This document outlines four basic types of chemical reactions: combination reactions, decomposition reactions, substitution reactions, and double replacement (ionic) reactions. It provides examples of each type of reaction and describes their key characteristics, such as two or more substances combining in a combination reaction or one element replacing another in a substitution reaction. Knowing the different reaction types allows chemists to predict products from given reactants.

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!