Topic 8: Acids and Bases
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Theories of acids and bases
Properties of acids and bases
Strong and weak acids and bases
The pH-scale
8.1 Brönsted-Lowry
Acid - base Theory
• The protolytic reaction:
Reaction of an acid with water:
HCl (g) + H2O ↔ H3O+ + Cl-
Reaction of a base with water:
NH3 (g) + H2O ↔ OH- + NH4+
• Draw the reaction formulas when these
substances react with water as acids:
HNO3
H3PO4
H2SO4
CH3COOH
• Draw the reaction formulas when these
substances react with water as bases:
CO32CH3NH2
Brönsted-Lowry
Acid - base Theory
• Acid: proton donor
• Base: proton acceptor
Conjugate acid-base pair
CH3COOH + H2O  CH3COO- + H3O+
Acid1
Base2
Base1
Acid2
State for each reaction which reactant
is acid and which is base
a) HSO4- + H2O ↔ SO42- + H3O+
b) NH3 + H2O ↔ NH4+ + OHc) HCO3- + H2O ↔ CO32- + H3O+
d) HCO3- + H2O ↔ H2CO3 + OHe) H3O+ + OH- ↔ 2 H2O
State for each reaction which reactant
is acid and which is base
a) HSO4- + H2O ↔ SO42- + H3O+
acid
base
b) NH3 + H2O ↔ NH4+ + OHbase acid
c) HCO3- + H2O ↔ CO32- + H3O+
acid base
d) HCO3- + H2O ↔ H2CO3 + OHbase acid
e) H3O+ + OH- ↔ 2 H2O
base acid
Which of these are conjugated
acid/base-pairs?
a) HSO4-/SO42b) H2SO4-/SO42c) NH3 /NH4+
d) HCO3-/CO32e) CO32-/H2CO3
f) H3O+/OH-
Which of these are conjugated
acid/base-pairs?
a) HSO4-/SO42c) NH3 /NH4+
d) HCO3-/CO32What is the rule?
The carboxyl group
Amphiprotic
• Water can act both as an acid and as a base;
H3O+  H2O  OH-
• Such compounds are said to be amphiprotic
(ampholytic).
Monoprotic  Polyprotic
• Monoprotic: CH3COOH  CH3COOAcetic acid
• Diprotic:
HOOC-COOH  -OOC-COO-
• Triprotic:
H3PO4  PO43-
Oxalic acid
Phosphoric acid
• Polyprotic
Brönsted-Lowry
Acid - base Theory
• Acid: proton donor
• Base: proton acceptor
Arrhenius
Acid-Base Theory
• Acid: H+ Hydrogen ion / Proton
– Acidic solutions contain H+ / H3O+
(oxonium,
hydroxonium or
hydronium ion)
• Base: OH- Hydroxide ion
– Alkaline solutions contain OHAlkaline = Water soluble base
Lewis
Acid-Base Theory
• Lewis acid: electron pair acceptor,
– e.g. H+, AlCl3, BF3
• Lewis base: electron pair donor,
– e.g. OH-, NH3
• A Lewis acid-base reaction involves the
formation of a covalent bond. The Lewis base
provides the electrons in that bond. This kind
of covalent bond is called dative covalent
bonds (see topic 13) or co-ordinate covalent
bond.
• Its no difference between a normal covalent
bond and a dative covalent bond except the
origin of the electrons. Sometimes an arrow is
used instead of a line to show that it's a dative
bond e.g. H3NBF3.
• The term Lewis acid is often just used for acids
that aren’t Brönstedt acids
• The formation of complex ions, topic 13, is
usually Lewis acid-base reactions
Exercises page 144-2
8.2 Properties of acids in solution
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They have pH<7
They taste sour
They react with bases and metals
Where can you find:
Hydrochloric acid, HCl
Sulphuric acid, H2SO4
Acetic acid, CH3COOH
Carbonic acid, H2CO3
Properties of bases in solution
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They have pH>7
They feel ”slippery”
They react with acids
Where can you find:
Ammonia, NH3
Sodium bicarbonate, NaHCO3
Calcium carbonate, CaCO3
8.3 Strong acids
Totally dissociated
• Hydrochloric acid: HCl + H2O  Cl- + H3O+
chloride
• Nitric acid:
HNO3+ H2O  NO3- + H3O+
nitrate
• Sulphuric acid: H2SO4 + H2O  SO42- + 2 H3O+
sulphate
Start
End
HCl + H2O  H3O+ +Cl100%
0%
0%
100%
Weak acids
Partially dissociated
• Ethanoic acid, (Acetic acid)
CH3COOH + H2O  CH3COO- + H3O+
ethanoate ion
(acetate ion)
• Carbonic acid,
H2CO3 H2O 
CO32- + H3O+
carbonate ion
(HCO3- hydrogen carbonate)
CH3CH2COOH + H2O  H3O+ + CH3CH2COOstart 100%
0%
end 99%
1%
If the concentration is the same for the strong
and the weak acid:
– The strong acid is more acidic than the weak acid
– The strong acid has a higher concentration of
hydroxonium ions than the weak acid
– The strong acid has higher conductivity
Strong bases
Containing the OH- ion
• All group I hydroxides:
NaOH(s) + H2O Na+ + OH• Group II hydroxides
Ba(OH)2 + H2O Ba2+ + 2 OH-
Weak bases
Partially dissociated
• Ammonia
NH3 + H2O  NH4+ + OH• Ethylamine
CH3CH2-NH2 + H2O  CH3CH2-NH3+ + OH-
The anions from carbonic acid;
CO32- and HCO3• Alkaline properties
• Often water soluble salts
• H2CO3 + H2O HCO3- + H3O+
Acid
Base
Base
Acid
• HCO3- + H2O  CO32- + H3O+
Acid
Base
Base
Conjugated acid and base pair
Acid
Conjugated acid and base pair
Indicators
Acidic
Neutral
Basic
Litmus
red
blue
BTB
(red) yellow
green
Blue
Phenolphthalein
colourless
colourless
cerise
Universal paper
(red)
(green)
(blue)
Some typical reactions of acids: salt
formations
• Neutralisation
• Reactions with metals or metal oxides
Neutralisation
Acid + base  salt + water
HCl + NaOH  NaCl + H2O
H2SO4 + KOH 
?
?
Draw reaction formulas
HCl + Ba(OH)2 
H2SO4 + LiOH 
HNO3 + Mg(OH)2 
Draw reaction formulas- facit
2 HCl + Ba(OH)2  2 H2O + BaCl2
H2SO4 + 2 LiOH  2 H2O + Li2SO4
2 HNO3 + Mg(OH)2  2 H2O + Mg(NO3)2
With metals
Acid + metal  salt + hydrogen gas
Mg + 2 HCl  MgCl2 + H2
Al + H2SO4 
?
?
Noble metals (Cu, Ag, Au) doesn’t react with HCl or H2SO4 . They
demand more oxidative acids (HNO3) and will then give other
gases than H2 (N2O)
Draw reaction formulas
Ca + HCl 
Al + HCl 
Na + CH3COOH 
Draw reaction formulas- facit
Ca + 2 HCl  CaCl2 + H2
2 Al + 6 HCl 2 AlCl3 + 3 H2
2 Na + 2 CH3COOH  NaCH3COO + H2
With metal oxides
Acid + metal oxide  salt + water
CuO + 2 HCl  CuCl2 + H2O
To synthesise a salt from a noble metal you can’t start with
metal + acid (Why?)
Carbonates and hydrogen carbonates
Carbonates + acids salt+carbon dioxide+water
Na2CO3 + 2 HCl  2 NaCl + CO2 + H2O
K2CO3 + H2SO4 
CaCO3 + HCl 
Carbonates and hydrogen carbonates
Carbonates + acids salt+carbon dioxide+water
Na2CO3 + 2 HCl  2 NaCl + CO2 + H2O
K2CO3 + H2SO4  K2SO4 + CO2 + H2O
CaCO3 + HCl  CaCl2 + CO2 + H2O
8.4 The pH-scale
pH = -log[H+]
[H+] = [H3O+]= 10-pH
• pH = -log[H+] => change in one pH unit = 10
times difference in [H+]
• pH=5  pH= 3 => 100 times more acidic.
• pH=8  pH= 11 => 1000 times more basic.
• pH-meter, pH-paper
• [H+] = 10-pH
Exercises 1-2 page 148