Chapter 14Chapter 14Aqueous Equilibria: Acids Aqueous Equilibria: Acids andand

Arrhenius Acid: A substance that dissociates in water to produce hydrogen ions, H+ or increases the concentration of hydrogen ion

Arrhenius Base: A substance that dissociates in water to produce hydroxide ions, OH-1 increases the concentration of hydroxide ions.

M+(aq) + OH(aq)MOH(aq)

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

The Brønsted-Lowry Theory

Acid DissociationAcid 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.

Acid-Base Concepts: The Acid-Base Concepts: The Brønsted-Lowry TheoryBrønsted-Lowry Theory

Conjugate Acid-Base Pairs: Chemical species whose formulas differ only by one hydrogen ion, H+

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

14.3 Hydrated Protons and 14.3 Hydrated Protons and Hydronium IonsHydronium Ions

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

[H(H2O)n]+

For our purposes, H+ is equivalent to H3O+.

n = 4 H9O4+

n = 1 H3O+

n = 2 H5O2+

n = 3 H7O3+

Due to high reactivity of the hydrogen ion, it is actually hydrated by one or more water molecules.

ExamplesExamples Write a balanced equation for the dissociation of each

of the following Bronsted-Lowry acids in water

HCl (aq)

H2CO3(aq)

HSO4-(aq)

ExamplesExamples 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)

14.2 Acid Strength and Base 14.2 Acid Strength and Base StrengthStrength

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

BaseAcidBaseAcid

Weaker acid + Weaker baseStronger acid + Stronger base

With equal concentrations of reactants and products, what will be the direction of reaction?

The reaction almost goes completely to the right, to the formation of a weaker acid. Why?

Acid Strength and Base Acid Strength and Base StrengthStrengthWeak Acid: An acid that is only partially dissociated in water and is thus a weak electrolyte.

ExamplesExamples 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)

14.4 Dissociation of 14.4 Dissociation of WaterWater

Water autoionizationWater 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?

Dissociation of WaterDissociation of Water Acidic:

[H3O+] > [OH]

Neutral:

[H3O+] = [OH]

Basic:

[H3O+] < [OH]

ExamplesExamples At 50oC the value of Kw is 5.5 x 10-14. 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?

The pH ScaleThe pH Scale

© 2012 Pearson Education, Inc.

pH= log[H3O+]

[H3O+] = 10 pH

Acidic: pH < 7Neutral: pH = 7Basic: pH > 7

The pH ScaleThe pH ScaleThe hydronium ion concentration for lemon juice is approximately 0.0025 M. What is the pH when [H3O+] = 0.0025 M?

pH = log(0.0025) = 2.60

2 significant figures

2 decimal places

ExamplesExamples

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

Ways to measure pHWays to measure pH Three ways to measure pH Litmus paper

◦ red-to-blue:

◦ basic, pH > 8

blue-to-red:◦ acidic, pH < 5

An indicator A pH meter

14.6 Measuring pH14.6 Measuring pHAcid-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 BColor A

14.7 pH of strong acids and 14.7 pH of strong acids and strong basesstrong 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

Strong BasesStrong 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]

The pH in Solutions of The pH in Solutions of Strong Acids and Strong Strong Acids and Strong BasesBases What is the pH of a 0.050 M solution of HNO3?

What is the pH of a 0.100 M solution of HCl?

ExampleExample 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

14.8 Equilibria in Solutions of 14.8 Equilibria in Solutions of Weak AcidsWeak Acids

Acid-Dissociation Constant:

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

[H3O+][A]

[HA]Ka =

pKa = - log Ka

Calculating Equilibrium Calculating Equilibrium Concentrations in Solution Concentrations in Solution of Weak acidsof 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-)

ExampleExampleDetermine the concentration of all species

present (H3O+, CH3CO2-, CH3CO2H) and pH of a

0.0100 M CH3CO2H

Ka = 1.8 x 10-5

ExampleExampleDetermine the concentration of all species

present (H3O+, CH3CO2-, CH3CO2H) and pH

of a 1.00 M CH3CO2H

Ka = 1.8 x 10-5

Equilibria in Solutions of Equilibria in Solutions of Weak AcidsWeak Acids

The pH of 0.250 M HF is 2.036. What are the values of Ka and pKa for hydrofluoric acid?

H3O+(aq) + F(aq)HF(aq) + H2O(l)

Percent Dissociation in Percent Dissociation in Solutions of Weak AcidsSolutions of Weak Acids

Percent dissociation =[HA]initial

[HA]dissociated

x 100%

ExampleExample Find the pH, [-OH] and percent ionization of a 0.100

M HClO2 solution

Ka = 1.1 x 10-2

ExampleExample A generic weak acid, formula = HA, has a

concentration of 0.200 M and is 1.235% dissociated. Determine the Ka.