Types of Reactions

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

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

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

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

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

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)

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

4) Double Replacement Reactions: Two elements in the reactants exchange places. AgNO3(aq) + NaCl(aq) AgCl(s) + NaNO3(aq) ZnS(s) + 2HCl(aq) ZnCl2(aq) + H2S(g)

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

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

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) 2NO(g) D. C2H4(g) + 2O2(g) 2CO2(g) + 2H2O(g) Single Replacement Double Replacement Synthesis or Combination Combustion

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) 4CO2(g) + 6H2O(g) G. PbCl2(aq) + K2SO4(aq) 2KCl(aq) + PbSO4(s) H. K2CO3(s) K2O(aq) + CO2(g) Synthesis or Combination Single Replacement Combustion Double Replacement Decomposition

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) 2Al(NO3)2(aq) + 3H2SO4(aq)→ Al2(SO4)3(aq) +6HNO3(aq) Single Replacement. Double Replacement & Neutralization. Single Replacement. Double Replacement.

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

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.

Exothermic Reactions: Examples: Exothermic Reactions: 4 Fe(s) + 3 O2 (g) 2 Fe2 O3(s) + Q Heat 4 Fe(s) + 3 O2 (g) 2 Fe2 O3(s) + 1,625 kJ 4 Fe(s) + 3 O2 (g) 2 Fe2 O3(s) ∆H= - 1,625 kJ

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

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