Acids, Bases, & Salts
I. Properties of Acids & Bases
A. Properties of Acids
1. Aqueous solutions have a sour taste
2. Acids change the color of acid-base indicators
3. Some acids react with active metals to release
hydrogen
Zn(s) + H2SO4(aq)  ZnSO4(aq) + H2(g)
4. Acids react with bases to produce salts and water
HCl(aq) + NaOH(aq)  NaCl(aq) + H2O(l)
5. Acids conduct electric current
B. Properties of Bases
1.
2.
3.
4.
5.
Aqueous solutions of bases have a bitter taste
Bases change the color of acid-base indicators
Dilute aqueous solutions of bases feel slippery
Bases react with acids to produce salts and water
Bases conduct electric current
II. Arrhenius Acids and Bases - Svante
Arrhenius, Swedish chemist (1859-1927)
A. Arrhenius Acid – A chemical compound that
increases the concentration of hydrogen ions, H+,
in aqueous solution
[H+] > [OH-]
B. Arrhenius Base – A substance that increases
the concentration of hydroxide ions, OH-, in
aqueous solution
[H+] < [OH-]
C. Aqueous Solutions of Acids
1. Acids are molecular compounds that ionize in
solution
HNO3 + H2O  H3O+ + NO3H2SO4 + H2O  H3O+ + HSO4H2O + HCl  H3O+ + Cl-
D. Strength of Acids
1. Strong acids ionize completely in solution
2. Weak acids ionize only slightly and are weak
electrolytes
Strong Acids
Weak Acids
H2SO4
HSO4-
HClO4
H3PO4
HCl
HF
HNO3
CH3COOH
HBr
H2CO3
HI
H2S
HCN
HCO3-
E. Aqueous Solutions of Bases
1. Ionic bases dissociate to some extent when
placed in water
NaOH (s)
Na+ (aq) + OH − (aq)
1. Basic solutions are referred to as “alkaline”
2. Molecular bases react with water to produce
hydroxide ions
NH3 (g) + H2O (l) ↔ NH4+ (aq) + OH- (aq)
F. Strength of Bases
1. Strength of ionic bases is linked to solubility
a) High solubility = strong base
b) Low solubility = weak base
2. Molecular bases tend to be weak regardless of
solubility
III. Bronsted-Lowry Acids & Bases
A. Bronsted-Lowry Acid – molecule or ion
that donates a proton (H+)
B. Bronsted-Lowry Base – molecule or ion
that accepts protons
1. Hydroxide ions (OH-) are acceptor of ionic
bases
IV. Conjugate Acids & Bases
A. Conjugate Base
1. The species that remains after an acid
has given up a proton
H3PO4 (aq) + H2O (l) ↔ H3O+ (aq) + H2PO4-(aq)
acid
conjugate base
2. The stronger an acid, the weaker its
conjugate base
B. Conjugate Acid
1. The species that is formed when a base
gains a proton
H3PO4 (aq) + H2O (l) ↔ H3O+ (aq) + H2PO4- (aq)
base
conjugate acid
2. The stronger a base, the weaker its
conjugate acid
IV. Ionization Constants
A. Ionization constant = Ka
B. Compares to relative
strength of acids
(high Ka = stronger acid)
B. pH and pOH
1. H2O(l) + H2O(l) ↔ H3O+ + OH2. At 25oC, [H3O+] and [OH-] = 1.0 x 10-7 mol/L,
and remains constant in pure water and
dilute aqueous solutions.
3. This constant, Kw, is called the ionization
constant of water.
Kw = [H3O+][OH-] = (1.0 x 10-7)(1.0 x 10-7) = 1.0 x 10-14 mol/L
V. pH Scale & [H3O+] [OH-]
A. pH Scale
1. pH < 7 : acid
[H3O+] > [OH-]
2. pH = 7 : neutral
[H3O+] = [OH-]
3. pH > 7 : basic/alkaline [H3O+] < [OH-]
B. Formulas for pH, pOH, [H3O+], & [OH-]
pH = - log [H3O+] pOH = - log [OH-] pH + pOH = 14
Calculating pH:
[H3O+] must be in scientific notation!
Calculate the log of the [H3O+]
Multiply by -1
If [H3O+] = 0.0261, then pH =
0.0261 = 2.61 x 10-2
(log 2.61 x 10-2) x -1 = 1.58
(In calculators type: 2.61, EE, 2, +/-, log, +/-)
Calculating pOH
[OH-] must be in scientific notation!
Calculate the log of the [OH-]
Multiply by -1
If [OH-] = 0.0000000048, then pOH =
0.0000000048 = 4.8 x 10-9
log 4.8 x 10-9 x-1 = 8.3
(In calculators type: 4.8, EE, 9, +/-, log, +/-)
Calculating [H3O+] and [OH-]
• From pH & [OH-] from [H3O+] (and vice versa)
[H3O+] = antilog (-pH)
If pH = 7.52, what are the [H3O+] and [OH-]
concentrations?
a) [H3O+] = antilog (-7.52)
(In calculators type: 7.52, +/-, 2nd, 10x or 7.52, +/-, 2nd,
log)
b) [H3O+] =
c) [H3O+] [OH-] = 1.0 x 10-14
d) [OH-] = 1.0 x 10-14 = 1.0 x 10-14 = 3.3 x 10-7 M OH[H3O+]