19.5 Salts of Carboxylic Acids 743 QUANTITATIVE RELATIONSHIPS INVOLVING CARBOXYLIC ACIDS uppose you take two flasks, one containing pure This relationship is one form of the Henderson- water and the other a buffer solution main- Hasselbalch equation It is a useful relationship in tained at a pH of 7.0 If you add 0. 1 mol of acetic chemistry and biochemistry. one rarely needs to cal- acid to each one and the final volume in each flask is culate the ph of a solution-pH is more often mea- 1 L, how much acetic acid is present at equilibrium? sured than calculated. It is much more common that low much acetate ion? In other words, what is the one needs to know the degree of ionization of an extent of ionization of acetic acid in an unbuffered acid at a particular pH, and the Henderson-Hassel- edium and in a buffered one? balch equation gives that ratio The first case simply involves the ionization of a For the case at hand the solution is buffered at weak acid and is governed by the expression that de- pH=7.0. Therefore fines k for acetic acid: [CH3CO 18×10 IH ICH3CO2] =18×10 [CH3 CO,H [CH3 CO2 H] A very different situation exists in an aqueous solu- Since ionization of acetic acid gives one h for each tion maintained at ph =7.0 from the situation in CH3CO2, the concentrations of the two ions are pure water. We saw earlier that almost all the acetic equal, and setting each one equal to x gives: acid in a 0.1 M solution in pure water was nonion- ized. At pH 7.0, however, hardly any nonionized 18×10-5 acetic acid remains; it is almost completely converted to its carboxylate ion Solving for x gives the acetate ion concentration as: This difference in behavior for acetic acid in pure water versus water buffered at ph =7.0 has X=1.3×10 some important practical consequences. Biochemists usually do not talk about acetic acid (or lactic acid, or Thus when acetic acid is added to pure water, the ra- salicylic acid, etc. ) They talk about acetate (and lac- tio of acetate ion to acetic acid is tate, and salicylate). Why? It's because biochemists are concerned with carboxylic acids as they exist in di- CHo21=13×10=0013 lute aqueous solution at what is called biological pH 0.1 lological fluids are naturally buffered. The ph of blood, for example, is maintained at 7. 2, and at this Only 1.3% of the acetic acid has ionized. Most of it h carboxylic acids are almost entirely converted to (98.7%)remains unchanged Now think about what happens when the same their carboxylate anions. amount of acetic acid is added to water that is Daics An alternative form of the Henderson-Hassel- buffered at pH =7.0. Before doing the calculation, let us recognize that it is the [CH3 CO2 1/(CH3CO2H] CH3 CO2-I ratio in which we are interested and do a little alge- pH=pka +log (CH3 CO2H From this equation it e seen that whe K,=田cHo2 ICH3 CO2= [CH3,H then the second term is lo 1=0, and ph=pKa. This means that when the ph of a solution is equal to the pka of a weak acid, the con- centration of the acid and its conjugate base are [CHa CO equal. This is a relationship worth remembering [CH, CO,H [H+ Back Forward Main MenuToc Study Guide ToC Student o MHHE Website19.5 Salts of Carboxylic Acids 743 QUANTITATIVE RELATIONSHIPS INVOLVING CARBOXYLIC ACIDS S uppose you take two flasks, one containing pure water and the other a buffer solution main￾tained at a pH of 7.0. If you add 0.1 mol of acetic acid to each one and the final volume in each flask is 1 L, how much acetic acid is present at equilibrium? How much acetate ion? In other words, what is the extent of ionization of acetic acid in an unbuffered medium and in a buffered one? The first case simply involves the ionization of a weak acid and is governed by the expression that de- fines Ka for acetic acid: Ka   1.8  105 Since ionization of acetic acid gives one H for each CH3CO2 , the concentrations of the two ions are equal, and setting each one equal to x gives: Ka   1.8  105 Solving for x gives the acetate ion concentration as: x  1.3  103 Thus when acetic acid is added to pure water, the ra￾tio of acetate ion to acetic acid is   0.013 Only 1.3% of the acetic acid has ionized. Most of it (98.7%) remains unchanged. Now think about what happens when the same amount of acetic acid is added to water that is buffered at pH  7.0. Before doing the calculation, let us recognize that it is the [CH3CO2 ] ⁄[CH3CO2H] ratio in which we are interested and do a little alge￾braic manipulation. Since Ka  then  Ka [H] [CH3CO2 ] [CH3CO2H] [H][CH3CO2 ] [CH3CO2H] 1.3  103 0.1 [CH3CO2 ] [CH3CO2H] x2 0.1 x [H][CH3CO2 ] [CH3CO2H] This relationship is one form of the Henderson– Hasselbalch equation. It is a useful relationship in chemistry and biochemistry. One rarely needs to cal￾culate the pH of a solution—pH is more often mea￾sured than calculated. It is much more common that one needs to know the degree of ionization of an acid at a particular pH, and the Henderson–Hassel￾balch equation gives that ratio. For the case at hand, the solution is buffered at pH  7.0. Therefore,   180 A very different situation exists in an aqueous solu￾tion maintained at pH  7.0 from the situation in pure water. We saw earlier that almost all the acetic acid in a 0.1 M solution in pure water was nonion￾ized. At pH 7.0, however, hardly any nonionized acetic acid remains; it is almost completely converted to its carboxylate ion. This difference in behavior for acetic acid in pure water versus water buffered at pH  7.0 has some important practical consequences. Biochemists usually do not talk about acetic acid (or lactic acid, or salicylic acid, etc.). They talk about acetate (and lac￾tate, and salicylate). Why? It’s because biochemists are concerned with carboxylic acids as they exist in di￾lute aqueous solution at what is called biological pH. Biological fluids are naturally buffered. The pH of blood, for example, is maintained at 7.2, and at this pH carboxylic acids are almost entirely converted to their carboxylate anions. An alternative form of the Henderson–Hassel￾balch equation for acetic acid is pH  pKa log From this equation it can be seen that when [CH3CO2 ]  [CH3CO2H], then the second term is log 1  0, and pH  pKa. This means that when the pH of a solution is equal to the pKa of a weak acid, the con￾centration of the acid and its conjugate base are equal. This is a relationship worth remembering. [CH3CO2 ] [CH3CO2H] 1.8  105 107 [CH3CO2 ] [CH3CO2H] Back Forward Main Menu TOC Study Guide TOC Student OLC MHHE Website