Fig. 9-CO, p. 251

1. Chapter 7 Acids and Bases Fig. 9-CO, p. 251

2. Base Accepts Acid Donates Bronsted-Lowry Definition

3. The hydrogen ion is a proton H+

4. (base) (acid) H2O HCl + + H3O+ Cl-

5. Brønsted-Lowry Acids & Bases • We can use curved arrows to show the transfer of a proton from acetic acid to ammonia:

6. H2O NH3 + H2O NH3 OH- NH4+ + + OH- NH4+ +

7. H2O NH3 OH- NH4+ + + (acid) (base) (base) (acid)

8. Brønsted-Lowry Acids & Bases

9. Identify the Acid and the Base in Each Chemical Equation • HBr + H2O → H3O+ + Br- • H2O + CN- → HCN + OH- • HF + H2O → H3O+ + F- ←

12. Acid-Base Equilibria • HClis a strong acid, which means that the position of this equilibrium lies very far to the right. • In contrast, acetic acid is a weak acid, and the position of its equilibrium lies very far to the left.

13. Amphoteric A substance whose ability to behave as an acid is about the same as its ability to behave as a base.

14. -

16. For pure water… ] [ = 0.0000001 M OH- = 10-7 M ] [ = 0.0000001 M H3O+ = 10-7 M

17. ] [ ] [ = Kw OH- H3O+ RULE x 10-7 10-7 = 10-14 0.0000001 M X 0.0000001 M = 0.00000000000001 M

18. The equation for the ionization of water applies not only to pure water but also to any aqueous solution. • The product of [H3O+] and [OH-] in any aqueous solution is equal to 1.0 x 10-14. • For example, if we add 0.010 mol of HCl to 1.00 liter of pure water, it reacts completely with water to give 0.010 mole of H3O+. • In this solution, [H3O+] is 0.010 or 1.0 x 10-2. • This means that the concentration of hydroxide ion is:

20. pH and pOH • Because hydronium ion concentrations for most solutions are numbers with negative exponents, we commonly express these concentrations as pH, where: pH = -log [H3O+] • We can now state the definitions of acidic and basic solutions in terms of pH: • Acidic solution: one whose pH is less than 7.0. • Basic solution: one whose pH is greater than 7.0. • Neutral solution: one whose pH is equal to 7.0.

21. log 107 log 10-5 = = 7 -5 log 0.01 = 10-2 = -2 = 100,000 Antilog 5 Log 1.4 x 10-9 Antilog 5.3 = =

22. What is the pH of a solution with a hydronium ion conc. of 3.2 x 10-9M? ] [ pH = -log H3O+ [ ] = -log 3.2 x 10-9 = - (-8.5) = 8.5

23. What is the pH of a solution with a hydroxide ion conc. of .001 M?

24. The pH of a solution is 10.0. What is the [H3O+] and the [OH-]?

26. A buffer solution contains at least two components: A component to neutralize any incoming base A component to neutralize any incoming acid

27. pH Buffers • The most common buffers consist of approximately equal molar amounts of a weak acid and a salt of the weak acid; that is, approximately equal molar amounts of a weak acid and a salt of its conjugate base. • For example, if we dissolve 1.0 mole of acetic acid and 1.0 mole of its conjugate base (in the form of sodium acetate) in water, we have an acetate buffer.

28. p. 272

30. pH Buffers • The effect of a buffer can be quite dramatic • Consider a phosphate buffer prepared by dissolving 0.10 mole of NaH2PO4 (a weak acid) and 0.10 mole of Na2HPO4 (the salt of its conjugate base) in enough water to make 1 liter of solution.

31. The average pH of human blood is 7.4. • any change greater than 0.10 pH unit in either direction can cause illness. • Acidosis and Alkalosis Describe condtions when the pH of blood is too high or too low. Chemical Connections 9D, p. 277

32. Salt An ionic compound formed from the reaction between an acid and a base. Neutralization The reaction between an acid and a base to give a salt and water.

33. Salt formation Hydrogen chloride Sodium hydroxide Sodium chloride Water (acid) (base) (salt) HCl + NaOH → NaCl + H2O (acid) (base) (salt) (water)