1. Properties of acids Electrolytes: conduct electricity React to form salts Change the color of an indicator Have a sour taste

2. Properties of Bases Bitter taste “slippery feel” Electrolyte React with an acid to form a salt

3. Acid names: ide=  hydro______ic acid ite ==  ____________ous acid ate ==  _____________ic acid

4. Arrhenius Acids and Bases An acid contains a Hydrogen ion that easily disassociates A base has a hydroxide ion that easily disassociates

5. Bronsted-Lowry An acid is a hydrogen ion donor A base is a hydrogen ion acceptor. NH 3(aq) + H 2 O (aq)  NH 4 + (aq) + OH - (aq) Base Acid Conj. Acid Conj. Base

6. conjugate acids and bases These are the reverse reaction “reactants” HA (aq) + H 2 O (l)  H 3 O + (aq) + A - (aq) Acid Base conjugate conjugate acid base HCl + NaOH  HOH + Cl -

7. Lewis acid and bases An acid is an electron pair acceptor A base is an electron pair donator base acid NH 3 + BF 3 H 3 N BF 3

8. Acid Strength A strong acid is one for which the equilibrium lies far to the products side K a > 1. A weak acid is one for which the equilibrium lies far to the reactants side K a < 1. Table 14.1 on page 659.

9. Monoprotic: HCl Diprotic: H 2 SO 4 Triprotic H 3 PO 4

10. Oxyacids Acids where the acidic hydrogen is attached to an oxygen. Pictures on pages 658-659. Structural formulas given in table 14.8 on page 694.

11. Organic acids Acids with a carbon atom backbone. Commonly contain a carboxyl group. Acetic acid Benzoic acid

12. Amphoteric Substance A substance that can act as both an acid and a base. Water Ammonia

13. We can tell an acid from a base by using an indicator

14. Autoionization of water H 2 O + H 2 O  H 3 O + + OH - K w = [H 3 O + ] [OH - ] = 1.0 x 10 -14 [H 3 O + ] [OH - ] = [H + ] [OH - ] [H + ] = [OH - ] =1.0 x 10 -7 M

15. pH = -log [H + ] [H + ] = antilog (-pH) pOH = -log[OH - ] pK = -log K Significant figures in logarithms- the # of decimal places in the log = # of sig figs in the original number.

16. [H + ] [OH - ] = 1.0 x 10 -14 pH + pOH = 14 For strong acids the [H + ] = the molarity of the acid.

17. System for solving weak acid equilibrium problems List the major species in solution. Find any species that can produce H + and write a balanced equation for the reaction producing H +. Use the values for K for the reactions you have written to decide which reaction will dominate.

18. Write the equilibrium expression for the dominant reaction. List the initial concentrations of the species in the dominant reaction. Define the change needed to obtain equilibrium. Define x.

19. Write equilibrium concentrations in terms of x. Put equilibrium concentrations into equilibrium expression. Solve for x the “easy way” Use the 5% rule to see if approximation is valid. Calculate [H + ] and pH.

20. Calculate the pH of a 0.100M solution of hypochlorous acid.

21. Calculate the pH of a solution that contains 1.00M HCN and 1.00M HNO 2. Also calculate the concentration of cyanide ions.

22. % dissociation = amount dissociated x100 initial concentration For a given weak acid, the % dissociated increases as the acid becomes more dilute.

23. In a 0.100M solution, lactic acid (HC 3 H 5 O 3 )is 3.7% dissociated. Calculate the value of K a for this acid.

24. Bases B (aq) + H 2 O (l)  HB + (aq) + OH - (aq) Base acid conj conj acid base K b = [BH + ][OH - ] [B]

25. Bases Strong bases – hydroxides of group 1A metals and calcium, barium and strontium. Weak bases are commonly ammonia and substituted ammonia compounds. Table 14.3 on page 678.

26. Calculate the pH for a 15.0M solution of NH 3. K b = 1.8 x10 -5

27. Salts that produce neutral solutions Salts that consist of the cations of strong bases and the anions of strong acids have no effect on pH. KCl, NaNO 3, Ba(HSO 4 ) 2

28. Salts that produce basic solutions Salts that consist of cations of strong bases and the anion is the conjugate base of a weak acid. NaC 2 H 3 O 2, KNO 2, Sr(CN) 2

29. Salts that produce acidic solutions Salts that consist of a cation that is the conjugate acid of a weak base and the anion of a strong acid. NH 4 Cl A salt that contains a highly charged metal ion. AlCl 3 see sample exercise 14.20.

30. For any weak acid-conjugate base or weak base-conjugate acid K a x K b = K w

31. Calculate the pH of: 0.10M NaCl 0.10M NaF 0.10M NH 4 Cl

32. If the anion is the conjugate base of a weak acid and the cation is the conjugate acid of a weak base the K b must be compared to the K a. Which ever is greater will dominate.

33. Predict if the following will be acidic or basic NH 4 C 2 H 3 O 2 NH 4 CN NH 4 NO 2

34. Effect of structure on acid-base properties Acidic properties depend on two factors: H-X The strength of the bond. As the strength increases the acidity decreases The polarity of the bond. As polarity increases acidity increases

35. Strength of oxyacids Within a series of oxyacids as the number of oxygens increases the strength of the acid also increases. Table 14.8 page 694

36. Acid-base properties of oxides Depends on the electronegativity of the element bonded to oxygen. O-X Non-metal oxides in water will form acids. O-X is stronger than H-O in polar water. Metal oxides in water form bases. O-H stronger than O-X in polar water.

37. Common Ion Effect The shift in equilibrium position because of the addition of an ion already involved in the equilibrium. Equilibrium shifts away from the added component.

38. Buffered Solutions Resists change in pH when either hydrogen or hydroxide ions are added. Consist of a solution that contains both a weak acid and its salt or a weak base and its salt. Important Characteristics are on page 726.

39. Buffer Capacity Represents the amount of hydrogen or hydroxide ions the buffer can absorb without a significant change in pH. Best buffers contain a weak acid with a pKa as close as possible to the desired pH.

40. Indicators A substance that changes color depending on the pH of the solution it is in. When choosing an indicator we want the indicator end point and the titration equivalence point to be as close as possible.