1. Acids and Bases

2. Definitions of Acids and Bases
The Arrhenius Definition
An acid is a substance that dissociates in water to produce H+ ions.
HCl
A base is a substance that dissociates in water to produce OH- ions.
NaOH

3. Definitions of Acids and Bases
The Bronsted-Lowry Definition
An acid is any substance that can donate H+ ions (proton donor)
A base is any substance that can accept H+ ions. (proton acceptor)

4. The Hydronium Ion
the proton (H+) is strongly attracted to the electrons of water molecules and forms a hydronium ion (H3O+).

5. HCl + H2O  H3O+ + Cl-
NH3 + H2O  NH4+ + OH-

6. Water is amphoteric
-can act as an acid or base!

7. Definitions of Acids and Bases
Lewis Definition
-an acid is an electron pair acceptor
-base is an electron pair donor.

8. Properties of Acids and Bases
Appearance: both look much like water
Taste: Acids: sour (citrus fruits, pop)
Bases: bitter (soap)
Touch: Acids: if dilute, feel like water.
Bases: mild solutions feel smooth and slippery (soap)
Strong solutions of both will burn
Reaction with metals:
Acids: react vigorously with metals
Bases: do not react with metals

9. Strengths of Acids and Bases
Strong acid: completely dissociates in water.
Weak acids only partially dissociate in water.
1 mole of acetic acid in 1 liter of water only dissociates 0.4%.

10. Strengths of Acids and Bases
Strong base: have strong affinity for H+ ions (react completely with water and form OH- ions.)
Nearly any compound with the OH- ion included.
Weak base: only partially react with water to form OH- ions.

11. Strengths of Conjugate Acid-Base Pairs
The stronger the acid, the weaker its conjugate base; the stronger the base, the weaker its conjugate acid.

12. Autoionization of Water
Water naturally dissociates into H+ and OH- ions. Equilibrium is established between molecular water and the separated ions.
H2O H+ + OH-
Equilibrium constant for this reaction is Kw
Kw = 1.0 x = [H+][OH-]

13. Acid Dissociation Constant (Ka)
HA(aq) + H2O(l) H3O+(aq) + A-(aq)
Ka is a measure of the strength of an acid
The greater the Ka, the stronger the acid

14. Base Dissociation Constant (Kb)
B(aq) + H2O(l) HB+(aq) + OH-(aq)
Measure of the strength of a base

15. Converting from Ka to Kb
Ka = Kw/Kb
Kb=Kw/Ka

16. pH Scale Logarithmic scale
Each 1 unit increment = a tenfold change in concentration of H+ ions
pH = -log [H+]
pOH = -log [OH-]
pH = 14 – pOH
pOH = 14 – pH

17. What is the pH of a solution of 0.04 M HNO3?

18. What is the pH of a 0.89 molar solution of hydrocyanic acid?

19. What is the pH of a solution of 0.45 M ammonia?

20. What was the initial concentration of ammonia if the pH of the solution at equilibrium is 9.6?

21. Neutralization Reactions
Acid + Base  Water + Salt (Ionic Compound)

22. Titration
Determining the concentration of a solution by reacting it with known concentration of an opposite solution until an equivalence point is reached.
Equivalence point is when amount of H+ in solution equals amount of OH-.

23. Indicators
Chemicals that change color in the presence of an acid or base.
Ex: phenolphthalein
Bromothymol blue
Methyl orange
Purple cabbage juice
Etc.

24. 50 mL of hydrochloric acid required 39 mL of 1
50 mL of hydrochloric acid required 39 mL of 1.2 M sodium hydroxide to titrate it to the equivalence point. What is the concentration of the hydrochloric acid solution?

25. 100 mL of a solution of hydrofluoric acid was titrated to equivalence with 3.5mL of M sodium hydroxide. What was the initial concentration of the hydrofluoric acid?