1. Chapter 9 Charge-Transfer Reactions: Acids and Bases and Oxidation-Reduction Copyright  The McGraw-Hill Companies, Inc. Permission required for reproduction or display.

2. 9.1 Acids and Bases Acids: Taste sour, dissolve some metals, cause plant dye to change color Bases: Taste bitter, are slippery, are corrosive. Two theories that help us to understand the chemistry of acids and bases. 1.Arrhenius Theory 2.Brønsted-Lowry Theory

3. Arrhenius Theory of Acids and Bases Acid - a substance, when dissolved in water, dis- sociates to produce hydrogen ions –Hydrogen ion: H + also called “protons” HCl is an acid: HCl(aq)  H + (aq) + Cl - (aq) Base - a substance, when dissolved in water, dissociates to produce hydroxide ions. NaOH is a base NaOH(aq)  Na + (aq) + OH - (aq) Where does NH 3 fit? When it dissolves in water it is basic but it does not have OH - ions in it.

4. Brønsted-Lowry Theory of Acids and Bases Acid - proton donor Base - proton acceptor Notice that it is not defined using water. When writing the reactions, both accepting and donation are evident. HCl(aq) + H 2 O(l)  Cl - (aq) + H 3 O + (aq) Now let us look at NH 3 and see why it is a base. NH 3 (aq) + H 2 O(l) NH 4 +(aq) + OH-(aq) baseacid baseacid

5. Conjugate Acids and Bases The acid base reaction can be written in the general form: HA + B A- + HB+ Notice the reversible arrows. The products are also an acid and base called the conjugate acid and base. Conjugate Acid - what the base becomes after it accepts a proton. Conjugate Base - what the acid becomes after it donates its proton. Conjugate Acid-Base Pair - The acid and base on the opposite sides of the equation. 2 acid basebase acid

6. Strong and Weak Acids The reversible arrow isn’t always written. Some acids or bases essentially dissociate 100% and a one way arrow is used. Example: HCl + H 2 O  Cl - + H 3 O + HCl is called a strong acid-an acid that dissociates 100% Weak acid - one which does not dissociate 100%.

7. 9.2 Solutions of Acids and Bases Strength of Acids and Bases Acid and base strength - degree of dissociation –Not a measure of concentration, different thing Strong acids and bases - reaction with water is virtually 100% (Strong electrolytes) Strong Acids: –HCl, HBr, HIHydrochloric Acid, etc. –HNO 3 Nitric Acid –H 2 SO 4 Sulfuric Acid Strong Bases: –NaOH, KOH, Ba(OH) 2 (all are metal hydroxides)

8. Weak acids and bases - only a small percent dissociates. (Weak electrolytes) Weak acid example: –Acetic acid: Weak base example: –Ammonia: CH 3 COOH(aq) + H 2 O(l) CH 3 COO - (aq) + H 3 O + (aq) NH 3 (aq) + H 2 O(l) NH 4 + (aq) + OH - (aq)

9. The Dissociation of Water Pure water is virtually 100% molecular. Very small number of molecules dissociate –Dissociation of acids and bases is often called ionization. Called autoionization. Very weak electrolyte. H + is called the hydrogen ion. In pure water at room temperature: [H + ] = 1 x 10 -7 M [OH - ] = 1 x 10 -7 M H 2 O(l) H + (aq) + OH - (aq)

10. Therefore the equilibrium expression for: Remember, liquids are not included. This constant is called the ion product for water and has the symbol K w Since [H + ] = [OH - ] = 1.0 x 10 -7 M, what is the value for K w ? 1.0 x 10 -14. –Remember, it is without units. H 2 O(l) H + (aq) + OH - (aq)

11. The pH Scale pH scale - a scale that indicates the acidity or alkalinity of a solution. –Ranges from 0 (very acidic) to 14 (very basic) As we do the problems, keep in mind that since 1 x 10 -14 = [H + ][OH - ], –if we know one concentration, can calculate the other, –if add an acid, [H + ]  and [OH - ]  –if add a base, [OH - ]  and [H + ]  The pH of a solution is defined as: 3 pH = -log[H + ]

12. 9.3 Reactions between Acids and Bases Neutralization reaction - the reaction of an acid with a base to produce a salt and water. HCl(aq) + NaOH(aq)  NaCl(aq) + H 2 O(l) An analytical technique to determine the concentration of an acid or base is the titration. Titration involves the addition of measured amount of a standard solution (solution of known concentration) to neutralize the second, unknown solution. The equivalence point is when the moles H + and OH - are equal. 7

13. Polyprotic Substances The previous examples have the acid and base at a 1:1 combining ratio. –Not all acid-bases do this. Polyprotic substance - donates or accepts more than one proton per formula unit. H 2 SO 4 (aq) + 2NaOH(aq)  Na 2 SO 4 (aq) + 2H 2 O(l) Other polyprotics: Nitric Acid, Sulfuric Acid, and Phosphoric Acid.

14. 9.4 Acid-Base Buffers Buffer solution - solution which resists large changes in pH when either acids or bases are added. The Buffer Process Buffers consist of either –a weak acid and its salt or –a weak base and its salt Examples: –Acetic acid (CH 3 COOH) with sodium acetate (CH 3 COONa). –An equilibrium is established in solution between the acid and the salt anion. 8 CH 3 COOH(aq) + H 2 O(l) CH 3 COO - (aq) + H + (aq)

15. Addition of Base (OH - ) to a Buffer Solution. The OH - will react with the H +, removing it from the above equilibrium. Which way will the equilibrium shift? To the right. Addition of Acid (H + ) to a Buffer solution. The acid increases the concentration of H +. Which way will the equilibrium shift? To the left. CH 3 COOH(aq) + H 2 O(l) CH 3 COO - (aq) + H + (aq)

16. Buffer Capacity - a measure of the ability of a solution to resist large changes in pH when a strong acid or strong base is added. Preparation of a Buffer Solution CH 3 COOH(aq) + H 2 O(l) CH 3 COO - (aq) + H + (aq)

17. 9.5Oxidation-Reduction Reactions Oxidation - defined by one of the following –loss of electrons –loss of hydrogen atoms –gain of oxygen atoms Example: Na  Na + + e - –Oxidation half of the reaction Reduction - defined by one of the following: –gain of electrons –gain of hydrogen –loss of oxygen Example: Cl + e -  Cl - –Reduction half of the reaction 9

18. Na  Na + + e - Cl + e -  Cl - Na + Cl  Na + + Cl - Oxidizing Agent Is reduced Gains electrons Causes oxidation Reducing Agent Is oxidized Loses electrons Causes reduction

19. 9 Applications of Oxidation and Reduction Corrosion - the deterioration of metals caused by an oxidation-reduction process. –Example: rust (oxidation of iron) 4Fe(s) + 3O 2 (g)  2Fe 2 O 3 (s) Combustion of Fossil Fuels –Example: natural gas (methane) furnaces. CH 4 (g) + 2O 2 (g)  CO 2 (g) + 2H 2 O(g) Bleaching - Most bleaching agents are oxidizing agents. The oxidation of the stains produces compounds that do not have color.