1. Common Ion Effect Lewis Acids & Bases
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2. Common Ion…again
As with solubility problems, usually make the pH of acids & bases less dramatic than expected; lower concentrations of H+ & OH-
For example…a solution of NaF & HF
NaF  Na+ + F-
HF  H+ + F-
Common ion
Shifts equilibrium to the left, causes a decrease in [H+] & higher pH

3. Another example A solution of NH4Cl & NH3 NH4Cl  NH4+ + Cl-
NH3 + H2O  NH OH-
Shifts left, causes a decrease in [OH-] & lower pH

4. SAMPLE EXERCISE 17.1 Calculating the pH When a Common Ion Is Involved
What is the pH of a solution made by adding 0.30 mol of acetic acid (HC2H3O2) and 0.30 mol of sodium acetate (NaC2H3O2) to enough water to make 1.0 L of solution?
Solution  
Analyze: We are asked to determine the pH of a solution of a weak electrolyte (HC2H3O2) and a strong electrolyte (NaC2H3O2) that share a common ion, C2H3O2–.
Plan: In any problem in which we must determine the pH of a solution containing a mixture of solutes, it is helpful to proceed by a series of logical steps:
1. Identify the major species in solution, and consider their acidity or basicity.
2. Identify the important equilibrium that is the source of H+ and therefore determines pH.
3. Tabulate the concentrations of ions involved in the equilibrium.
4. Use the equilibrium-constant expression to calculate [H+] and then pH.
Solve: First, because HC2H3O2 is a weak electrolyte and NaC2H3O2 is a strong electrolyte, the major species in the solution are HC2H3O2 (a weak acid), Na+ (which is neither acidic nor basic and is therefore a spectator in the acid-base chemistry), and C2H3O2– (which is the conjugate base of HC2H3O2).
Second, [H+] and, therefore, the pH are controlled by the dissociation equilibrium of HC2H3O2:
(We have written the equilibrium using H+(aq) rather than H3O+(aq), but both representations of the hydrated hydrogen ion are equally valid.)
Third, we tabulate the initial and equilibrium concentrations much as we did in solving other equilibrium problems in Chapters 15 and 16:

5. SAMPLE EXERCISE 17.1 continued
The equilibrium concentration of C2H3O2– (the common ion) is the initial concentration that is due to NaC2H3O2 (0.30 M) plus the change in concentration (x) that is due to the ionization of HC2H3O2.
Now we can use the equilibrium-constant expression:
(The dissociation constant for HC2H3O2 at 25°C is from Appendix D; addition of NaC2H3O2 does not change the value of this constant.) Substituting the equilibrium-constant concentrations from our table into the equilibrium expression gives
Because Ka is small, we assume that x is small compared to the original concentrations of HC2H3O2 and C2H3O2– (0.30 M each). Thus, we can ignore the very small x relative to 0.30 M, giving

6. SAMPLE EXERCISE 17.1 continued
The resulting value of x is indeed small relative to 0.30, justifying the approximation made in simplifying the problem.
Finally, we calculate the pH from the equilibrium concentration of H+(aq):
Comment: In Section 16.6 we calculated that a 0.30 M solution of HC2H3O2 has a pH of 2.64, corresponding to [H+] = 2.3  10–3 M.Thus, the addition of NaC2H3O2 has substantially decreased [H+], as we would expect from Le Châtelier’s principle.
In your book page !

7. Factors Affecting Acid Strength
In oxyacids, in which an OH is bonded to another atom, Y, the more electronegative Y is, the more acidic the acid.

8. Factors Affecting Acid Strength
For a series of oxyacids, acidity increases with the number of oxygens.

9. Factors Affecting Acid Strength
Resonance (look for the double bond or π bond) in the conjugate bases of carboxylic acids stabilizes the base and makes the conjugate acid more acidic.

10. Oxides Basic oxides: ionic oxide dissolves in water
Acidic oxides: covalent oxide dissolves in water
CO2 + H2O  H2CO3 natural rain water
SO2 + H2O  H2SO3 acid rain
2NO2 + H2O  HNO3 + HNO2
Basic oxides: ionic oxide dissolves in water
CaO + H2O  Ca(OH)2 lime soil
K2O + H2O  2 KOH

11. Lewis Acids Lewis acids are defined as electron-pair acceptors.
Atoms with an empty valence orbital can be Lewis acids.

12. Lewis Bases Lewis bases are defined as electron-pair donors.
Anything that could be a Brønsted–Lowry base is a Lewis base.
Lewis bases can interact with things other than protons, however.

13. Metal-Ligand Bond
This bond is formed between a Lewis acid and a Lewis base (coordinate covalent bond).
The ligands (Lewis bases) have nonbonding electrons.
The metal (Lewis acid) has empty orbitals.

14. Metal-Ligand Bond
The coordination of the ligand with the metal can greatly alter its physical properties, such as color, or chemical properties, such as ease of oxidation.