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CHEMICAL EQUATIONS
An equal amount of matter exists both before and after the experiment.
CH4 (g) + 2O2 (g) → CO2 (g) + 2H2O
(g)
Coefficients are inserted to balance the equation.
CH
4 (g) + 2 O2 (g) → CO2 (g) + 2 H2O (g)
• Subscripts tell the number of atoms of each element in a molecule
Subscripts and Coefficients Give Different Information
• Coefficients tell the number of molecules
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REACTION TYPES Some simple patterns of chemical reactivity
By understanding chemical reactions we:
Become better acquainted with chemical reactions and their balanced equations
Can predict products of some of these reactions knowing only their reactants
Recognize patterns of reactivities for a class of substances, thus giving a broader understanding than merely memorizing a large number of unrelated reactions
Let’s look at the following classes of reactions
1. Combination reactions
2. Decomposition reactions
3. Displacement reactions
4. Metathesis reactions
5. Neutralisation reactions
6. Oxidation-reduction (redox) reactions
7. Combustion reactions
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1. COMBINATION REACTIONS
General rxn: A + B → C
Examples:
(i) C(s) + O2(g) → CO2(g)
(ii) S(s) + O2(g) → SO2(g)
(iii) SO2 (s) + ½O2(g) → SO3 (g)
(iv) N2(g) + 3H2(g) → 2NH3(g)
(v) 2Na(s) + ½O2(g) → Na2O(s)
(vi) 2Na + O2(g) → Na2O2 (s)
(vii) NH3 (g) + HCl (g) → NH4Cl(s)
(viii) C3H6 (g) + Br2(l) → C3H6Br2(ℓ)
(viii) 2Mg(s) + O2(g) → 2MgO(s)
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Mg(s) + O2
Formation of Oxides
Element + Oxygen → oxide (O
→ 2MgO (s)
2-
e.g. Mg(s) + O
)
2 → 2MgO (s) – magnesium oxide
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There are 3 types of oxides :
a) Acidic oxides:
(i) non-metal + oxygen → acidic oxide
Examples: C(s) + O2(g) → CO2(g)
S(s) + O2(g) → 2SO2(g)
2S(s) + 3O2(g) → 2SO3(g)
Si(s) + O2(g) → SiO2(s)
(ii) acidic oxide + base → salt + H2O
CO2(g) + 2NaOH(aq) → Na2CO3(aq) + H2O(ℓ) (sodium carbonate –
Glass making)
SO3(g) + 2NaOH(aq) → Na2SO4 (aq) + H2O(ℓ)
(iii) acidic oxide + H2O → acid
Example: CO2(g) + H2O(l) → H2CO3(aq) (weak acid)
SO3(g) + H2O(l) → H2SO4(aq) (strong acid)
SiO2(s) + H2O(l) → insoluble
Silica is insoluble inmost acids except HF
Silica glass used for crucibles
All glass used for general purposes contain silica
Common glass: Na2O, CaO, 6SiO2
b) Basic oxides
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(i) metal + oxygen → basic oxides
Examples: 4Na(s) + O2(g) (limited) → 2Na2O(s)
4Na(s) + 2O2(g) (excess) → 2Na2O2(s)
2Ca(s) + O2(g) + heat → 2CaO(s)(calcined lime)
2Cu(s)+ O2(g) → 2CuO(s)
(ii) basic oxide + acid → salt + H2O
Example: Na2O(s) + 2HCl(aq) →2NaCl(aq) + H2O(ℓ)
CaO(s) + 2HCl(aq) →CaCl2(aq) + H2O(ℓ)
CuO(s) + 2HCl(aq) → CuCl2(aq) + H2O(ℓ)
(iii) basic oxide + H2O → base
Example: (i) Na2O(s) + H2O(ℓ) → 2NaOH(aq)
(ii) CaO(s) + 2H2O(ℓ) → Ca(OH)2(aq) - slaked lime
Ca(OH)2(aq) also called lime water used for CO2 (g) test
(iii) Ca(OH)2 (aq) + CO2(g) → CaCO3(s) + H2O(ℓ)
CaO(s) – quick lime – used as a drying agent in labs
(iv) Na2O(s) + H2O(ℓ) →2NaOH(aq)
(v) CuO(s) + H2O (ℓ) → insoluble
All group 1 oxides, Na2O, Li2O, K2O are soluble in H2
oxides that reacts with an acid and a base
O
Group II oxides are sparingly soluble in water
c) Amphoteric oxides
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Examples:
(i) Al2O3(s) + 6HCl(aq) (acid) → 2AlCl3(aq) + 3H2O(ℓ)
Al2O3(s) + 2NaOH(aq) (base) → 2Na[Al(OH)4 ](aq)
Al3+(aq) + 3OH- (aq) (limited) → Al(OH)3(s) white gelatinous precipitate
Al(OH)3 (s) + OH- (aq) → [Al(OH)4]- (aq) – complex ion
(i) ZnO(s) + 2HCl(aq) (acid) → ZnCl2(aq) + H2(g)
ZnO(s) + NaOH(aq) (base) → Zn(OH)4]2- (aq) – complex ion
2. DECOMPOSITION REACTIONS
General reaction: AB + heat → A + B
Examples:
(i) 2HgO(s) + heat → 2Hg(ℓ) + O2(g)
(ii) 2MgO(s) + heat → 2Mg(s) + O2(g)
(iii) 2KClO3(s) + heat → 2KCl(s) + 3O2(g)
(iv) 2KNO3(s) + heat → 2KNO2(s) + O2(g)
(v) 2Pb(NO3)2(s) + heat → 2PbO(s) + 2NO2(s) + O2(g)
(vi) 2H2O2(s) + heat → 2H2O (ℓ) + O2 (g) – Disproportionation reaction
(vii) CaCO3(s) + heat → 2CaO(s) + CO2(g)
(viii) Cu(OH)2(s) + heat → CuO(s) + H2O(g)
(ix) NH4NO3(s) + heat → 2N2O(g) (nitrous oxide ) + 2H2
(x) 2N
O(ℓ)
2O (s) + heat → 2N2(g) + O2(g)
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3. DISPLACEMENT REATIONS
General reaction: A + BC → AC + B
Examples:
A. Hydrogen Displacement
1. 2Na(s) + H2O(l) → 2NaOH(aq) + H2(g)
2. Ca(s) + H2O(l) → 2Ca(OH)2 (aq) + H2(g)
3. 2Al(s) + 3H2O(g) → 2Al2O3(s) + H2(g)
4. Zn(s) + 2HCl(ℓ) → ZnCl2 (aq) + H2(g)
5. Cd(s) + H2SO4 (dil) → 2CdSO4 (aq) + H2(g)
6. 2Ag(s) + HCl(dil) → no reaction
7. 2Cu(s) + HCl (dil) → no reaction
B. Metal Displacement or (Single Displacement rxns)
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1. Zn(s) + CuSO4(aq) → ZnSO4(aq) + Cu(s)
Nett reaction: Zno + Cu2+(aq) → Zn2+(aq) + Cuo
Note: Cuo + ZnSO4(aq) → no rxn – Why?
C. Halogen Displacement
1. Cl2(g) + 2KBr(aq) → 2KCl + Br2(ℓ)
Nett reaction: Cl2(g) + 2Br- (aq) → 2Cl-(aq) + Br2(ℓ)
2. Cl2(g) + 2I-(aq) → 2Cl-(aq) + I2(aq)
Order of reactivity of halogens: F2 > Cl2 > Br2 > I
Reaction with Acid
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Activity Series of the Elements
Element Reaction with H2O/Steam
React very vigorously with acid solutions to
produce H
Li
2
React with water to give H2 K
Ba Ca Na
React with acids to produce H
Mg
2
React slowly with water but readily with steam
Al Zn Cr Fe Cd Co Ni Sn Pb
Do not react with acids to give H(Cu, Hg, and Ag react with HNO
2
3
H
but do not produce
Cu Hg Ag
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H2 Pt ) Au
4. METATHESIS REACTION (Double Displacement Rxns)
Gen rxn: AB + CD ⇌ AB + BC
(a) Precipitate formation
(b) Weak electrolyte formation
(c) Gas Formation
(a) Precipitate formation
1. AgNO3(aq) + NaCl(aq) → AgCl(s) + NaNO3(aq)
Nett ionic rxn: Ag+ (aq) + Cl-(aq) → AgCl(s)
2. BaCl2 + H2SO4 → BaSO4(s) + 2HCl
Nett ionic rxn:
Ba2+(aq) + 2Cl-(aq) + 2H+(aq) + SO42-(aq) → BaSO4(s) + 2H+ +2Cl-(aq)
3. MgSO4(aq) + Na2CO3(aq) → MgCO3(s) + Na2SO4(aq)
A knowledge of solubility rules is needed here
Soluble Compounds Exceptions
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Compounds containing alkali ions (Li+, Na+, Cs+) and the ammonium ion (NH4+) Nitrates (NO3-), bicarbonates (HCO3-), and the chlorates (ClO3-
Halides (Cl
)
-, Br-, I-
Sulfates (SO
)
42-
Halides of Ag, Hg
)
22+, and Pb2+ Sulfates of Ag+, Ca2+, Sr+, Ba2+, Hg22+, and Pb2+
Insoluble compounds Exceptions Carbonates (CO32-), phosphates (PO43-), chromates (CrO42-), Sulfides (S2-) Hydroxides (OH-
Compounds containing alkali metal ions and the ammonium ion. Compounds containing alkali metal ions and the Ba
)
2+ ion.
Solubility rules for Common Ionic Compounds in water at 25oC (Chang, 8th Edition)
(b) Weak Electrolyte Formation
Acid + Base → Salt + water
Example. HCl + NaOH → NaCl + H2O – neutralization rxn
(i) H+(aq)(acid) + A-(aq)(conjugate base) → HA (weak acid) Example: H+(aq) + CH3COO-(aq) → CH3COOH
CH3COOH is a weak acid which is relatively undissociated or unionized.
(ii) OH-(aq)(base) + B+
Example: OH
(aq)(conjugate acid) → BOH (weak base)
-(aq) + NH4+(aq) → NH4OH ⇌ NH3 (aq) + H2O ℓ (ℓ)
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NH3
Strong electrolyte
is a weak base which is relatively undissociated or unionized
Definition of an electrolyte: a substance when dissolved in water can conduct electricity.
Classification of solutes in aqueous solution
Weak electrolyte Non-electrolyte HCl CH3 (NHCOOH 4)2CO - urea HNO HF 3 CH3OH - methanol HClO HNO4 CH2 3CH2OH - ethanol H2SO NH4 C3 6H12O6 - glucose NaOH H2 CO 12H22O11 - sucrose Ba(OH) 2
Examples of weak electrolytes
1. H2CO3 – carbonic acid
H+ (aq) + CO32-(aq) → HCO3-(aq) (unstable)
H+ (aq) + HCO3-(aq) → H2CO3(aq)
2. H2SO3 – sulfurous acid
H+(aq) + SO32-(aq) → HSO3-(aq) (unstable)
H+(aq) + HSO3-(aq) → H2SO3(aq) – sulfurous acid
H2SO3 + Heat → SO2(g) + H2O(ℓ)
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3. Na2CO3(aq) + 2HCl(aq) → H2CO3(aq) + 2NaCl(aq)
H2CO3(aq) + heat → CO2(g) + H2O(ℓ)
4. Na2SO3(aq) + H2SO4(aq) → H2SO3(aq) + 2NaSO4(aq)
H2SO3(aq) + Heat → SO2(g) + H2O(ℓ)
5. NH4Cl(aq) + NaNO2(aq) → NH4NO2(aq) + NaCl(aq)
NH4NO2(aq) → N2(g) + 2H2O(ℓ)
(c) Gas Formation
1. 2HCl(aq) + Na2S(aq) → H2S(g) + 2NaCl(aq)
2. NaCN(aq) + HCl(aq) → HCN(g) + NaCl(aq)
3. H2SO4(aq) + NaCl(aq) → NaHSO4(aq) + HCl (g)
Gases that are slightly soluble in water:
CO2(g) , SO2(g), HCN(g), NH3(g), H2S(g), HF(g)
5. Acid -Base (Neutralisation ) Reactions
(i) Strong acid-strong base reaction
HCl(aq) + NaOH(aq) → NaCl(aq)(neutral solvent) + H2O(ℓ)
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(ii) Strong acid-weak base reaction
HCl(aq) + NH3(aq) → NH4Cl(aq)
(iii) Weak acid-strong base
CH3COOH(aq) + NaOH(aq) → NaCH3COO(aq) + H2O(ℓ)
(iv) Weak acid-weak base
CH3COOH(aq) + NH3(aq) → NH4CH3COO(aq)
(v) Diprotic acid – base reaction
H2SO4(aq) + NaOH(aq) → NaHSO4(aq) + H2O(ℓ)
(vi) Triprotic acid-base reaction
H3PO4(aq) + NaOH(aq) → NaH2PO4(aq) + H2
• An oxidation occurs when an atom or ion loses electrons.
O(ℓ)
6. Oxidation-Reduction (Redox) Reactions
• A reduction occurs when an atom or ion gains electrons.
Examples:
1. Zn(s) + Cu2+(aq) → Zn2+(aq) + Cu(s)
2. 14H+ (aq) + Cr2O72-(aq) + 6Cl-(aq) → 2Cr3+(aq) + 7H2O (ℓ) + 3Cl2(g)
3. 2Cr(OH)3(s) + 6ClO-(aq) → 2 CrO42-(aq) + 3Cl2(g) + 2OH-(aq) + 2H2O(l)
7. Combustion reactions
Most combustions reactions produce a flame and involve O2 from air.
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Hydrocarbons, CxHyOz, are combusted in air, they react with O2 to form CO2(g) and H2O.
General reactions of alkanes:
CnH2n+2 (alkane) + 3n+12
O2 → nCO2 + (n+1) H2O
Examples:
1. CH4(g) + O2(g) (limited) → C(s) + 2H2O(g)
2. 2CH4(g) + 3O2(g) (restricted) → 2CO(g) + 2H2O(g)
3. CH4(g) + excess 2O2(g) → CO2(g) + 2H2O(g) – complete combustion
4. CH3OH(g) + O2(g) → CO2(g) + 2H2O(g)
All the above combustion reactions are oxidations reactions
