Chemical ReactionsBy

MUHAMMAD FAHAD ANSARI12 IEEM 14

Most chemical reactions can be categorized into five general types:

• Synthesis, • Decomposition, • Single-Replacement, • Double-Replacement, and • Oxidation-Reduction.

Types of Chemical Reactions

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

• Decomposition: One substance breaks down to form two new substances:

AB → A + B

For example: 2 H 2 O → 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

• 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+

Reaction Rates

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.

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.

• 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.

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.

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.

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)

Decomposition of N2O5 (in CCl4) at 45 0C [N2O5] = 1.40 Mdecomposes completely (at time t = ∞, [N2O5] = 0 M)

Time, s Total volume O2, cm3

(at STP)0 0

432 1.32753 2.18

1116 2.891582 3.631986 4.102343 4.46

. .

. .

. .∞a 5.93

Decomposition of Ozone, O3

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.

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.

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.

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.

Exothermic ReactionsReactions 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

Endothermic ReactionsReactions 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 + heatEndothermic

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.

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.

Chemical Equilibrium

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.

Forward and Reverse Reaction:

N2 + 3H2 → 2NH3

2NH3 → N2 + 3H2

N2 + 3H2 ↔ 2NH3

Reactants Products

Chemical Mechanisms

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.

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.

Reaction Coordinate Diagrams

Figure %: Reaction coordinate diagram for an endothermic reaction

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

Reaction with three elementary steps

Figure %: Reaction coordinate diagram for a three-step reaction

• 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

Transformation Processes

• Dissolution of Species in Water • Solubility of Non-Aqueous-Phase Liquids• Dissolution and Precipitation of Solids

Table- Water Solubility of Selected Organic Liquids (at T=20oC)

Species Solubility (mg/L)Benzene 1.780Benzo(a)pyrene 0.0038Chloroform 8.200Dieldrin 0.2Ethylbenzene 152Ethylene dibromide 4.300n-Octane 0.72Naphthalene 31Phenol 93.0002,3,7,8-TCDD 0.0002Tetrachloroethylene 200Toluene 5351,1,1-Trichloroethane 4.400Trichloroethylene 1.100

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

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)

• 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 - SO4

2-, phosphate - PO4

3-, nitrate - NO3-, and carbonate - CO3

2-), dissolved in water.

• Salts can be divided into two types: those soluble in water, and those insoluble in water.

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+

• 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 - CO3

2- are only slightly soluble

Kinetic Molecular Theory of Gases

Properties of Gases1. Gases have low densities2. Gases are compressible and expandable3. Gases have no definite volume4. Gases have no definite shape5. Gases exert pressure6. Gases are diffusible7. The pressure exerted by a confines gas varies directly

with temperature8. Gases are liquefiable9. Gases are miscible10. Gases are stable with respect to volume and pressure

Kinetic Molecular Theory of Gases1. 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.

Thank You for Attention