Chapter Outline

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Aqueous Equilibria: Chemistry of the Water World

Chapter Outline

• 15.1 Acids and Bases: The BrØnsted–Lowry Model • 15.2 Acid Strength and Molecular Structure • 15.3 pH and the Autoionization of Water

• 15.4 Calculations Involving pH, Ka, and Kb • 15.5 Polyprotic Acids • 15.6 pH of Salt Solutions • 15.7 The Common-Ion Effect • 15.8 pH Buffers • 15.9 pH Indicators and Acid–Base Titrations • 15.10 Solubility Equilibria

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Acids

Have a sour taste. Vinegar owes its taste to acetic acid. Citrus fruits contain citric acid. React with certain metals to produce hydrogen gas. React with carbonates and bicarbonates to produce carbon dioxide gas Bases Have a bitter taste. Feel slippery. Many soaps contain bases.

Nomenclature Review – Ch 4, Section 4.2

You are only responsible for nomenclature taught in the lab. These ions are part of many different acids and you need to know them!

3- 2- - PO4 , HPO4 , H2PO4 H3PO4 2- - SO4 , HSO4 H2SO4 2- - SO3 , HSO3 H2SO3 2- - CO3 , HCO3 H2CO3 - - HNO HNO NO3 , NO2 3, 2 2- - S , HS H2S - - C2H3O2 (CH3COO ) HC2H3O2

binary acids, oxoacids HCl, HClO4

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Strong and Weak Acids

A Brønsted acid is a proton donor A Brønsted base is a proton acceptor

Strong Acid: Completely ionized

- + HNO3(aq) + H2O(ℓ) → NO3 (aq) + H3O (aq) (H+ donor) (H+ acceptor)

Weak Acid: Partially ionized

- + HNO2(aq) + H2O(ℓ) ⇌ NO2 (aq) + H3O (aq) (H+ donor) (H+ acceptor)

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Hydronium Ion

Conjugate Acid-Base Pairs

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Weak Acids

reordered

stronger − + CH3COO (aq) [H3O (aq)] Kc = CH3COOH(aq) [H2O(l)]

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Strong and Weak Bases

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Weak Bases

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Relative Strengths of Acids/Bases

Leveling Effect:

+ • H3O is the strongest H+ donor that can exist in water. • Strong acids all have the same strength in water; they are completely converted + into H3O ions.

Relative Strengths of Acids/Bases

Leveling Effect Bases: - + OH is the strongest H acceptor that can exist in H2O

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Chapter Outline

• 15.1 Acids and Bases: The BrØnsted–Lowry Model • 15.2 Acid Strength and Molecular Structure • 15.3 pH and the Autoionization of Water

Acid Strength and Molecular Structure

H2SO4 is a stronger acid because –

1. The -2 charge is delocalized over 4 oxygen atoms compared to three 2. the larger number of oxygens in

H2SO4 creates a greater electronegativity effect and consequent weakening of the O-H bond.

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The Acid-Base Properties of Water Water is amphoteric - which means that it can behave either as an acid or a base

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

autoionization of water H+ + - H O + H O [ H O H] + H O H H H

conjugate base acid + - H2O + H2O H3O + OH equivalent conjugate expressions acid base

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Chapter Outline

• 15.1 Acids and Bases: The BrØnsted–Lowry Model • 15.2 Acid Strength and Molecular Structure • 15.3 pH and the Autoionization of Water

pH and the Autoionization of Water

[H O+][OH-] 2 H O(l) = H O+ + OH- 3 2 3 Kc = 2 [H2O]

+ - What is the concentration of H3O and OH in pure water? Using the RICE table -

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pH - A Measure of Acidity

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

Solution neutral [H+] = [OH-] [H+] = 1 x 10-7 pH = 7 acidic [H+] > [OH-] [H+] > 1 x 10-7 pH < 7 basic [H+] < [OH-] [H+] < 1 x 10-7 pH > 7

The pH Scale

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pH, pOH, and K

pOH is defined the same way as pH -

pOH = - log[OH-]

“p-functions” are very common in chemistry, i.e. the negative log of any physical constant is calculated the same way.

Since Ka and Kb values for weak acids and bases tend to be very small, it’s convenient to take the negative log of these values as well

pKa = - log[Ka] • The smaller the Ka, the weaker the acid • The weaker the acid, the larger the pKa • The same concepts apply for weak bases pKb = - log[Kb]

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Useful Equation for Acid-Base Calculations

1. Starting with Kw

+ - -14 Kw = [H3O ][OH ] = 1.00 x 10

2. Taking the negative log of both sides -

+ - - log Kw = - log [H3O ][OH ]

-14 + - - log(1.00 x 10 ) = - log [H3O ] - log[OH ]

14 = pH + pOH

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Chapter Outline

• 15.1 Acids and Bases: The BrØnsted–Lowry Model • 15.2 Acid Strength and Molecular Structure • 15.3 pH and the Autoionization of Water

Weak Acids

• Most acids are weak. How do you know if an acid is weak? • Because it’s not one of the 6 strong ones you’ve memorized! HCl hydrochloric HBr hydrobromic HI hydroiodic

HNO3 nitric HClO4 perchloric The monster in the movie “Alien’” had blood that contained H2SO4 sulfuric “molecular acid” and ate through six decks of the spaceship!

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General Weak Acid

Equilibrium Equation and Ka

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

[H O+][A-] K = 3 a [HA]

If you measure the pH of a solution containing a weak acid, you can calculate the equilibrium constant

Calculating Ka for a weak acid when the pH is known, e.g 0.100 M, pH = 2.20

+ - HA(aq) + H2O(l) = H3O (aq) + A (aq) 0.100 M

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Percent Ionization

+ [H ]equil Percent Ionization = X 100% [퐻퐴]initial

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Weak Bases

Weak bases frequently contain nitrogen because the lone pair makes a good proton acceptor

NH3 ammonia

NH(CH3)2 C6H5NH2 dimethylamine aniline

General Weak Base

Equilibrium Equation and Kb

+ - B(aq) + H2O(l) = BH (aq) + OH (aq)

[BH+][OH-] K = b [B]

Ordinary bleach contains the weak base ClO-

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Relationship Between Ka and Kb

+ - - [H3O ][A ] [HA][OH ] K = K = a [HA] b [A-]

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Chapter Outline

• 15.1 Acids and Bases: The BrØnsted–Lowry Model • 15.2 Acid Strength and Molecular Structure • 15.3 pH and the Autoionization of Water

Polyprotic Acids

Two or more ionizable protons Ka1 > Ka2 > Ka3

H H+ -3 Ka1 = 7.11 x 10

H H+ -8 Ka2 = 6.32 x 10

H+ -13 Ka3 = 4.5 x 10 H

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Chapter Outline

• 15.1 Acids and Bases: The BrØnsted–Lowry Model • 15.2 Acid Strength and Molecular Structure • 15.3 pH and the Autoionization of Water

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pH of Salt Solutions

1. Neutral Salts (pH = 7) are from strong electrolytes (100% ionization) (a)Ionic Compounds:

NaCl(aq)  Na+(aq) + Cl-(aq)

base conj. acid + - HCl(aq) + H2O(l)  H3O (aq) + Cl (aq) acid conj. base

Infinitely Infinitely strong weak

pH of Salt Solutions

2. Basic Salts (pH > 7) are conjugate bases of weak acids

+ - HClO(aq) + H2O(l) = H3O (aq) + ClO (aq) weak conj. acid base

- - ClO (aq) + H2O(l) = OH (aq) + HClO(aq) conj. base pH > 7

K 1.0 x 10-14 K (HClO) = 2.9 x 10-8 K = w K = a b b -8 Ka 2.9 x 10 - -7 Kb (ClO ) = 3.4 x 10

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pH of Salt Solutions

3. Acidic Salts (pH < 7) are conjugate acids of weak bases

- + NH3(aq) + H2O(l) = OH (aq) + NH4 (aq) weak conj. base acid

+ + NH4 (aq) + H2O(l) = H3O (aq) + NH3(aq) conj. acid pH < 7

K 1.0 x 10-14 K (NH ) = 1.8 x 10-5 K = w K = b 3 a a -5 Kb 1.8 x 10 + -10 Ka (NH4 ) = 5.6 x 10

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Calculating the pH of Solutions of Weak Acids and Bases: Use the RICE Table as Before

1. Calculating the pH of a Solution of a Basic Salt

- - -8 ClO (aq) + H2O(l) = OH (aq) + HClO(aq) Ka = 2.9 x 10 0.100 M - x x x K = Kw b K 2 a -7 x Kb = 3.4 x 10 = -7 0.100 M - x Kb = 3.4 x 10

x2 -5 3.4 x 10-7 = Since Ka is < 10 , assume 0.100 M that x << 0.100 M

x = (3.4 x 10-7)(0.100)

x = [OH-] = 1.9 x 10-4 M Assumption OK

pOH = - log (1.9 x 10-4) = 3.7 pH = 14 - 3.7 = 10.3

Calculating the pH of Solutions of Weak Acids and Bases: Use the RICE Table as Before

2. Calculating the pH of a Solution of an Acidic Salt (Ex. 15.8)

What is the pH of a 0.25 M solution of NH4Cl?

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Chapter Outline

• 15.1 Acids and Bases: The BrØnsted–Lowry Model • 15.2 Acid Strength and Molecular Structure • 15.3 pH and the Autoionization of Water

The Common Ion Effect

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

+ - CH3COONa (s) Na (aq) + CH3COO (aq) common + - ion CH3COOH (aq) H (aq) + CH3COO (aq)

The common ion effect can be used to produce a BUFFER SOLUTION = a solution of a weak acid or base and it's

conjugate, e.g. CH3COOH and CH3COONa

By controlling the ratio of weak acid/base to it's conjugate, we can shift the equilibrium to whatever [H+] and therefore pH we want.

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The Henderson-Hasselbach Equation

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

+ - [H3O ][A ] Ka = [HA]

1. Solution using the RICE table

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2. Solution using the Henderson-Hasselbach equation

Chapter Outline

• 15.1 Acids and Bases: The BrØnsted–Lowry Model • 15.2 Acid Strength and Molecular Structure • 15.3 pH and the Autoionization of Water

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pH Buffers

Calculate the response of buffers to an influx of acid or base as compared to the same amount in pure water, p.686.

-5 - -4 [H2CO3] = 1.3 x 10 M and [HCO3 ] = 1.0 x 10 M 1.0 L samples of river water and pure water, and then add 10.0 -3 mL of 1.0 x 10 HNO3 (a) Calculate the pH change in pure water

pH Buffers

(b) Calculate the pH change in the buffer

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A Buffer in Action

Weak Acid and its Salt Weak Base and its Salt - + e.g. CH3COO /CH3COOH e.g. NH4 /NH3

Add H+

Add OH-

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Chapter Outline

• 15.1 Acids and Bases: The BrØnsted–Lowry Model • 15.2 Acid Strength and Molecular Structure • 15.3 pH and the Autoionization of Water

pH Indicators

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Acid-Base Titrations In a titration a solution of accurately known concentration (titrant) is added gradually added to another solution of unknown concentration (analyte) 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 base to unknown acid UNTIL The indicator changes color (pink)

Acid-Base Titrations (p. 692): Strong Acid

NaOH + HCl  NaCl + H2O titrant analyte

20.0 mL pH = 7 0.100 M each

(a) Initial pH = - log (0.100 M) = 1.00

(b) pH eq. pt. = 7

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Acid-Base Titrations (p. 692): Weak Acid

NaOH + CH3COOH  NaCH3COO + H2O titrant analyte

20.0 mL 0.100 M each

(a) Initial pH = use RICE table

(b) pH eq. pt. < 7, RICE table again

Titration Calculations: Concentration of the Unknown

aA + bB  products

acid base

MA, VA MB, VB Remember: M x V = moles unknown known

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Chapter Outline

• 15.1 Acids and Bases: The BrØnsted–Lowry Model • 15.2 Acid Strength and Molecular Structure • 15.3 pH and the Autoionization of Water

Solubility Equilibria

Calculating the concentration of sparingly soluble salts in solution (compounds that violated the solubility rules do ionize to a very limited extent)

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Solubility Equilibria

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Solubility Equilibria and the Solubility Product Ksp

Another equilibrium constant that allows calculation of the amount of a compound that will dissolve in water.

AgCl(s)

Ca3(PO4)2(s)

Calculating the Molar Solubility from Ksp

2+ - -12 Mg(OH)2(s) = Mg (aq) + 2 OH (aq) Ksp = 5.6 x 10

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General Equation for Calculating the Molar Solubility

n+ m- AmBn(s) = m A (aq) + n B (aq)

Common Ion Effect

2+ 2- -11 BaSO4(s) = Ba (aq) + SO4 (aq) Ksp = 9.1 x 10

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Need to include the hydrolysis of the conjugate acid or base

Ksp and Q

Q > Ksp forms too much product so it precipitates

Q = Ksp at equilibrium

Q < Ksp not enough products so no precipitate forms