1. IV. IONIC EQUILIBRIUM / A. Review of Electrolytes

1. IV. IONIC EQUILIBRIUM / A. Review of Electrolytes

IV. IONIC EQUILIBRIUM

A. Review of Electrolytes

1. Electrolytes: Substances that dissociate, or ionize, in water to produce “free” ions.

They include salts, acids and bases.

2. Salts

a) Soluble salts dissociate or ionize completely in water, and are strong electrolytes.

Example(1): The following salts are soluble in water. Write the equation showing the products produced when each is placed in water.

a) NaCl

b) BaCl2

c) Na3PO4

b) Insoluble salts dissociate or ionize only partially in water, and are weak electrolytes.

Example(2): The following salts are classified as insoluble in water. Write the equation showing the equilibrium produced when each is placed in water.

a) AgCl

b) PbI2

c) Ca3(PO4)2

3. Acids: are proton (H+) donors.

Acids are covalent compounds. So unlike the salts, the ionization of an acid is

not a simple process of dissolving. An acid produces ions in water by a

reaction with the water.

a) Strong acids react completely (or nearly) with water.

Example(3): Write the equation showing the reaction of the following strong acids with water.

a) HCl

b) HNO3

c) HClO4

b) Weak acids react only partially with water.

Example(4): Write the equilibrium equation showing the reaction of the following weak acids with water.

a) HC2H3O2

b) HNO2

c) HClO

4. Bases: produce hydroxide ion in water.

a) Ionic compounds containing OHˉ dissolve in water to produce free OHˉ.

Example(5): The following hydroxide compounds are soluble in water. Write the equation showing the products produced when each is placed in water.

a) NaOH

b) Ba(OH)2

Example(6): The following hydroxide compounds are classified as insoluble in water. Write the equation showing the products produced when each is placed in water.

a) AgOH

b) Ca(OH)2

b) Covalent bases react with the water to produce the OHˉ.

Bases: are proton (H+) acceptors.

Example(7): The following covalent compounds behave as weak proton (H+) acceptors in water. Write the equation showing their reaction with water.

a) NH3

b) C2H5NH2

1. IV. IONIC EQUILIBRIUM / B. Auto-Ionization of Water

B. Auto-Ionization of Water

1. Equation and equilibrium expression, Kc

2. Kw, The ion product

Example(1): Use Kw to find the [H3O+] and [OH¯] in pure water.

1. IV. IONIC EQUILIBRIUM / C. The pH Scale

C.The pH Scale

1. Definition of pH

2. pH of strong acids and bases

a) Strong acids

Example(1): If 0.1 moles of HCl are added to water to make 1 liter of solution, what is the pH of the solution?

Example(2): What is the pH of a 0.01M solution of HCl.

Example(3): What is the pH of a 0.001M solution of HCl?

[H+]0.10.010.001//0.0000001

pH 1 2 3 //7

Example(4): What is the pH of a 1.6 x 10–4 Molar solution of HCl?

b) Strong bases

Example(5): If 0.01 moles of NaOH are added to water to make one liter of solution, what is the pH?

Example(6): What is the pH of a 0.1M solution of NaOH?

Example(7): Complete the table below

[H+] / 1x10–1 / 1x10–2 / 1x10–3 / …
…
…
… / 1x10–7 / ….
….
….
….
[OH–] / 1x10–2 / 1x10–1
[H+][OH–]
pH / 1 / 2 / 3 / 7 / 12 / 13

3. Interpretation of the pH scale

Example(8): a) Which is a more acidic solution, pH=3 or pH=1.

b) How much more acidic is one solution than the other? (i.e. by what factor is the [H+] greater?)

Example(9): a) If the pH of a solution is decreased by 4 units is the solution becoming more acidic or basic?

b) By what factor does the acidity change? (i.e. by what factor does the [H+] change?)

Example(10): Which is a more basic solution, pH=8 or pH=10? By what factor?

1. IV. IONIC EQUILIBRIUM / D. Weak Acid and Base Equilibrium

D. Weak Acid and Base Equilibrium

1. The equilibrium constant for weak acids, Ka

2. Using Ka to find equilibrium concentrations and pH.

Example(1): a) In a 0.50 M solution of acetic acid, what are the equilibrium concentrations of HC2H3O2, C2H3O2‾, and H+? Ka = 1.8 x 10‾5

b) What is the pH of the solution? .

Example(2): a)In a 0.020 M solution of HClO acid, what are the equilibrium concentrations of HClO,ClO‾, and H+? Ka = 3.0 x 10‾8.

b) What is the pH of the solution?

3. Using pH to find Ka

a) High pH, small Ka

Example(3): The pH of a 0.20 M solution of the acid HCN is measured to be 5.00. Determine the value of the Ka for HCN.

b) Low pH, large Ka

Example(4): The pH of a 0.10 M solution of the acid HClO2 is measured to be 1.57, determine the value of the Ka for HClO2.

4. The equilibrium constant for weak bases, Kb

5. Using Kb to find equilibrium concentrations and pH.

Example(5): Find the concentrations of all species in a 0.10M solution of NH3 and the pH of solution. Kb = 1.8 x 10‾5.

6. Using pH to find Kb

Example(6): The pH of a 1.00M solution of C5H5N is 9.6, find the Kb.

7. Acid and base properties of salts

a) Salts made from a strong acid and strong base

Example(7): HCl + NaOH → NaCl + H2O

b) Salts made from a strong acid and weak base

Example(8): HCl + NH3 → NH4Cl

c) Salts made from a weak acid and strong base

Example(9): HC2H3O2 + NaOH → NaC2H3O2 + H2O

1. IV. IONIC EQUILIBRIUM / E. The Common Ion Effect

E. The Common Ion Effect

1. Review and application of LeChatelier’s Principle

Example(1): What is the pH of a 0.50M solution of HC2H3O2 that is also 0.50M NaC2H3O2?

Ka = 1.8 x 10‾5.

Example(2): What is the pH of a 2.0M solution of HF that is also 0.50M NaF? Ka = 6.8 x 10‾4.

Example(3): What is the pH of a 0.15M solution of NH3 that is also 0.30M NH4Cl? Kb = 1.8 x 10‾5.

2. Buffered solutions.

a)Buffered Solutions: maintains an essentially constant pH when a small amount

of acid or base is added to it.

b) A buffered solution is made from a weak acid and a salt of the acid, or a weak base and a salt of the base.

Example(4): What salt would you use with the following acids and bases to make a buffered solution?

a) HC2H3O2

b) HF

c) H2CO3

d) NH3

3. How a buffer works

Example(5): HF and NaF

Example(6): NH3 and NH4Cl

4. Polyprotic acids

a) Polyprotic Acid: has more than one ionizable H’s per molecule.

Example(7): Give the formulas of three polyprotic acids.

Example(8): Write the equations showing the stepwise ionization of H3PO4 in water.

Example(9): Write the expression for the equilibrium constant for each of the above steps.

Example (10): For H2S Ka1 = 9.5 x 10‾8 and Ka2 = 1 x 10‾19. For a 2.00 M solution of H2S find the concentrations of all species at equilibrium.

Example (11): For H2SO4 Ka1 = “large” and Ka2 = 1.0 x 10‾2. For a 2.5 M solution of H2SO4 find the concentrations of all species at equilibrium.

1. IV. IONIC EQUILIBRIUM / F. Solubility Product

F. Solubility Product

1. The equilibrium constant for dissolving a salt, Ksp

2. Finding Ksp from solubility

Example(1): At 25ºC the solubility of AgCl is 1.3 x 10‾5 moles per liter. Determine the Ksp for AgCl.

Example(2): At 25ºC the solubility of PbI2 is 1.5 x 10‾3 moles per liter. Determine the Ksp for PbI2.

3. Using Ksp to find the molar solubility

Example(3): For AgBr the Ksp = 5.0 x 10‾13 at 25ºC. Find its solubility in moles per liter.

Example(4): For Ag2CrO4 the Ksp = 1.1 x 10‾12 at 25ºC, find its solubility.

Example(5): At 25ºC the Ksp of Fe(OH)3 is 1.6 x 10‾39. Find the concentrations of all ions in a saturated solution.

4. Using Ksp to predict if a precipitate will form

a) If the ion product > Ksp a precipitate will form

b) If the ion product ≤ Ksp a precipitate will not form

Example(6): If a solution is 1.0 x 10‾5M NaCl and 2.0 x 10‾4M AgNO3, will the precipitate AgCl form?

Ksp of AgCl = 1.8 x 10‾10 at 25ºC.

Example(7): If a solution is 1.0 x 10‾3M NaI and 2.0 x 10‾3M Pb(NO3)2, will the precipitate PbI2 form?

Ksp of PbI2 = 1.4 x 10‾8 at 25ºC.