This document summarizes the key points about types of chemical reactions and factors that affect reaction rates. It discusses the 5 main types of chemical reactions: synthesis, decomposition, single-replacement, double-replacement, and oxidation-reduction. It also outlines characteristics of each type and provides examples. Additionally, the document covers concepts like activation energy, catalysts, temperature effects, and concentration effects on reaction rates. Finally, it briefly discusses chemical equilibrium and kinetic molecular theory of gases.

1. ChemiCal
ReaCtions
By
MUHAMMAD FAHAD ANSARI
12 IEEM 14

2. Types of Chemical Reactions
Most chemical reactions can be categorized into five
general types:
• Synthesis,
• Decomposition,
• Single-Replacement,
• Double-Replacement, and
• Oxidation-Reduction.

3. Characteristics of each type
• Synthesis: Two substances combine to form
one new substance:
In general: A + B → AB
For example:
2Na + Cl2 → 2NaCl
CaO + H2O → Ca(OH)2

4. • Decomposition: One substance breaks down to form
two new substances:
AB → A + B
For example: 2 H 2O → 2 H 2+ O 2
• Single Replacement: An element and a compound
react such that the element replaces one other
element in the compound:
A + BC → AC + B
For example: Mg + 2HCl → MgCl2 + H2

5. • Double Replacement: Two compounds react with each
other in such a way that they exchange partners with
each other:
AB + CD → AD + CB
For example:
NaBr + HCl → NaCl + HBr
• Oxidation-reduction: One or more elements in the
reaction changes its oxidation state during the reaction:
A 3+ → A 6+
For example: Cr 3+ → Cr 6+

6. Reaction Rates

7. Reaction Rates
Few questions related to factors that influence reaction
rates:
• Why substances such as gasoline be heated before
they will start burning?
• To what does the term catalytic refer in the catalytic
conversion process used in auto emission control
systems?
• Why do wood shavings burn more rapidly than
pieces of wood?
The answers to these and other questions involve the
rate of chemical reactions.

8. Terms
• Activation Energy - The difference in energy
between the reactants and the transition state that
is the energy barrier the reactants must overcome to
achieve a chemical reaction.
• Catalyst - A substance that lowers the activation
energy for a chemical reaction without being
chemically altered by the reaction.
• Homogeneous Catalyst - A catalyst that is in the
same phase as the reactants. While a heterogeneous
catalyst is not.

9. • Kinetics - The study of the rate and mechanism of chemical
reactions.
• Mechanism - The series of elementary steps that combine to
produce the path molecules take from reactant(s) to
product(s) in a chemical reaction.
• Rate - The speed of a reaction measured in amount or
reagent consumed or product produced per unit time.
• Reaction Coordinate Diagram - A plot of free energy versus
the reaction coordinate for a reaction that provides a pictorial
representation of the lowest energy path from reactants to
products.
• Transition State - The species with the highest energy
between reactants and products on a reaction coordinate
diagram, it is a short-lived species that represents a
combination of product-like and reactant-like properties.

10. Energy changes and chemical kinetics
• Chemical reactions are typically accompanied by
energy changes.
• The equation for the synthesis of ammonia from its
elements is,
N 2 + 3 H 2 → 2 NH 3
but that reaction takes place only under very special
conditions— at a high temperature and pressure and
in the presence of a catalyst.

11. Rate of Chemical Reaction
• The rate or speed with which a reactants disappears
or a product appears.
• The rate at which the concentration of one of the
reactants decreases or one of the products increases
with time.

12. For Example
• The decomposition of Dinitrogen Pentoxide
(N2O5) in an inert solvent carbon tetrachloride
(CCl4)
2 N2O5 (in CCl4) → 2N2O4 (in CCl4) + O2
• A typical unit for a rate of reaction
(mol per L per s)

13. Decomposition of N2O5 (in CCl4) at 45 0C [N2O5] = 1.40 M
decomposes completely (at time t = ∞, [N2O5] = 0 M)
Time, s Total volume O2, cm3
(at STP)
0 0
432 1.32
753 2.18
1116 2.89
1582 3.63
1986 4.10
2343 4.46
. .
. .
. .
∞a 5.93

14. Decomposition of Ozone, O3

15. Factors Affecting the Rate of Reactions
All the factors involved in the transformation of
reactants to products in chemical changes are called
chemical kinetics.
The major factors that affect the rate of reactions are:
• The nature of the reactants,
• Temperature,
• Catalysts, and
• The reactants concentration.

16. Nature of the reactants:
Different substances react at different rates.
• In general, the readiness with which any two
substances react together depends largely on the
structure of the atoms of the elements and the
nature of the bonds that hold the atoms together in
the substance.

17. Temperature:
• In general, an increase in temperature results in an
increase in rate of reaction.
Before reaction can occur between two substances,
the bonds already existing between the atoms of
each substance must be weakened and broken.
An increase in temperature results in greater kinetic
energy, (energy of motion), of the individual atoms
of each substance.

18. The greater energy of motion causes the atoms to
pull away from one another. As a result, the existing
bonds are weakened or broken.
The energy that must be added to weakened bonds
so that substances will react is called the energy of
activation.

19. Exothermic Reactions
Reactions that give off (evolve) heat are called
exothermic reactions.
A reaction will be exothermic if the energy of the
products is less than the energy of the reactants.
All combustion reactions are exothermic reactions such
as:
C + O2 → CO2 + heat
Mg + F2 → MgF2 + heat
N2 + 3H2 → 2NH3 + heat

20. Endothermic Reactions
Reactions that absorb heat are called endothermic reactions.
A reaction is endothermic when the energy of the products is
greater than the energy of the reactants.
For example:
2KClO3 + heat → 2KCl + 3O2
CaCO3 + heat → CaO + CO2
2NH3 + heat → N2 + 3H2
One reaction of equilibrium will evolve heat (exothermic) and
the other reaction will absorb heat (endothermic).
Exothermic
N2 + 3H2 -------→ 2NH3 + heat
Endothermic

21. Catalysts:
A catalyst is a substance that changes the rate of a reaction
without itself being used up in the process. Overall catalysts
effect is to change the reaction rate by lowering the energy of
reactants activation.
• There are two types of catalysts--heterogeneous catalysts and
homogeneous catalysts.
• The difference lies in whether the catalyst is in the same
phase (solid, liquid, or gas) as the reactants.
• A homogeneous catalyst is in the same phase as the reactants
while a heterogeneous catalyst is not. An enzyme is a
biological homogeneous catalyst.

23. Reactants Concentration:
In general, increasing the concentration of reactants increases
the rate of reaction.
For example, substances burn more rapidly in pure oxygen
than in air.
• The rate of the reaction is proportional to the product of the
concentration of the reactants. But the rate of the reaction is
equal to the product of the concentration of the reactants
times a proportionality constant (k).
• The constant K takes into account the effect of the nature of
the reactants, temperature, pressure, and catalyst on the
reaction.

26. Chemical Equilibrium

27. Chemical Equilibrium
• Chemical Equilibrium: Exists when two opposing
reactions are occurring simultaneously and at the
same rate.
• Reaction rates and equilibrium are interrelated.
Example: The maintenance of body weight is an
example of a kind of equilibrium.
• When the rate of the forward reaction equals the
rate of the reverse reaction, the system is said to be
in chemical equilibrium.

28. Forward and Reverse Reaction:
N2 + 3H2 → 2NH3
2NH3 → N2 + 3H2
N2 + 3H2 ↔ 2NH3
Reactants Products

29. Chemical Mechanisms

30. Mechanism
• By describing how atoms and molecules interact to
generate products.
• Mechanisms help us to understand how the world
around us functions at a fundamental level.
• A mechanism is a series of elementary steps whose
sum is the overall reaction.
• An elementary step is a reaction that is meant to
represent a single collision or vibration that leads to
a chemical change.

31. Reaction Coordinate Diagrams
• The progress of a reaction on its way from
reactants to products by graphing the energy
of the species versus the reaction coordinate.

32. Reaction Coordinate Diagrams
Figure %: Reaction coordinate diagram for an endothermic reaction

33. Figure %: A reaction coordinate diagram for a
single-step reaction

34. Reaction with three elementary steps
Figure %: Reaction coordinate diagram for a three-step reaction

35. • At higher temperatures the average kinetic energy of the
molecules increases. Therefore, at higher temperatures more
molecules have an energy greater than the activation energy:
Figure %: distributions for T1 greater than T2

36. Transformation Processes
• Dissolution of Species in Water
• Solubility of Non-Aqueous-Phase Liquids
• Dissolution and Precipitation of Solids

37. Table- Water Solubility of Selected Organic Liquids (at T=20oC)
Species Solubility (mg/L)
Benzene 1.780
Benzo(a)pyrene 0.0038
Chloroform 8.200
Dieldrin 0.2
Ethylbenzene 152
Ethylene dibromide 4.300
n-Octane 0.72
Naphthalene 31
Phenol 93.000
2,3,7,8-TCDD 0.0002
Tetrachloroethylene 200
Toluene 535
1,1,1-Trichloroethane 4.400
Trichloroethylene 1.100

38. Table-Solubility Products for Some Ionic Solids
(at T=25oC)
Compound Equilibrium relationship Ksp
Aluminum hydroxide Al(OH)3 ↔ Al3+ + 3OH- 1 X 10-32 M4
Cadmium hydroxide Cd(OH)2 ↔ Cd2+ + 2 OH- 2 X 10-14 M3
Calcium carbonate CaCO3 ↔ Ca2+ + CO32- 5 X 10-9 M2
Calcium fluoride CaF2 ↔ Ca2+ + 2F- 3 X 10-11 M3
Calcium hydroxide Ca(OH)2 ↔ Ca2+ + 2 OH- 8 X 10-6 M1
Calcium phosphate Ca3(PO4)2 ↔ 3Ca2+ + 2 PO43- 1 X 10-27 M5
Calcium sulfate CaSO4 ↔ Ca2+ + SO4 2- 2 X 10-5 M2
Chromium(III) hydroxide Cr(OH)3 ↔ Cr3+ + 3 OH- 6 X 10-31 M4
Iron(II) hydroxide Fe(OH)2 ↔ Fe2+ + 2 OH- 5 X 10-15 M3
Iron(III) hydroxide Fe(OH)3 ↔ Fe3+ + 3 OH- 6 X 10-38 M4
Magnesium carbonate MgCO3 ↔ Mg2+ + CO32- 4 X 10-5 M2
Magnesium hydroxide Mg(OH)2 ↔ Mg2+ + 2 OH- 9 X 10-12 M3
Nickel hydroxide Ni(OH)2 ↔ Ni2+ + 2 OH- 2 X 10-16 M3

39. Dissolution and Precipitation of Solids
• Precipitation reactions occur when different salt
solutions are mixed, which result in the formation of
an insoluble salt, or precipitate.
• When precipitation occurs, the cation of one of the
soluble salts interacts with the anion of the other
soluble salt to form an insoluble salt.
NaCl (aq) + AgNO3 (aq) --> AgCl (s) + NaNO3 (aq)

40. • Most of the precipitation reactions that involve aqueous
salt solutions.
• Salts are compounds which consist of metal cations like
Na+, Ca2+, Cu2+ (or the one nonmetal molecular ion such
as ammonium - NH4+)
ionically bonded to nonmetal anions such as Cl-, (or
molecular anions such as hydroxide - OH-, sulfate - SO42-,
phosphate - PO43-, nitrate - NO3-, and carbonate - CO32-),
dissolved in water.
• Salts can be divided into two types: those soluble in
water, and those insoluble in water.

41. Some Simple Solubility Rules
• Nitrate NO3- salts are soluble
• Salts containing Group 1 metals (Li, Na, K, 1+ charge)
and NH4+ are soluble
• Most Cl-, Br-, and I -salts are soluble,
exceptions of salts that contain Ag+ and, Pb2+

42. • Most sulfate SO42- are soluble
the exceptions of salts containing barium, Ba2+, Pb2+
and Ca2+
• Most hydroxides OH- are just slightly soluble,
the exceptions of the very soluble NaOH and KOH
• Most phosphates PO43- and carbonates - CO32- are
only slightly soluble

44. Kinetic Molecular Theory of
Gases

45. Properties of Gases
1. Gases have low densities
2. Gases are compressible and expandable
3. Gases have no definite volume
4. Gases have no definite shape
5. Gases exert pressure
6. Gases are diffusible
7. The pressure exerted by a confines gas varies directly
with temperature
8. Gases are liquefiable
9. Gases are miscible
10. Gases are stable with respect to volume and pressure

46. Kinetic Molecular Theory of Gases
1. Gases consists tiny particles-molecules, or atoms in the case
the inert gas elements-that are widely separated in space.
2. Molecules of a gas are in constant random motion, traveling
in straight lines at high velocities until they colloid with one
another.
3. Gas molecules behave as if their collisions are perfectly
elastic; the molecules rebound from collision with no
apparent loss of energy.
4. The kinetic energy (energy of motion) of gas molecules
varies directly with temperature expressed in degrees
Kelvin. At any given temperature the molecules have a
certain average kinetic energy.

47. Thank You for Attention