# Chapter 15 Acid–Base Equilibria.

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Chapter 15 Acid–Base Equilibria

15.1 Solutions of Acids or Bases Containing a Common Ion

An application of Le Châtelier’s principle.

HCN(aq) + H2O(l) H3O+(aq) + CN-(aq)
Example HCN(aq) + H2O(l) H3O+(aq) + CN-(aq) Addition of NaCN will shift the equilibrium to the left because of the addition of CN-, which is already involved in the equilibrium reaction. A solution of HCN and NaCN is less acidic than a solution of HCN alone. Copyright © Cengage Learning. All rights reserved

Buffered Solution – resists a change in pH. They are weak acids or bases containing a common ion. After addition of strong acid or base, deal with stoichiometry first, then the equilibrium. Copyright © Cengage Learning. All rights reserved

Henderson–Hasselbalch Equation
For a particular buffering system (conjugate acid–base pair), all solutions that have the same ratio [A–] / [HA] will have the same pH. Copyright © Cengage Learning. All rights reserved

Exercise What is the pH of a buffer solution that is 0.45 M acetic acid (HC2H3O2) and 0.85 M sodium acetate (NaC2H3O2)? The Ka for acetic acid is 1.8 × 10–5. pH = 5.02 pH = –logKa + log([C2H3O2–] / [HC2H3O2]) = –log(1.8 × 10–5) + log(0.85 M / 0.45 M) = 5.02 Copyright © Cengage Learning. All rights reserved

Determined by the magnitudes of [HA] and [A–].
The amount of protons or hydroxide ions the buffer can absorb without a significant change in pH. Determined by the magnitudes of [HA] and [A–]. A buffer with large capacity contains large concentrations of the buffering components. Copyright © Cengage Learning. All rights reserved

Buffer capacity measures how well a solution resists changes in pH when acid or base is added. The greater the buffer capacity, the less the pH changes.  The amount of H+ or OH- that buffered solution can absorb without a significant change in pH

2) Magnitudes of [HA] and [A-]
 the capacity of a buffered soln. Ex : soln A : 5.00 M HOAc M NaOAc soln B : 0.05 M HOAc M NaOAc pH change when 0.01 mol of HCl(g) is added

3) [A-] / [HA] ratio  the pH of a buffered soln.

The pH Curve for the Titration of 50. 0 mL of 0. 200 M HNO3 with 0

The pH Curve for the Titration of 100. 0 mL of 0. 50 M NaOH with 1

Concept Check Consider a solution made by mixing 0.10 mol of HCN (Ka = 6.2 x 10–10) with mol NaOH in 1.0 L of aqueous solution. What are the major species immediately upon mixing (that is, before a reaction)? HCN, Na+, OH–, H2O Major Species: HCN, Na+, OH-, H2O. Copyright © Cengage Learning. All rights reserved

K = 1.8 x 109 Let’s Think About It… HC2H3O2(aq) + OH–  C2H3O2–(aq)

Steps Toward Solving for pH
HC2H3O2(aq) + H2O H3O+ + C2H3O2-(aq) Initial 0.170 M ~0 0.030 M Change –x +x Equilibrium 0.170 – x x x Ka = 1.8 x 10–5 pH = 3.99 Copyright © Cengage Learning. All rights reserved

Exercise Calculate the pH of a mL solution of M acetic acid (HC2H3O2), which has a Ka value of 1.8 x 10–5. pH = 2.87 pH = 2.87 Major Species: HC2H3O2, H2O Dominant reaction: HC2H3O2(aq) + H2O  H3O+(aq) + C2H3O2-(aq) The mL is not necessary to know for this problem. Copyright © Cengage Learning. All rights reserved

Concept Check Calculate the pH of a solution made by mixing mL of a M solution of acetic acid (HC2H3O2), which has a Ka value of 1.8 x 10–5, and 50.0 mL of a 0.10 M NaOH solution. pH = 4.74 pH = 4.74 Major Species: HC2H3O2, Na+, OH-, H2O Dominant reaction: HC2H3O2(aq) + OH-(aq)  H2O + C2H3O2-(aq) In general, the best acid (HC2H3O2 in this case) will tend to react with the best base (OH- will always be the best base, if present). Discuss how to calculate the K value for this reaction (note -- it is not a Ka or Kb expression). In this case, K = Ka/Kw = 1.8 x 109. We can assume the reaction goes to completion. In this problem, OH- is limiting. After the reaction takes place, the major species are: HC2H3O2, C2H3O2-, H2O Primary reaction is: HC2H3O2(aq) + H2O  H3O+(aq) + C2H3O2-(aq) When students solve this, make sure they include the initial concentration of the conjugate base (they tend to forget and call it zero). Copyright © Cengage Learning. All rights reserved

Concept Check Calculate the pH of a solution at the equivalence point when mL of a M solution of acetic acid (HC2H3O2), which has a Ka value of 1.8 x 10–5, is titrated with a 0.10 M NaOH solution. pH = 8.72 pH = 8.72 Major Species: HC2H3O2, Na+, OH-, H2O Dominant reaction: HC2H3O2(aq) + OH-(aq)  H2O + C2H3O2-(aq) In general, the best acid (HC2H3O2 in this case) will tend to react with the best base (OH- will always be the best base, if present). At the equivalence point, there are the same number of moles of HC2H3O2(aq) + OH-(aq) to exactly react due to the 1:1 ratio in the balanced equation. This turns out to be mol of both the acid and base in this case. This means that mL of NaOH had to be added. After the reaction takes place, the major species are: C2H3O2-, H2O, Na+ Primary reaction is: C2H3O2-(aq) + H2O  OH-(aq) + HC2H3O2(aq) Now use an ICE chart to determine the equilibrium concentrations and finally the pH. Take note that the initial concentration of C2H3O2-(aq) is M (0.010 mol / L). Copyright © Cengage Learning. All rights reserved

The pH Curve for the Titration of 50. 0 mL of 0. 100 M HC2H3O2 with 0

The pH Curves for the Titrations of 50. 0-mL Samples of 0
The pH Curves for the Titrations of 50.0-mL Samples of 0.10 M Acids with Various Ka Values with 0.10 M NaOH Copyright © Cengage Learning. All rights reserved

The pH Curve for the Titration of 100. 0mL of 0. 050 M NH3 with 0

Kjeldahl Nitrogen Analysis : developed in 1883, the analysis remains one of the most widely used methods for determining nitrogen in organic substances such as protein, cereal, and flour. P.228

Melamine contaminated milk

Marks the end point of a titration by changing color.

The Acid and Base Forms of the Indicator Phenolphthalein

The Methyl Orange Indicator is Yellow in Basic Solution and Red in Acidic Solution

HIn(aq)  H+(aq) + In-(aq) pKHIn
1) Usually a weak organic acid or base that has distinctly different colors in its nonionized & ionized forms. HIn(aq)  H+(aq) + In-(aq) pKHIn nonionized ionized form form

1) 2) 3)

Useful pH Ranges for Several Common Indicators