1. Chapter 14 Aqueous Equilibria: Acids and
Chemistry: McMurry and Fay, 6th Edition
Chapter 14: Aqueous Equilibria: Acids and Bases
Chapter 14 Aqueous Equilibria: Acids and
4/25/ :04:27 PM
The Brønsted-Lowry Theory
Arrhenius Acid: A substance that dissociates in water to produce hydrogen ions, H+ or increases the concentration of hydrogen ion
H+(aq) + A(aq)
HA(aq)
Arrhenius Base: A substance that dissociates in water to produce hydroxide ions, OH-1 increases the concentration of hydroxide ions.
Ammonia, NH3, can’t be explained properly using Arrhenius theory.
M+(aq) + OH(aq)
MOH(aq)
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2. Acid Dissociation What happens when an acid dissolves in water?
HCl (aq) + H2O (l) → Cl–(aq) + H3O+ (aq)
Water acts as a Bronsted-Lowry base and takes a proton (H+) from the acid.
As a result, the conjugate base of the acid and a hydronium ion are formed.

3. Acid-Base Concepts: The Brønsted-Lowry Theory
Chemistry: McMurry and Fay, 6th Edition
Chapter 14: Aqueous Equilibria: Acids and Bases
4/25/ :04:27 PM
Acid-Base Concepts: The Brønsted-Lowry Theory
Brønsted-Lowry Acid: A substance that can transfer hydrogen ions, H+. In other words, a proton donor
Brønsted-Lowry Base: A substance that can accept hydrogen ions,
H+. In other words, a proton acceptor
Conjugate Acid-Base Pairs: Chemical species whose formulas differ only by one hydrogen ion, H+
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4. 14.3 Hydrated Protons and Hydronium Ions
Chemistry: McMurry and Fay, 6th Edition
Chapter 14: Aqueous Equilibria: Acids and Bases
4/25/ :04:27 PM
14.3 Hydrated Protons and Hydronium Ions
H+(aq) + A(aq)
HA(aq)
Due to high reactivity of the hydrogen ion, it is actually hydrated by one or more water molecules.
n = 4 H9O4+
n = 1 H3O+
n = 2 H5O2+
n = 3 H7O3+
[H(H2O)n]+
For our purposes, H+ is equivalent to H3O+.
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5. Examples
Write a balanced equation for the dissociation of each of the following Bronsted-Lowry acids in water
HCl (aq)
H2CO3(aq)
HSO4-(aq)

6. Examples
In the following equation, label each species as an acid or a base. Identify the conjugated acid-base pair
a. NO2-(aq) + HF(aq) HNO2(aq) + F-(aq)
b. HClO4(aq) + H2O(l) H3O+(aq) + ClO4-(aq)

7. 14.2 Acid Strength and Base Strength
Chemistry: McMurry and Fay, 6th Edition
Chapter 14: Aqueous Equilibria: Acids and Bases
4/25/ :04:27 PM
14.2 Acid Strength and Base Strength
H3O+(aq) + A(aq)
HA(aq) + H2O(l)
Acid
Base
Acid
Base
With equal concentrations of reactants and products, what will be the direction of reaction?
Weaker acid + Weaker base
Stronger acid + Stronger base
The direction of reaction will be in the direction of the weaker acid and weaker base.
The reaction almost goes completely to the right, to the formation of a weaker acid. Why?
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8. Acid Strength and Base Strength
Chemistry: McMurry and Fay, 6th Edition
Chapter 14: Aqueous Equilibria: Acids and Bases
Acid Strength and Base Strength
4/25/ :04:27 PM
Weak Acid: An acid that is only partially dissociated in water and is thus a weak electrolyte.
We dealt with weak acids initially in Ch 4 (Reactions in Aqueous Solution).
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9. Chemistry: McMurry and Fay, 6th Edition
Chapter 14: Aqueous Equilibria: Acids and Bases
4/25/ :04:27 PM
Copyright © 2011 Pearson Prentice Hall, Inc.

10. Examples
If you mix equal amount concentrations of reactants and products, decide which species (reactants or products) are favored at the completion of the reaction
H2SO4(aq) + NH3(aq) NH4+(aq) + HSO4-(aq)
HF(aq) + NO3-(aq) HNO3(aq) + F-(aq)
H2S(aq) + F-(aq) HF(aq) + HS-(aq)

11. 14.4 Dissociation of Water Chemistry: McMurry and Fay, 6th Edition
Chapter 14: Aqueous Equilibria: Acids and Bases
4/25/ :04:27 PM
14.4 Dissociation of Water
This value of Kw is only at 25 °C.
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12. Water autoionization As we have seen, water is amphoteric.
In pure water, a few molecules act as bases and a few act as acids.
This is referred to as autoionization.
H2O (l) + H2O (l) ⇌ H3O+ (aq) + OH- (aq
Kw = [H3O+] [OH−] at 25oC
This special equilibrium constant is referred to as the ion-product constant for water, Kw
Do you expect to have the same value of Kw at a temperature that is differ than 25oC?

13. Dissociation of Water Acidic: [H3O+] > [OH] Neutral:
Chemistry: McMurry and Fay, 6th Edition
Chapter 14: Aqueous Equilibria: Acids and Bases
4/25/ :04:27 PM
Dissociation of Water
Acidic:
[H3O+] > [OH]
Neutral:
[H3O+] = [OH]
Basic:
[H3O+] < [OH]
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14. Examples
At 50oC the value of Kw is 5.5 x What are the concentration of H3O+ and –OH in a neutral solution at 50oC?
Calculate the [-OH] in a solution with [H3O+] = 7.5 x 10-5M. Is this solution basic, acidic or neutral?

15. The pH Scale pH= log[H3O+] [H3O+] = 10 Acidic: pH < 7
Chemistry: McMurry and Fay, 6th Edition
Chapter 14: Aqueous Equilibria: Acids and Bases
4/25/ :04:27 PM
The pH Scale
pH= log[H3O+]
pH
[H3O+] = 10
Acidic: pH < 7
Neutral: pH = 7
Basic: pH > 7
The scale is logarithmic which means each unit change for pH corresponds to a factor of 10 change in hydronium ion concentration.
Also, since pH is dependent on Kw, this scale is only good at 25 °C.
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Copyright © 2011 Pearson Prentice Hall, Inc.

16. Chemistry: McMurry and Fay, 6th Edition
Chapter 14: Aqueous Equilibria: Acids and Bases
4/25/ :04:27 PM
The pH Scale
The hydronium ion concentration for lemon juice is approximately M. What is the pH when [H3O+] = M?
2 significant figures
pH = log(0.0025) = 2.60
2 decimal places
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17. Examples
Calculate the [H3O+] and pH of a solution with [-OH] = 6.4 x 10-8M
Calculate the concentration of H3O+ and –OH for a solution with a pH of 5.50

18. Ways to measure pH Three ways to measure pH Litmus paper blue-to-red:
red-to-blue:
basic, pH > 8
blue-to-red:
acidic, pH < 5
An indicator
A pH meter

19. Chemistry: McMurry and Fay, 6th Edition
Chapter 14: Aqueous Equilibria: Acids and Bases
14.6 Measuring pH
4/25/ :04:27 PM
Acid-Base Indicator: A substance that changes color in a specific pH range. Indicators exhibit pH-dependent color changes because they are weak acids and have different colors in their acid (HIn) and conjugate base (In) forms.
H3O+(aq) + In(aq)
HIn(aq) + H2O(l)
Color A
Color B
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20. 14.7 pH of strong acids and strong bases
A strong monoprotic acds – 100% dissociated in aqueous solution (HClO4, HCl, HBr, HI, H2SO4 HNO3 etc…)
Contains a single dissociable proton
HA(aq) H2O(l)  H3O+(aq) + A-(aq)
pH = - log H3O+
[H3O+] = [A-] = initial concentration of the acid
Undissociated [HA] = 0

21. Strong Bases Alkali metal hydroxide, MOH Water-soluble ionic solids
NaOH(aq)  Na+(aq) + -OH(aq)
Exits in aqueous solution as alkali metal cations and hydroxide anions
Calculate pH from [-OH]
Alkaline earth metal hydroxide, M(OH)2 where M= Mg, Ca, Sr, Ba
Less soluble than alkali hydroxide, therefore lower [-OH]

22. The pH in Solutions of Strong Acids and Strong Bases
Chemistry: McMurry and Fay, 6th Edition
Chapter 14: Aqueous Equilibria: Acids and Bases
4/25/ :04:27 PM
The pH in Solutions of Strong Acids and Strong Bases
What is the pH of a M solution of HNO3?
What is the pH of a M solution of HCl?
pH < 7 which makes sense since it’s an acid.
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23. Example What is the pH of a 0.025 M solution of NaOH?
Calculate the pH of a solution by dissolving 0.15 g of CaO in enough water to make 0.500L of limewater Ca(OH)2

24. 14.8 Equilibria in Solutions of Weak Acids
Chemistry: McMurry and Fay, 6th Edition
Chapter 14: Aqueous Equilibria: Acids and Bases
4/25/ :04:27 PM
14.8 Equilibria in Solutions of Weak Acids
H3O+(aq) + A(aq)
HA(aq) + H2O(l)
Acid-Dissociation Constant:
[H3O+][A]
Ka =
[HA]
pKa = - log Ka
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25. Chemistry: McMurry and Fay, 6th Edition
Chapter 14: Aqueous Equilibria: Acids and Bases
4/25/ :04:27 PM
Copyright © 2011 Pearson Prentice Hall, Inc.

26. Calculating Equilibrium Concentrations in Solution of Weak acids
Step 1: Write the balance equation for all possible proton transfer (both acids and water)
Step 2: Identify the principle reaction (the reaction that has larger Ka)
Step 3: Generate an ICE table
Step 4: Solve for x
Step 5: Calculate pH and all other concentrations (HA, H3O+ and A-)

27. Example
Determine the concentration of all species present (H3O+, CH3CO2-, CH3CO2H) and pH of a M CH3CO2H
Ka = 1.8 x 10-5

28. Example
Determine the concentration of all species present (H3O+, CH3CO2-, CH3CO2H) and pH of a 1.00 M CH3CO2H
Ka = 1.8 x 10-5

29. Equilibria in Solutions of Weak Acids
Chemistry: McMurry and Fay, 6th Edition
Chapter 14: Aqueous Equilibria: Acids and Bases
4/25/ :04:27 PM
Equilibria in Solutions of Weak Acids
The pH of M HF is What are the values of Ka and pKa for hydrofluoric acid?
H3O+(aq) + F(aq)
HF(aq) + H2O(l)
Worked Example 14.5 p553.
While the initial hydronium ion concentration is not 0, it’s insignificant when compared to that which comes from the dissociation of HF.
pH comes from the equilibrium concentration of hydronium ion.
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30. Percent Dissociation in Solutions of Weak Acids
Chemistry: McMurry and Fay, 6th Edition
Chapter 14: Aqueous Equilibria: Acids and Bases
4/25/ :04:27 PM
Percent Dissociation in Solutions of Weak Acids
[HA]dissociated
Percent dissociation =
x 100%
[HA]initial
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31. Example
Find the pH, [-OH] and percent ionization of a M HClO2 solution
Ka = 1.1 x 10-2

32. Example
A generic weak acid, formula = HA, has a concentration of M and is 1.235% dissociated. Determine the Ka.