1 Types of Reactions Types of Reactions Chemical

1 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 2

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

1) Synthesis or Combination Reactions: Two or more elements form only one product. 2 Mg(s) + O 2(g) 2 Mg. O(s) 2 Na(s) + Cl 2(g) 2 Na. Cl(s) SO 3(g) + H 2 O(l) H 2 SO 4(aq) 4

2) Decomposition Reactions: One substance splits into two or more simpler substances.

2) Decomposition Reactions: One substance splits into two or more simpler substances. 2 Hg. O(s) 2 KCl. O 3(s) 2 Hg(l) + O 2(g) 2 KCl(s) + 3 O 2(g) 6

3) Single Replacement Reactions: One element takes the place of a different element in another reacting compound. 7

3) Single Replacement Reactions: One element takes the place of a different element in another reacting compound. Zn(s) + 2 HCl(aq) Fe(s) + Cu. SO 4(aq) Zn. Cl 2(aq) + H 2(g) Fe. SO 4(aq) + Cu(s) 8

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

4) Double Replacement Reactions: Two elements in the reactants exchange places. Ag. NO 3(aq) + Na. Cl(aq) Zn. S(s) + 2 HCl(aq) Ag. Cl(s) + Na. NO 3(aq) Zn. Cl 2(aq) + H 2 S(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) + Na. OH(aq) → Na. Cl(aq) + H 2 O(l) acid base salt water H 2 SO 4(aq) + 2 KOH(aq) → K 2 SO 4(aq) + 2 H 2 O(l) acid base salt water

5) Combustion Reactions: • a carbon-containing compound burns in oxygen gas to form carbon dioxide (CO 2) and water (H 2 O) • energy is released as a product in the form of heat CH 4(g) + 2 O 2(g) CO 2(g) + 2 H 2 O(g) + energy 12

Learning Check: Classify each of the following reactions: A. 2 Al(s) + 3 H 2 SO 4(aq) Al 2(SO 4)3(s) + 3 H 2(g) Single Replacement B. Na 2 SO 4(aq) + 2 Ag. NO 3(aq) C. N 2(g) + O 2(g) D. C 2 H 4(g) + 2 O 2(g) Ag 2 SO 4(s) + 2 Na. NO 3(aq) Double Replacement 2 NO(g) Synthesis or Combination 2 CO 2(g) + 2 H 2 O(g) Combustion E. H 2 SO 4(aq) + 2 Na. OH (aq) → Na 2 SO 4 (aq) + 2 H 2 O Double Replacement & Neutralization. 13

F. 3 Ba(s) + N 2(g) Ba 3 N 2(s) Synthesis or Combination G. 2 Ag(s) + H 2 S(aq) Ag 2 S(s) + H 2(g) Single Replacement H. 2 C 2 H 6(g) + 7 O 2(g) 4 CO 2(g) + 6 H 2 O(g) Combustion I. Pb. Cl 2(aq) + K 2 SO 4(aq) 2 KCl(aq) + Pb. SO 4(s) Double Replacement J. K 2 CO 3(s) K 2 O(aq) + CO 2(g) Decomposition 14

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.

Examples: Exothermic Reactions: 4 Fe(s) + 3 O 2 (g) 2 Fe 2 O 3(s) + Q Heat 2 Fe 2 O 3(s) + 1, 625 k. J 2 Fe 2 O 3(s) ∆H= - 1, 625 k. J

Examples: Endothermic Reactions: 2 SO 3 (g) + Q 2 SO 2(g) + O 2(g) Heat 2 SO 3 (g) + 198 k. J 2 SO 3 (g) 2 SO 2(g) + O 2(g) ∆H= 198 k. J

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