2. • One system which is convenient at an
introductory level involves classifying
reactions into four basic types.
• Knowing the different types of reaction
allows the chemist to predict the products
of the reactions of different substance.c

3. TYPES OF REACTIONS:
1.COMBINATION REACTION
2.DECOMPOSITON REACTION
3.SUBSTITUTION REACTION
4.DOUBLE REPLACEMENT REACTION

4. 1. COMBINATION REACTION
• Type of reaction in which two
or more substances combine
to form a more complex
substance.
• Also called synthesis reaction
• Form:
A + B -> AB

5. EXAMPLES:
• C ₍s₎ + O₂₍g₎ -> CO ₂₍g₎
• 2 Ca ₍s₎ + O ₂₍g₎ -> 2CaO ₍s₎
• N ₂₍g₎ + 3 H ₂₍g₎ -> 2NH ₃₍g₎
• 2 Na ₍s₎ + Cl ₂₍g₎ -> 2NaCl ₍s₎

6. 2. DECOMPOSITION REACTION
Various conditions can cause a substance to
decompose:
• The most common type of a decomposition is
cause by a rise in temperature. This is called
thermal decomposition.
Example:
2 HgO ₍s₎ ->(heat) 2Hg ₍g₎ + O ₂₍g₎
2KMnO ₄₍s₎ ->(heat) K ₂MnO ₄₍s₎ + MnO ₂₍g₎ + O ₂₍g₎

7. • If such a decomposition proceeds
very fast, an explosion may occur
Example:
2NH₄NO ₃₍s₎ ->(heat) 4H₂O ₍g₎ +
2N ₂₍g₎ + O ₂₍g₎

8. • Light is another factor which may
occur chemical decompositions to
occur. Such decompositions are
known as photo-chemical
decompositions.
Example:
2AgBr ₍s₎ ->(light) 2Ag ₍s₎ + Br ₂₍g₎

9. • Although catalysts cannot cause chemicals to decompose,
they can speed up the rate of a chemical decomposition.
The addition of a tiny amount of manganese dioxide to a
solution of hydrogen peroxide will immediately result in a
vigorous evolution of oxygen gas.
Example:
2H ₂O ₂₍aq₎ ->(manganese dioxide/catalyst)2 H ₂O ₍l₎ + O ₂₍g₎

10. • Electricity can also be used to decompose
compounds, resulting in electrolytic
decomposition.
Examples:
2NaCl ->(electrolysis) 2Na + Cl ₂
PbBr ₂ ->(electrolysis) Pb + Br ₂

11. 3. SUBSTITUTION REACTION
• Also called single replacement reaction
• Form:
C + AB -> CB + A
• more active element replaces a less active element in
the compound.
Example:
Fe + CuSO₄ -> FeSO₄ + Cu
2 Na + 2H₂O -> 2NaOH + H₂

12. Metals:
Lithium(Li) Chromium(Cr) Mercury(Hg)
Potassium (K) Iron(Fe) Platinum(Pt)
Barium(Ba) Cobalt(Co) Gold(Au)
Calcium(Ca) Nickel(Ni)
Sodium(Na) Tin(Sn)
Magnesium(Mg) Lead(Pb)
Aluminum(Al) Hydrogen(H₂)
Manganese(Mn) Copper(Cu)
Zinc(Zn) Silver(Ag)

13. Halogens:
Flourine(F)
Chlorine(Cl)
Bromine(Br)
Iodine(I)

14. 4. DOUBLE REPLACEMENT REACTION
• Also called ionic reaction
• Characterized by the exchange of ions
(charged atoms or molecules) between
two compounds
• Form:
AB + CD -> AD + CB

15. Several types of double replacement reactions are
frequently encountered:
1. A reaction between a base such as NaOH, KOH or
Ca(OH)₂ in aqueous solution and an acid such as HCl,
H₂SO₄ or CH₃COOH. Reactions of this type are called
neutralizations and the products are always water
molecules and salt.
HCl + NaOH -> NaCl + H₂O
H₂SO₄ + 2KOH -> K₂SO₄ + 2H₂O

16. 2. A reaction between solutions of two soluble
salts, a soluble salt and an acid, or a soluble
salt and an alkali which, by exchanging cations
and anions, can produce one insoluble
compound.
Examples:
NaCl + AgNO₃ -> NaNO₃ + AgCl
BaCL₂ + MgSO₄ -> MgCl₂ + BaSO₄

17. Activity:
Given the type of reaction, predict the
products and balance the equation.
1. Combination: Li + F₂
2. Decomposition: CuCO₃
3. Substitution: Mg + HCl
4. Ionic reaction: AgNO₃ + BaCl₂

18. End …
Thank you!