Acid Reactions I
Acids & Active Metals:
• Single replacement reactions
• Active metals (K, Na, Ca, Mg, Al, Zn, Fe, Sn)
• Produce H2
• Oxidation-reduction reactions
Example: calcium solid + HCl
Zn and HCl
Ca (s) + 2 HCl (aq)  CaCl2 (aq) + H2 (g)
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Acid Reactions II
Acids, Carbonates and Bicarbonates
• H+ from acid transferred to carbonate
• Carbonate breaks down rapidly
• Products:
• CO2
• H2O
• salt (ionic compound)
Example:
NaHCO3 (aq) + HCl (aq)  H2CO3 (aq) + NaCl (aq)
CO2 (g) + H2O (l)
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Acid Reactions III
Neutralization
• Reactants: acid + base
• Products: water + salt
Example:
Molecular equation:
HCl (aq) + NaOH (aq)  HOH (l) + NaCl (aq)
Net ionic equation:
H+ (aq) + OH- (aq)  H2O (l)
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Acid-Base Titration
• Titration: Neutralize an unknown acid sample with a
known amount of base
• Indicator: changes color when pH changes
Example: phenolphthalein
Colorless
(acidic)
Add
base
Add
base
Pink
(basic)
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Acid-Base Properties of Salt Solutions
• Salts dissolve in water
• Which salts form…
• Neutral Solutions?
cation from strong base and anion from strong acid
NaNO3
NaNO3 (s)  Na+ (aq) + NO3 - (aq)
• Basic Solutions?
• Acidic Solutions?
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Basic and Acidic Salt Solutions
• Basic solutions form from salts that contain:
cation from strong base and anion from weak acid
NaF
NaF (s)


H 2O
Na+ (aq) + F - (aq)
• Acidic solutions form from salts that contain:
cation from weak base and anion from strong acid
NH4Cl
NH4Cl (s)


H 2O
NH4+ (aq) + Cl - (aq)
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Buffers
• Buffer solutions resist a change in pH
– Ex. Blood (pH ~ 7.4)
• Acid must neutralize small amounts of base
• Base must neutralize small amounts of acid
• Acid and base must not neutralize each other
Use conjugate acid-base pairs!
CH3COOH (aq) + H2O (l)
Added in as salt (NaCH3COO)
CH3COO - (aq) + H3O+ (aq)
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Buffers
• weak acid + its salt (containing conjugate base)
CH3COOH (aq) + H2O (l)
CH3COO - (aq) + H3O+ (aq)
You have prepared a buffer with 0.90 M CH3COOH and
1.2 M CH3COO-. (Ka for acetic acid = 1.8 x 10-5).
What is the equilibrium expression for this buffer?
What is the [H3O+] in this solution?
What is the pH of this solution?
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Making a buffer
Henderson-Hasselbalch equation
 [base] 

pH  pK a  log 
 [acid] 
We want to make 100. mL of a pH 4 citric
acid/sodium citrate buffer with an acid
concentration of 0.5 M.
We are given solid sodium citrate (294 g/mol) and
5.0 M citric acid. The pKa of citric acid is 3.15.
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