1. Chemistry 100
Chapter 14
Acids and Bases

2. Acids and Bases
Acids: sour
Bases: bitter or salty

3. Arrhenius definition:
Acids and Bases
Arrhenius definition:
(If H2O is involved.)
Acid: produces H3O+
CH3COOH(aq) + H2O(l) CH3COO-(aq) + H3O+(aq)
H3O+ (Hydronium ion):
H+(aq) + H2O(l) H3O+(aq)
Base: produces OH-
H2O
NaOH(s) Na+(aq) + OH-(aq)
NH3(aq) + H2O(l) NH4+(aq) + OH-(aq)

4. Brønsted and Lowry definition: (If H2O is not involved.)
Acids and Bases
Brønsted and Lowry definition: (If H2O is not involved.)
Acid: donates H+ (proton) a proton donor
Base: accepts H+ (proton) a proton acceptor
HCl H2O Cl H3O+
acid
base
Conjugate
base
Conjugate
acid
Conjugate acid-base pair
Conjugate acid-base pair

5. Acids and Bases
HCl H2O Cl H3O+
Proton (H+) is transferred.

6. Acids and Bases CH3COOH + NH3 CH3COO- + NH4+
Conjugate
Conjugate acid-base pair
C6H5OH + H2O C6H5O H3O+
acid
base
Conjugate
base
Conjugate
acid
Conjugate acid-base pair
Conjugate acid-base pair

7. Weak acid and base: is partially ionized in aqueous solution.
Acids and Bases
Weak acid and base: is partially ionized in aqueous solution.
produces less H+ and OH-
CH3COOH(aq) + H2O(l) CH3COO-(aq) + H3O+(aq)
NH3(aq) + H2O(l) NH4+(aq) + OH-(aq)
Strong acid and base: is completely ionized in aqueous solution.
produces more H+ and OH-
HCl(aq) + H2O(l) Cl-(aq) + H3O+(aq)
NaOH(aq) + H2O(l) Na+(aq) + OH-(aq)

8. Electrolytes + - Electrolyte: conducts an electric current.
Na+
Cl-
bulb
Electrolyte: conducts an electric current.
Ionization (Dissociation)
NaCl → Na+ + Cl-
strong electrolytes: molecules dissociate completely into ions (NaCl).
weak electrolytes: molecules dissociate partially into ions (CH3COOH).
nonelectrolytes: molecules do not dissociate into ions (DI water).

9. Acids and Bases Strong acid/base  Strong electrolyte
Weak acid/base  Weak electrolyte
HCl(aq) + H2O(l) Cl-(aq) + H3O+(aq)

10. Acids and Bases
A strong acid contains a weak conjugate base.

11. Acids and Bases Monoprotic acids HCl Diprotic acids H2SO4
Triprotic acids
H3PO4
Amphiprotic: it can act as either an acid or a base.
HCl(aq) + H2O(l) Cl-(aq) + H3O+(aq)
NaOH(aq) + H2O(l) Na+(aq) + OH-(aq)
base
acid

12. Organic acids: contain carboxyl group (-COOH).
Acids and Bases
Oxyacids: acidic H is attached to an oxygen atom.
H2SO4 H3PO4 HNO3
Organic acids: contain carboxyl group (-COOH).
They are usually weak.
CH3COOH

13. Naming binary acids Hydro -ide ion -ic acid Anion : + HF
F-: flouride ion Hydroflouric acid
HCl
Cl-: chloride ion Hydrochloric acid
H2S
S2-: sulfuride ion Hydrosulfuric acid

14. Naming ternary acids -ite ion -ous acid Anion: -ate ion -ic acid HNO2
NO2-: Nitrite ion Nitrous acid
HNO3
NO3-: Nitrate ion Nitric acid
H2CO3
CO32-: carbonate ion carbonic acid
H2SO3
SO32-: sulfurite ion sulfurous acid

15. Ionization constant HA + H2O A- + H3O+ [A-] [H3O+] K = [HA] [H2O]
Equilibrium constant
K =
not for strong acids
[HA] [H2O]
Ka = K [H2O] =
[A-] [H3O+]
[HA]
Acid ionization constant
Ka < 1
- Log Ka = pKa
Ka ↑ or pKa ↓
Stronger acid

16. Ionization of water H2O(l) + H2O(l) ⇌ H3O+ (aq) + OH- (aq)
KW = [H3O+] [OH-] = (1×10-7) (1×10-7)
[H3O+] [OH-] = 1×10-14
pH + pOH = 14

17. [H+] and [OH-] [H+] = [OH-] Neutral solution
[H+] > [OH-] Acidic solution
[H+] < [OH-] Basic solution

18. pH and pOH pH = - log [H3O+] or -log [H+] pOH = - log [OH-] pH scale:7 14
Acid
Neutral
Base
[H3O+] ↑ and [OH-] ↓
[H+] = [OH-]
[H3O+] ↓ and [OH-] ↑

19. pH meter and pH indicators

20. Nature & pH indicators Bigleaf Hydrangea In basic soil (alkaline)
In acidic soil

22. pH of strong acids 0.10 M HCl  pH = ? [H3O+] = [H+] = 0.10 M
HCl(aq) + H2O(l) Cl-(aq) + H3O+(aq)
1 mol
1 mol
1 mol
0.10 M
0.10 M
0.10 M
[H3O+] = [H+] = 0.10 M
pH = -log [H+]
pH = -log (0.10) = 1.00

23. Acid Base Reactions
Neutralization: reaction between an acid and a base.
Acid + Base Salt + Water
KOH(aq) + 2HCl(aq) KCl(aq) + H2O(l)
2NaOH(aq) + H2SO4(aq) Na2SO4(aq) + 2H2O(l)
Strong acid reacts with strong base
to produce the weaker acid and weaker base.
(This is the direction of a reaction)

24. Titration (Neutralization reaction)B A
MB: known
VB: known
MA: unknown
VA: known
Equivalence point:
Equal amount of acid (H+) and base (OH-) (pH = 7).
H2SO4(aq) + 2NaOH(aq)  Na2SO4(aq) + 2H2O(l)
Acid
Base

25. Practice: Titration (Neutralization reaction)
H2SO4(aq) + 2NaOH(aq)  Na2SO4(aq) + 2H2O(l)
Find the concentration of NaOH solution if 143 mL of this solution is completely neutralized
by L of M H2SO4 solution?
0.150 mol H2SO4
1L H2SO4 solution
1 mol H2SO4
2 mol NaOH ×
0.205 L H2SO4 solution ×
= mol NaOH solution M
mol NaOH
M =
M =
= M
V (L)
0.143 L NaOH

26. Buffers pH stays constant. Buffer
Acid or Base
pH stays constant.
Buffer
A buffer resists changes in pH when limited
amounts of acid or base are added.

27. Buffers Our blood is a buffer solution. pH of blood ≈ 7.4
Acid
Acid
pH of blood ≈ 7.4
Base
Shock Absorber
Base

28. Weak Acid + its Conjugate base (in equilibrium)
Buffer Composition
Weak Acid + its Conjugate base (in equilibrium)
salt of the weak acid
CH3COOH + CH3COO-Na+
CH3COOH / CH3COO-
Or it can be weak base with it’s conjugate acid.

29. Buffers pH of blood = between 7.35 and 7.45
Carbonate buffer H2CO3 / HCO3-
Phosphate buffer H2PO4- / HPO42-
Proteins buffer

30. How do buffers work? HCO3- + H3O+ → H2CO3 + H2O
Carbonate buffer H2CO3 / HCO3-
If we eat an acidic food:
HCO3- + H3O+ → H2CO3 + H2O
H2CO3 + OH- → HCO3- + H2O
If we eat a basic food:

31. Henderson-Hasselbalch equation
pH of Buffers
HA(aq) A-(aq) H+(aq)
Weak acid
Conjugate base
[Conjugate Base]
pH = pKa + log
[Weak Acid]
Henderson-Hasselbalch equation
[Weak Acid]: concentration of the weak acid
[Conjugate Base]: concentration of its conjugate base
pKa of the weak acid