1. Acids and Bases
Chapters 14 and 15

2. Properties of acids Aqueous solutions have a sour taste.
Acids change the color of acid-base indicators.
Some acids react with active metals to release hydrogen gas, H2.
Acids react with bases to produce salts and water (i.e. a neutralization reaction)
Some acids conduct electric current.

3. Properties of bases Aqueous solutions taste bitter.
Bases change the color of acid-base indicators.
Dilute aqueous solutions of bases feel slippery (ex. soap).
Bases react with acids to produce salts and water. (i.e. neutralization reactions)
Bases conduct electric current.

4. Theories of acids and bases
Theory
Acid properties
Base properties
Arrhenius
(1887)
H+ ion produced
OH- ion produced
Brønsted-Lowry
(1923)
Proton donor
Proton acceptor
Lewis
e- pair acceptor
e- pair donor

5. Strong vs. weak
Strong acids and bases are those that ionize completely in aqueous solution. They are also typically strong electrolytes.
Weak acids and bases are those that do not ionize completely in aqueous solution. They are also typically weak electrolytes.

6. Common strong and weak acids
Strong acids
Weak acids
H2SO4 + H2OH3O+ + HSO4-
HSO4- + H2O  H3O+ + SO42-
HClO4 + H2O H3O+ + ClO4-
H3PO4 + H2OH3O++ H2PO4-
HClO3 + H2O H3O+ + ClO3-
HF + H2O  H3O+ + F-
HCl + H2O  H3O+ + Cl-
CH3COOH +H2OH3O+ + CH3COO-
HNO3 + H2O  H3O+ + NO3-
H2CO3 +H2OH3O+ + HCO3-
HBr + H2O  H3O+ + Br-
HCO3- + H2OH3O+ + CO32-
HI + H2O  H3O+ + I-
H2S + H2O  H3O+ + HS-
HCN + H2O  H3O+ + CN-
Page 474 of your textbook

7. Common strong and weak bases
Strong bases
Weak bases
Ca(OH)2 Ca2+ + 2OH-
NH3 + H2O  NH4+ + OH-
Sr(OH)2  Sr2+ + 2OH-
C6H5NH2+ H2O  C6H5NH3+ + OH-
Ba(OH)2 Ba2+ + 2OH-
NaOH  Na+ + OH-
KOH  K+ + OH-
RbOH  Rb+ + OH-
CsOH  Cs+ + OH-
p. 475 of text

8. Conjugate acid-base reactions
HF (aq) + H2O (l)  F- (aq) + H3O+ (aq) c.a.1 c.b.2 c.b.1 c.a.2

9. Amphoteric substances are those that
can react as either an acid or a base.

10. Self-ionization of water
H2O(l) + H2O(l)  H3O+ (aq) + OH- (aq) Kw = [H3O+] [OH-] = 1 x Measurements at 25˚C show that the [H3O+] and [OH-] are each 1 x 10-7 M

11. [H3O+] and [OH-] are inversely proportional
When something is acidic [H3O+] > [OH-];
[H3O+] >1 x 10-7 M and [OH-] < 1 x 10-7 M
When something is basic [H3O+] < [OH-];
[H3O+] <1 x 10-7 M and [OH-] > 1 x 10-7 M
[H3O+] and [OH-] are inversely proportional

13. pH means proportion of H+ ion
scale from 0 – 14
pH = -log[H3O+] pOH = -log[OH-]
As Kw = 1 x we can conclude that
pH + pOH = 14

14. Basic pH 7-14 [H3O+] < [OH-]; [H3O+] < 1 x 10-7 M
pOH
[OH-] 1 14
1 x M
1 x 10-1 M 13
1 x M 2
1 x 10-2 M 12
1 x M 3
1 x 10-3 M 11
1 x M 4
1 x 10-4 M 10
1 x M 5
1 x 10-5 M 9
1 x 10-9 M 6
1 x 10-6 M 8
1 x 10-8 M 7
1 x 10-7 M
Acidic pH 0-7
[H3O+] > [OH-];
[H3O+] > 1 x 10-7 M
[OH-] < 1 x 10-7 M
pH = 7 NEUTRAL
[H3O+] = [OH-]
1 x 10-7 M
Basic pH 7-14
[H3O+] < [OH-];
[H3O+] < 1 x 10-7 M
[OH-] > 1 x 10-7 M

16. pH of common items

17. Acid – base indicators
compounds whose colors are sensitive to pH. Indicators change colors because they are either weak acids or weak bases. In basic solution HIn  H+ + In- In acidic solution

18. Indicators come in many different colors
Indicators come in many different colors. The pH range over which an indicator changes color is called its transition interval.

19. Blue litmus paper turns red under acidic conditions and red litmus paper turns blue under basic (i.e. alkaline) conditions

21. A pH meter determines the pH of a solution by measuring the voltage between the two electrodes that are placed in the solution.

22. Neutralization reactions
DR between strong acids and bases always form water and a salt (an ionic compound formed from the cation of the base and the anion of the acid). The net ionic equation is:
H3O+ (aq) + OH- (aq)  2H2O (l)

23. Titration is the controlled addition and measurement of the amount of a solution of known concentration required to react completely with a measured amount of a solution of unknown concentration.

24. Buffers
solutions composed of a weak acid or base and its salt (ex. CH3COOH and NaCH3COO) and can withstand large changes in pH.
Buffer capacity is the amount of acid or base a buffer solution can absorb without a significant change in pH. This has very important biological implications.

25. Buffer examples
Ex. Blood maintains a pH of 7.4. The conjugate acid-base pair that acts as a buffer system is H2CO3 and HCO3-.
When there is excess acid present
H3O+ + HCO3-  H2O + H2CO3
When there is excess base present
OH- + H2CO3  H2O + HCO3-