1. COMMON ION EFFECT LEWIS ACIDS & BASES [16.11-17.1]

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 _____, causes a ___crease in [H + ] & ______ pH

3. Another example  A solution of NH 4 Cl & NH 3 NH 4 Cl  NH 4 + + Cl - NH 3 + H 2 O  NH 4 + + OH - Shifts ______, causes a ___crease in [OH - ] & ____ 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 (HC 2 H 3 O 2 ) and 0.30 mol of sodium acetate (NaC 2 H 3 O 2 ) 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 (HC 2 H 3 O 2 ) and a strong electrolyte (NaC 2 H 3 O 2 ) that share a common ion, C 2 H 3 O 2 –. 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. (We have written the equilibrium using H + (aq) rather than H 3 O + (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: Solve: First, because HC 2 H 3 O 2 is a weak electrolyte and NaC 2 H 3 O 2 is a strong electrolyte, the major species in the solution are HC 2 H 3 O 2 (a weak acid), Na + (which is neither acidic nor basic and is therefore a spectator in the acid-base chemistry), and C 2 H 3 O 2 – (which is the conjugate base of HC 2 H 3 O 2 ). Second, [H + ] and, therefore, the pH are controlled by the dissociation equilibrium of HC 2 H 3 O 2 :

5. SAMPLE EXERCISE 17.1 continued The equilibrium concentration of C 2 H 3 O 2 – (the common ion) is the initial concentration that is due to NaC 2 H 3 O 2 (0.30 M) plus the change in concentration (x) that is due to the ionization of HC 2 H 3 O 2. Now we can use the equilibrium-constant expression: (The dissociation constant for HC 2 H 3 O 2 at 25°C is from Appendix D; addition of NaC 2 H 3 O 2 does not change the value of this constant.) Substituting the equilibrium-constant concentrations from our table into the equilibrium expression gives Because K a is small, we assume that x is small compared to the original concentrations of HC 2 H 3 O 2 and C 2 H 3 O 2 – (0.30 M each). Thus, we can ignore the very small x relative to 0.30 M, giving

6. 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 HC 2 H 3 O 2 has a pH of 2.64, corresponding to [H + ]  2.3  10 –3 M.Thus, the addition of NaC 2 H 3 O 2 has substantially decreased [H + ], as we would expect from Le Châtelier’s principle. 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. In your book page 722-723!

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  Acidic oxides: covalent oxide dissolves in water CO 2 + H 2 O  H 2 CO 3 natural rain water SO 2 + H 2 O  H 2 SO 3 acid rain 2NO 2 + H 2 O  HNO 3 + HNO 2  Basic oxides: ionic oxide dissolves in water CaO + H 2 O  Ca(OH) 2 lime soil K 2 O + H 2 O  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.