chapter outline€¢ 15.7 the common-ion effect • 15.8 ph buffers • 15.9 ph indicators and...

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Aqueous Equilibria:

Chemistry of the Water World

Chapter Outline

2

• 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

React with certain metals to produce hydrogen gas.

Have a bitter taste.

Feel slippery. Many soaps contain bases.

Have a sour taste. Vinegar owes its taste

to acetic acid. Citrus fruits contain citric acid.

Bases

React with carbonates and bicarbonates to produce

carbon dioxide gas

Nomenclature Review –

Ch 4, Section 4.2

PO43-, HPO4

2-, H2PO4-

SO42-, HSO4

-

SO32-, HSO3

-

CO32-, HCO3

-

NO3-, NO2

-

S2-, HS-

C2H3O2- (CH3COO-)

binary acids, oxoacids

You are only responsible for nomenclature taught in the lab. These ions

are part of many different acids and you need to know them!

H3PO4

H2SO3

H2SO4

H2CO3

HNO3, HNO2

H2S

HC2H3O2

HCl, HClO4

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Strong and Weak Acids

HNO3(aq) + H2O(ℓ) → NO3-(aq) + H3O

+(aq)

(H+ donor) (H+ acceptor)

Strong Acid: Completely ionized

Weak Acid: Partially ionized

HNO2(aq) + H2O(ℓ) ⇌ NO2-(aq) + H3O

+(aq)

(H+ donor) (H+ acceptor)

A Brønsted acid is a proton donor

A Brønsted base is a proton acceptor

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Hydronium Ion

Conjugate Acid-Base Pairs

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Weak Acids

reordered

Kc =CH3COO

−(aq) [H3O+(aq)]

CH3COOH(aq) [H2O(l)]

stronger

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Strong and Weak Bases

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

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• 15.8 pH Buffers



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

O

H

H + O

H

H O

H

H H O H - +

[ ] +

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

H2O + H2O H3O+ + OH-

acid conjugate

base

base conjugate

acid

autoionization of water

equivalent

expressions

H+

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Chapter Outline

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• 15.8 pH Buffers



pH and the Autoionization of Water

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

Using the RICE table -

[H3O+][OH-]

Kc = [H2O]2

2 H2O(l) = H3O+ + OH-

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pH - A Measure of Acidity

pH = -log [H+]

[H+] = [OH-]

[H+] > [OH-]

[H+] < [OH-]

Solution

neutral

acidic

basic

[H+] = 1 x 10-7 [H+] > 1 x 10-7

[H+] < 1 x 10-7

pH = 7 pH < 7

pH > 7

pH [H+]

The pH Scale

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pH, pOH, and K

pOH = - log[OH-]

pOH is defined the same way as pH -

“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]

pKb = - log[Kb]

• The smaller the Ka, the weaker the acid

• The weaker the acid, the larger the pKa

• The same concepts apply for weak bases

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Useful Equation for Acid-Base Calculations

1. Starting with Kw

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

2. Taking the negative log of both sides -

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

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

14 = pH + pOH

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Chapter Outline

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• 15.8 pH Buffers



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

H2SO4 sulfuric

The monster in the movie

“Alien’” had blood that contained

“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)

Ka = [H3O

+][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

Percent Ionization = [H

+]equil

[𝐻𝐴]initial X 100%

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Weak Bases

NH3

ammonia

C6H5NH2

aniline

NH(CH3)2

dimethylamine

Weak bases frequently contain nitrogen because

the lone pair makes a good proton acceptor

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

Kb = [BH+][OH-]

[B]

General Weak Base

Equilibrium Equation and Kb

Ordinary bleach

contains the

weak base ClO-

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Relationship Between Ka and Kb

Kb = [HA][OH-]

[A-] Ka =

[H3O+][A-]

[HA]

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Chapter Outline

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• 15.8 pH Buffers



Polyprotic Acids

Two or more ionizable protons

H H+

H+ H

H

H+

Ka1 = 7.11 x 10-3

Ka2 = 6.32 x 10-8

Ka3 = 4.5 x 10-13

Ka1 > Ka2 > Ka3

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Chapter Outline

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• 15.8 pH Buffers



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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)

HCl(aq) + H2O(l) H3O+(aq) + Cl-(aq)

base conj. acid

Infinitely

strong

Infinitely

weak

acid conj. base

pH of Salt Solutions

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

HClO(aq) + H2O(l) = H3O+(aq) + ClO-(aq)

ClO-(aq) + H2O(l) = OH-(aq) + HClO(aq)

weak

acid

conj.

base

Ka (HClO) = 2.9 x 10-8 Kb = Kw

Ka

Kb = 1.0 x 10-14

2.9 x 10-8

conj. base pH > 7

Kb (ClO-) = 3.4 x 10-7

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

pH of Salt Solutions

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

NH4+(aq) + H2O(l) = H3O

+(aq) + NH3(aq)

weak

base

conj.

acid

Kb (NH3) = 1.8 x 10-5 Ka = Kw

Kb

Ka = 1.0 x 10-14

1.8 x 10-5

conj. acid pH < 7

Ka (NH4+) = 5.6 x 10-10

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Calculating the pH of Solutions of Weak Acids and Bases:

Use the RICE Table as Before

ClO-(aq) + H2O(l) = OH-(aq) + HClO(aq)

0.100 M - x

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

x x

Kb = 3.4 x 10-7 = x2

0.100 M - x

Since Ka is < 10-5, assume

that x << 0.100 M 3.4 x 10-7 = x2

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

Ka = 2.9 x 10-8

Kb = Kw

Ka

Kb = 3.4 x 10-7

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

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• 15.8 pH Buffers



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)

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

common

ion

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)

Ka = [H3O

+][A-]

[HA]

1. Solution using the RICE table

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2. Solution using the Henderson-Hasselbach equation

Chapter Outline

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• 15.8 pH Buffers



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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.

[H2CO3] = 1.3 x 10-5 M and [HCO3-] = 1.0 x 10-4 M

1.0 L samples of river water and pure water, and then add 10.0

mL of 1.0 x 10-3 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

e.g. CH3COO-/CH3COOH

Weak Base and its Salt

e.g. NH4+/NH3

Add H+

Add OH-

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Chapter Outline

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• 15.8 pH Buffers



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

20.0 mL

0.100 M

each

analyte titrant

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

(b) pH eq. pt. = 7

pH = 7

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Acid-Base Titrations (p. 692):

Weak Acid

NaOH + CH3COOH NaCH3COO + H2O

20.0 mL

0.100 M

each

analyte titrant

(a) Initial pH = use RICE table

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

(c) pH midpoint = Henderson-

Hasselbach eqn.

Titration Calculations:

Concentration of the Unknown

aA + bB products

acid base

MA, VA MB, VB

known unknown

Remember: M x V = moles

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Chapter Outline

69








• 15.8 pH Buffers



Calculating the concentration of

sparingly soluble salts in solution

(compounds that violated the

solubility rules do ionize to a very

limited extent)

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

Mg(OH)2(s) = Mg2+(aq) + 2 OH-(aq) Ksp = 5.6 x 10-12

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General Equation for Calculating the Molar Solubility

AmBn(s) = m An+(aq) + n Bm-(aq)

Common Ion Effect

BaSO4(s) = Ba2+(aq) + SO42-(aq) Ksp = 9.1 x 10-11

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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

chapter outline€¢ 15.7 the common-ion effect • 15.8 ph buffers • 15.9 ph indicators and...

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