6 5 4 3 2 1 pOH
12 ACID BASE EQUILIBRIA TITRATION A solution of strong base, NaOH, can be titrated Five drops of 0.01 M HCl were added to each of with a solution of strong acid, HCl. The overall the wells in the 1 x 12 tray. chemical reaction for the titration is the following:
NaOH + HCl NaCl + H2O 4.1
►Is the solution in each of these wells acidic, neutral, or basic? Are the solutions at equilibrium?
To each well above, drops of 0.01 M NaOH were added. The numbers above the wells represents the drops of NaOH added. 0 1 2 3 4 5 6 7 8 9 10 11 4.2
NaOH + HCl NaCl + H2O 4.3
►Which solution is acidic? Which solution is neutral? Which solution is basic? + Where are the water molecules that are shown ►So, as [H ] increases, pH decreases, and - coming from? vice versa. What happens to pH as [OH ] Write the net ionic equation for this titration increases? reaction. DEFINITION: Is [H+] higher in the well with 1 drop of base pH = -log [H+] added, or 5 drops added? What about pH? Is [H+] higher in the well with 5 drop of base ►What is the pH of the 0.01 M HCl that we added, or 10 drops added? What about pH? 13 started with? ACID BASE EQUILIBRIA
TITRATION
4.4 A Thompson-Markow Universal Indicator
Titration of Strong Acid B with Strong Base
Titration of Weak Acid C with Strong Base
D Titration of Strong Acid with Weak Base
E Titration of Weak Acid with Weak Base
The following graphs show the titration curves of experiments C, D, and E.
►Using the graphs, give a molecular descrip- tion of each titration from pH 0 to 2, pH 4 to 6, and pH 8 to 10. Describe the molecular changes taking place in the wells at these pH regions.
C D 14 E 4.5 14 4.6 14 4.7
pH 7 pH 7 pH 7
0 7 14 0 7 14 0 7 14 Drops of 0.01 M NaOH added Drops of 0.01 M HCl added Drops of 0.01 M WB added to 0.05 M weak acid to 0.05 M weak base to 0.05 weak acid
14 ACID BASE EQUILIBRIA
TITRATION
4.8 A Thompson-Markow Universal Indicator
Titration of Strong Acid B with Strong Base
Titration of Weak Acid C with Strong Base
D Titration of Strong Acid with Weak Base
E Titration of Weak Acid with Weak Base
NaOH + HCl NaCl + H2O 4.10 B � 4.9 14
�
� pH 7 �
�
� 0 7 14 Drops of 0.01 M NaOH Added �
� ►Experiment B is a strong acid titrated with a strong base. Try to produce the titration curve for this experiment. The chemLog is � shown at the right to help you with this. ��
��
���
15 ACID BASE EQUILIBRIA
TITRATION
NaOH + HCl NaCl + H2O
A solution of HCl is shown in each of the drawings. ►Draw a sketch for each sample after the 4.11 4.12 NaOH shown has been added. The water mol- � � ecules are shown for you by the pink background. � � � � � �
►Are the drawings of HCl in solution (before NaOH is added) at equilibrium?
►Which of the sketches that you drew is at equilibrium?
►Is the pH of each of the solutions that you have drawn acidic, neutral, or basic?
A LIMITING REAGENT is a reactant that governs the maximum yield of product that is possible. For example, when 1 NaOH is added to a solu- tion with 10 HCl, the NaOH is the limiting reagent because it governs how many NaCl and water molecules can be formed in the reaction. 4.13 4.14
� � � � � � � � ►What is the limiting reagent for each of the sketches that you drew?
16 ACID BASE EQUILIBRIA
TITRATION
4.15 ►Write the chemical equations for the process shown in pictures 4.15 and 4.16.
►Is the pH of the final solution greater than 7, equal to 7, or less than 7?
►What is the pH of the water?
►Draw a picture of the final solution after one
more Ba(OH)2 is added. Again, the water is shown as a pink background.
4.16 � � �� � � �
►Is the pH of the solution you drew greater than 7, equal to 7, or less than 7? Explain.
►Is the solution you drew at equilibrium?
17 ACID-BASE EQUILIBRIA
BUFFERS
In experiments, we often need to control the [H+] Acidic buffers are solutions containing weak acids of the solution. This is done with buffers. A buff- and the salts of weak acids, eg. CH3COOH and + ered solution resists changes in [H ]. Buffers can CH3COONa. Basic buffers are solutions contain- have a predetermined pH, with acidic buffers at ing weak bases and the salts of weak bases, eg. any chosen value between pH 0-7, and basic buf- NH3 and NH4Cl. fers at any chosen value between pH 7-14. CH COOH H+ + CH COO– + – 3 3 CH3COOH H + CH3COO CH COONa Na+ + CH COO– + – 3 3 CH3COONa Na + CH3COO 5.1 5.2 Two dilute acetic acid solutions of acetic acid acetate ion at the same proton concentration. sodium
5.3 5.4
HCl NaOH
Add acid.
Add base.
When acid is added, the buffer When base is added, the buffer accepts protons. donates protons.
5.5 5.6
NaOH HCl
Add more acid.
Add more base.
The buffer capacity is an indication of the amount At which point in each of the additions above of acid or base that can be added before the buffer ► was the buffer capacity exceeded? loses its ability to resist the pH change. 18 ACID-BASE EQUILIBRIA
HENDERSON-HASSELBALCH EQUATION
How can we design a buffer with a specifi ed pH? Let’s calculate the pH of a buffer solution in which
Let’s examine the chemical reaction for solution 0.040 M CH3COONa (aq) and 0.080 M CH3COOH (aq) containing a general weak acid and its salt - a are mixed at 25 °C. buffer.
HA H+ + A– ►Identify the acid and its conjugate base in this NaA Na+ + A– mixture. Identify the ions from the salt in this mixture.
►Write the Ka expression for this reaction. 5.7
After rearranging the Ka expresstion to determine what [H+] is equal to, we determine that
+ [HA] [H ] = Ka x [A–]
We can now determine the pH. Write the pH expres- sion by taking the negative log of each side of the [H+] expression.
pH = – log [H+] = – log [HA] Ka x ( [A–] )
[HA] = – log Ka + – log – ►Write the equilibrium equation for acetic acid, ( [A ] ) and rearrange the Ka for this equation to give [H+]. [A–] = pKa + log ( [HA] )
Since acidic buffers are solutions of weak acids and their salt, HA loses only a tiny fraction of its protons. What are the dominant species (the most ►Use the Henderson-Hasselbalch equation to concentrated) in solution? They HA, becasue HA fi nd the pH. is a weak acid, and A–, because it comes from the complete dissociation of the salt NaA. Because they are so dominant in solution, we can approxi- mate [HA] and [A–] of a buffer by their initial con- centrations. The Henderson-Hasselbalch equation states that: [A–] pH ≈ pKa + log ([HA] ) initial
19 ACID-BASE EQUILIBRIA
POLYPROTIC ACIDS AND POLYBASIC BASES
Acids which can donate 2 or more protons are called polyprotic. Bases which can accept 2 or more protons are polybasic bases. Good exam- ples are phosphoric acid (H3PO4) and the phos- 3– + phate ion (PO4 ). Polyprotic acids release H ions in a step-wise manner:
+ H PO – 1. H3PO4 H + 2 4 – + HPO 2– 2. H2PO4 H + 4
– + PO 3– 3. H2PO4 H + 4
►Identify and write down the conjugate acids and bases in these dissociations.
►For H3PO4,, the pKa = 2.1. – For H2PO4 , the pKa = 7.2. 2– For HPO4 , the pKa = 13.
Calculate the Ka values.
►Write expressions for the Ka’s and Kb’s for all the acids and bases in these dissociation reactions.
►Why are the values of Ka’s and Kb’s very different from each other?
6.1 Thompson-Markow Universal Indicator
Titration of 1M H3PO4 with NaOH 1 2 3 4 5 6 7 8 9 10 11 12 Drops of NaOH added
6.2 14
7 pH
0 7 14
Drops of 0.01 M NaOH added to 0.05 M weak acid 20