ACIDS AND BASES II

IB CHEMISTRY GR.12

Topic 18

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TOPICS

18.1 Calculations Involving Acids and Bases 18.2 Buffer Solutions 18.3 Salt Hydrolysis 18.4 Acid-Base Titrations 18.5 Indicators

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IB STANDARDS

18.1.1 State the expression for the ionic product constant of water (Kw).

18.1.2 Deduce [H+(aq)] and [OH-(aq)], for water at different temperatures given Kw values.

18.1.3 Solve problems involving [H+(aq)], [OH-(aq)], pH and pOH.

18.1.4 State the equation for the reaction for any weak acid or weak base with water, and hence deduce the expressions for Ka and Kb. Chem

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IB STANDARDS

• 18.1.5 Solve problems involving solutions of weak acids and bases using the expressions:

Ka x Kb = Kw

pKa + pKb = pKw

pH + pOH = pKw

• 18.1.6 Identify the relative strengths of acids and bases using values of Ka, Kb, pKa and pKb.

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H2O (l) H+ (aq) + OH- (aq)

The Ion Product of Water

Kc =[H+][OH-]

[H2O][H2O] = constant

Kc[H2O] = Kw = [H+][OH-]

The ion-product constant (Kw) is the product of the molar concentrations of H+ and OH- ions at a particular temperature.

At 250CKw = [H+][OH-] = 1.0 x 10-14

[H+] = [OH-]

[H+] > [OH-]

[H+] < [OH-]

Solution Is

neutral

acidic

basic

KW IS TEMPERATURE DEPENDENT

Temp oC Kw [H+] in pure water

pH

0 1.5 x 10-15 3.9 x 10-6 7.47

20 6.8 x 10-15 8.2 x10-6 7.08

25 1.0 x 10-14 1.0 x10-7 7.00

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pH = - log [H+], [H+] = 10-pH

pOH = -log [OH-], [OH-] = 10-pOH

[H+] = [OH-]

[H+] > [OH-]

[H+] < [OH-]

Solution isneutral

acidic

basic

[H+] = 1 x 10-7

[H+] > 1 x 10-7

[H+] < 1 x 10-7

pH = 7

pH < 7

pH > 7

At 250C

pH [H+]

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[H+][OH-] = Kw = 1.0 x 10-14

-log [H+] – log [OH-] = 14.00

pH + pOH = 14.00

At 250C

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Example 1: The pH of rainwater collected in a certain region of the northeastern United States on a particular day was 4.82. What is the H+ ion concentration of the rainwater?

pH = -log [H+]

[H+] = 10-pH = 10-4.82 = 1.5 x 10-5 M

Example 2: The OH- ion concentration of a blood sample is 2.5 x 10-7 M. What is the pH of the blood?

pH + pOH = 14.00

pOH = -log [OH-] = -log (2.5 x 10-7) = 6.60

pH = 14.00 – pOH = 14.00 – 6.60 = 7.40

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HA (aq) + H2O (l) H3O+ (aq) + A- (aq)

Weak Acids (HA) and Acid Ionization Constants

HA (aq) H+ (aq) + A- (aq)

Ka =[H+][A-][HA]

Ka is the acid ionization constant

Ka

weak acidstrength

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NH3 (aq) + H2O (l) NH4+ (aq) + OH- (aq)

Weak Bases and Base Ionization Constants

Kb =[NH4

+][OH-][NH3]

Kb is the base ionization constant

Kb

weak basestrength

Solve weak base problems like weak acids except solve for [OH-] instead of [H+].

IONIZATION CONSTANTS OF CONJUGATE ACID-BASE PAIRS

Consider a weak acid (HA) and its conjugate base (A-) in water:HA (aq) H+(aq) + A-(aq)

Ka = [H+][A-] / [HA]

A-(aq) + H2O(l) HA(aq) + OH-(aq) Kb = [HA][OH-] /[A-]

Ka x Kb = [H+][OH-] = Kw

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IB STANDARDS

18.2.1 Describe the composition of a buffer solution and explain its action.

18.2.2 Solve problems involving the composition and pH of a specified buffer system.

18.3 Deduce whether salts form acidic, alkaline or neutral aqueous solutions.

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BUFFER SOLUTIONS

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A buffer solution is a solution of:

1. A weak acid or a weak base and2. The salt of the weak acid or weak base

Both must be present!

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The common ion effect is the shift in equilibrium caused by the addition of a compound having an ion in common with the dissolved substance.

The presence of a common ion suppresses the ionization of a weak acid or a weak base.

Example: A buffer consisting of CH3COONa (salt) and CH3COOH (weak acid).

CH3COONa (s) Na+ (aq) + CH3COO- (aq)

CH3COOH (aq) H+ (aq) + CH3COO- (aq)

common ion

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• A buffer solution has the ability to resist changes in pH upon the addition of small amounts of either acid or base.

Add strong acid

H+ (aq) + CH3COO- (aq) CH3COOH (aq)

Add strong base

OH- (aq) + CH3COOH (aq) CH3COO- (aq) + H2O (l)

How buffers work

Consider an equal molar mixture of CH3COOH and CH3COONa

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Consider mixture of salt NaA and weak acid HA.

HA (aq) H+ (aq) + A- (aq)

NaA (s) Na+ (aq) + A- (aq)Ka =

[H+][A-][HA]

[H+] =Ka [HA]

[A-]

-log [H+] = -log Ka - log[HA]

[A-]

-log [H+] = -log Ka + log [A-][HA]

pH = pKa + log [A-][HA]

pKa = -log Ka

HENDERSON-HASSELBALCH EQUATION

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pH = pKa + log[conjugate base]

[acid]

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pH of buffers can be determined from Henderson-Hasselbalch Equation.

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Exercise: Which of the following are buffer systems? (a) KF/HF (b) KBr/HBr, (c) Na2CO3/NaHCO3

(a) KF is a weak acid and F- is its conjugate basebuffer solution

(b) HBr is a strong acidnot a buffer solution

(c) CO32- is a weak base and HCO3

- is its conjugate acidbuffer solution

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20= 9.20

Problem: Calculate the pH of the 0.30 M NH3/0.36 M NH4Cl buffer system. What is the pH after the addition of 20.0 mL of 0.050 M NaOH to 80.0 mL of the buffer solution?

NH4+ (aq) H+ (aq) + NH3 (aq)

pH = pKa + log[NH3][NH4

+]pKa = 9.25 pH = 9.25 + log

[0.30][0.36]

= 9.17

NH4+ (aq) + OH- (aq) H2O (l) + NH3 (aq)

start (moles)

end (moles)

0.029 0.001 0.024

0.028 0.0 0.025

pH = 9.25 + log[0.25][0.28]

[NH4+] =

0.0280.10

final volume = 80.0 mL + 20.0 mL = 100 mL

[NH3] = 0.0250.10

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Acid-Base Properties of Salts

Neutral Solutions:

Salts containing an alkali metal or alkaline earth metal ion (except Be2+) and the conjugate base of a strong acid (e.g. Cl-, Br-, and NO3

-).

NaCl (s) Na+ (aq) + Cl- (aq)H2O

Basic Solutions:

Salts derived from a strong base and a weak acid.

NaCH3COOH (s) Na+ (aq) + CH3COO- (aq)H2O

CH3COO- (aq) + H2O (l) CH3COOH (aq) + OH- (aq)

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Acid-Base Properties of Salts

Acid Solutions:

Salts derived from a strong acid and a weak base.

NH4Cl (s) NH4+ (aq) + Cl- (aq)

H2O

NH4+ (aq) NH3 (aq) + H+ (aq)

Salts with small, highly charged metal cations (e.g. Al3+, Cr3+, and Be2+) and the conjugate base of a strong acid.

Al(H2O)6 (aq) Al(OH)(H2O)5 (aq) + H+ (aq)3+ 2+

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Acid-Base Properties of Salts

Solutions in which both the cation and the anion hydrolyze:

• Kb for the anion > Ka for the cation, solution will be basic

• Kb for the anion < Ka for the cation, solution will be acidic

• Kb for the anion Ka for the cation, solution will be neutral

IB STANDARDS

18.4.1Sketch the general shapes of graphs of pH and volume for titrations of strong and weak acids and bases and explain their important features.

18.5.1 Describe qualitatively the action of acid-base indicator.

18.5.2 State and explain how the pH range of an acid-base indicator relates to its pKa value.

18.5.3 Identify an appropriate indicator for a titration given the equivalence point of titration and the pH range of the indicator. 24

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TitrationsIn a titration a solution of accurately known concentration is added gradually to another solution of unknown concentration until the chemical reaction between the two solutions is complete.

Equivalence point – the point at which the reaction is complete

Indicator – substance that changes color at (or near) the equivalence point

Slowly add baseto unknown acid

UNTIL

The indicatorchanges color

(pink)

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Strong Acid-Strong Base Titrations

NaOH (aq) + HCl (aq) H2O (l) + NaCl (aq)

OH- (aq) + H+ (aq) H2O (l)

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Weak Acid-Strong Base Titrations

CH3COOH (aq) + NaOH (aq) CH3COONa (aq) + H2O (l)

CH3COOH (aq) + OH- (aq) CH3COO- (aq) + H2O (l)

CH3COO- (aq) + H2O (l) OH- (aq) + CH3COOH (aq)

At equivalence point (pH > 7):

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Strong Acid-Weak Base Titrations

HCl (aq) + NH3 (aq) NH4Cl (aq)

NH4+ (aq) + H2O (l) NH3 (aq) + H+ (aq)

At equivalence point (pH < 7):

H+ (aq) + NH3 (aq) NH4Cl (aq)

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Acid-Base IndicatorsHIn (aq) H+ (aq) + In- (aq)

10[HIn][In-] Color of acid (HIn) predominates

10[HIn][In-] Color of conjugate base (In-) predominates

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The titration curve of a strong acid with a strong base.C

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Which indicator(s) would you use for a titration of HNO2 with KOH ?

Weak acid titrated with strong base.

At equivalence point, will have conjugate base of weak acid.

At equivalence point, pH > 7

Use cresol red or phenolphthalein

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Acid-Base Properties of Salts

Acid Solutions:

Salts derived from a strong acid and a weak base.

NH4Cl (s) NH4+ (aq) + Cl- (aq)

H2O

NH4+ (aq) NH3 (aq) + H+ (aq)

Salts with small, highly charged metal cations (e.g. Al3+, Cr3+, and Be2+) and the conjugate base of a strong acid.

Al(H2O)6 (aq) Al(OH)(H2O)5 (aq) + H+ (aq)3+ 2+

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Acid-Base Properties of Salts

Solutions in which both the cation and the anion hydrolyze:

• Kb for the anion > Ka for the cation, solution will be basic

• Kb for the anion < Ka for the cation, solution will be acidic

• Kb for the anion Ka for the cation, solution will be neutral

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Chemistry In Action: Antacids and the Stomach pH Balance

NaHCO3 (aq) + HCl (aq)

NaCl (aq) + H2O (l) + CO2 (g)

Mg(OH)2 (s) + 2HCl (aq)

MgCl2 (aq) + 2H2O (l)

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