1. Types of Reactions

2. Types of Reactions:
Chemical reactions can be classified as:
1) Synthesis or combination reactions
2) Decomposition reactions
3) Single replacement reactions
4) Double replacement reactions
5) Combustion reactions

3. 1) Synthesis or Combination Reactions:
Two or more elements form only one product.

4. 1) Synthesis or Combination Reactions:
Two or more elements form only one product.
2Mg(s) O2(g) MgO(s)
2Na(s) Cl2(g) NaCl(s)
SO3(g) H2O(l) H2SO4(aq)

5. 2) Decomposition Reactions:
One substance splits into two or more simpler
substances.
2HgO(s) Hg(l) + O2(g)
2KClO3(s) KCl(s) + 3O2(g)

6. 3) Single Replacement Reactions:
One element takes the place of a different
element in another reacting compound.

7. 3) Single Replacement Reactions:
In a single replacement reaction, one element
takes the place of a different element in
another reacting compound.
Zn(s) + 2HCl(aq) ZnCl2(aq) + H2(g)
Fe(s) + CuSO4(aq) FeSO4(aq) + Cu(s)

8. 4) Double Replacement Reactions:
Two elements in the reactants exchange places.

9. 4) Double Replacement Reactions:
Two elements in the reactants exchange places.
AgNO3(aq) + NaCl(aq) AgCl(s) + NaNO3(aq)
ZnS(s) HCl(aq) ZnCl2(aq) + H2S(g)

10. Double replacement reactions in which
acids and bases react to produce a salt
and water are also classified as Neutralization reactions.
Examples:
HCl(aq) + NaOH(aq) → NaCl(aq) + H2O(l)
acid base salt water
H2SO4(aq) + 2KOH(aq) → K2SO4(aq) + 2H2O(l)
acid base salt water

11. 5) Combustion Reactions:
a carbon-containing compound burns in oxygen gas to form carbon dioxide (CO2) and water (H2O)
energy is released as a product in the form of heat
CH4(g) + 2O2(g) CO2(g) + 2H2O(g) + energy

12. Learning Check: Classify each of the following reactions:
A. 2Al(s) + 3H2SO4(aq) Al2(SO4)3(s) + 3H2(g)
B. Na2SO4(aq) + 2AgNO3(aq) Ag2SO4(s) + 2NaNO3(aq)
C. N2(g) + O2(g) NO(g)
D. C2H4(g) + 2O2(g) CO2(g) + 2H2O(g)
Single Replacement
Double Replacement
Synthesis or Combination
Combustion

13. E. 2Ag(s) + H2S(aq) Ag2S(s) + H2(g)
3Ba(s) + N2(g) Ba3N2(s)
E. 2Ag(s) + H2S(aq) Ag2S(s) + H2(g)
F. 2C2H6(g) + 7O2(g) CO2(g) + 6H2O(g)
G. PbCl2(aq) + K2SO4(aq) KCl(aq) + PbSO4(s)
H. K2CO3(s) K2O(aq) + CO2(g)
Synthesis or Combination
Single Replacement
Combustion
Double Replacement
Decomposition

14. Double Replacement & Neutralization.
Exercise: Classify the following reactions:
(1) 2HCl (aq) + Zn(s) → ZnCl2(aq) + H2(g)
(2) H2SO4(aq) + 2NaOH (aq) → Na2SO4 (aq) + 2H2O
(3) MgCl2(aq) + 2Na(s) → 2 NaCl (aq) + Mg(s)
(4) Al(NO3)2(aq) + 3H2SO4(aq)→ Al2(SO4)3(aq) +6HNO3(aq)
Single Replacement.
Double Replacement & Neutralization.
Single Replacement.
Double Replacement.

15. (5) CH4(g) + 2O2(g) → CO2 (g) + 2H2O(g)
Combustion.
Combustion.

16. Thermal energy (heat) is released during the reaction.
A chemical Reaction can be classify also as:
Exothermic
Endothermic
Thermal energy (heat) is released during the reaction.
Thermal energy (heat) is absorbed during the reaction.

17. Exothermic Reactions:
Examples:
Exothermic Reactions:
4 Fe(s) + 3 O2 (g) Fe2 O3(s) + Q
Heat
4 Fe(s) + 3 O2 (g) Fe2 O3(s) + 1,625 kJ
4 Fe(s) + 3 O2 (g) Fe2 O3(s) ∆H= - 1,625 kJ

18. Endothermic Reactions:
Examples:
Endothermic Reactions:
2SO3 (g) Q SO2(g) + O2(g)
Heat
2SO3 (g) kJ SO2(g) + O2(g)
2SO3 (g) SO2(g) + O2(g) ∆H= 198 kJ

19. ∆H › 0
Endothermic process
∆H ‹ 0
Exothermic process