CHEMICAL REACTIONS

 INTRODUCTION

A chemical reaction is a process in which one or more substances (called reactants) are changed into new substances (called products) with different chemical properties and compositions.
During a chemical reaction, bonds in the reactants are broken and new bonds are formed to produce the products.

Example: 2H2+O2→2H2O

In this reaction, hydrogen and oxygen combine to form water.

REACTANTS: Reactants are the starting substances that undergo chemical change during a reaction. They are present before the reaction begins and are used up as the reaction proceeds.  Reactants determine the type of reaction that will occur and the products that will be formed.

Example: In the reaction:

  Zn+2HCl→ZnCl2+H2

Zinc (Zn) and Hydrochloric acid (HCl) are the reactants.

PRODUCTS: Products are the new substances formed as a result of a chemical reaction. They have different physical and chemical properties from the reactants.  Products appear after the reaction and are usually written on the right-hand side of a chemical equation.

Example: In the reaction:

  Zn+2HCl→ZnCl2+H2

Zinc chloride (ZnCl₂) and hydrogen gas (H₂) are the products.

REACTION TIME: Reaction time is the period between the start and completion of a chemical reaction. It shows how long it takes for the reactants to completely change into products.

Example: When magnesium reacts with hydrochloric acid to produce hydrogen gas, the time taken for the magnesium to disappear completely is the reaction time.
Note: A short reaction time means the reaction is fast, while a long reaction time means the reaction is slow.
 Example: Explosion of gunpowder – very fast (short reaction time)
 Rusting of iron – very slow (long reaction time)

REACTION RATE: The rate of a chemical reaction is the speed at which reactants are converted to products. It measures how fast or slow a reaction takes place.

Mathematically,

  Rate of reaction= change in concentration of reactant or product (mol/dm3)change in time (secs) 

The unit of rate of reaction is mol/dm3s-1 or gdm-3s-1
Rate of reaction can also be expressed as: change in number of mole or mass of reactant or producttime taken for change
Then, the unit of rate is mols-1 or gs-1

Example: The rate of gas formation in the reaction between magnesium and hydrochloric acid, or the rate of color change in the reaction between sodium thiosulphate and hydrochloric acid.

Note:
Fast reactions have high reaction rates.
Slow reactions have low reaction rates.
Reaction rate depends on several factors such as temperature, concentration, surface area, and catalysts.

Example: When 0.5g of calcium trioxocarbonate (iv) was added to excess dilute hydrochloric acid, carbon (iv) oxide was evolved. The complete reaction took 5 minutes. What was the rate of reaction?

SOLUTION:
Rate of reaction = change in mass of reactanttime taken for change

  = (0.5 - 0)g5 x 60 = 0.5300 = 1.67 x 10-3gs-1

COLLISION THEORY

The collision theory states that for a chemical reaction to occur the reactant particles must collide and they must collide with a certain minimum amount of energy known as Activation Energy. 
Reacting particles are in continuous motion, thus they possess energy and they also collide with one another. Not all collisions results in a chemical reaction. Collisions, which result in a chemical reaction are called EFFECTIVE COLLISIONS. The minimum amount of energy required by reacting particles for chemical reaction to occur is called ACTIVATION ENERGY. Activation Energy is the Energy Barrier the reactants must overcome for reaction to occur. It is the minimum energy required for bond breaking for a chemical reaction to occur.

A chemical reaction occurs only when the energy of the colliding reactant particles is equal to more than the activation energy. Activation Energy must be equal to energy barriers also for chemical reaction to occur. Every reaction has its own energy of activation. Reactions with low activation energy have a high rate of reaction and occur spontaneously. Reactions with high activation energy have a low rate of reaction and are not spontaneous.

The graph above shows the concept of activation energy for endothermic and exothermic reactions. From the graph, it can be seen that:

• Both exothermic and endothermic reactions generally require an initial input of energy to overcome the activation energy barrier.
• Exothermic reactions once started, continue to proceed without any further external energy supply. E.g. burning of kerosene.
• Endothermic reactions need a continuous supply of external energy to proceed. E.g. cooking of rice.

FACTORS AFFECTING RATE OF CHEMICAL REACTION

The speed of a reaction depends on several factors that affect how frequently and effectively particles collide. These factors include the nature of the reacting substances, concentration (or pressure), temperature, effect of light, surface area of the reactant and the presence of a catalyst.

(i) Nature of the Reactants: If all other factors are kept constant, different substances will have a different rate of reaction with dilute HCl for example. When dilute HCl reacts with zinc, iron and gold under the same condition, hydrogen gas is evolved fast with zinc, slow with iron and no gas evolved with iron. The difference is due to the chemical nature of the elements as they naturally possess different amounts of energy content. Also, different substances react at different speeds depending on their chemical nature and bond type. 

• Ionic compounds react faster because they involve simple ion exchange and do not require bond breaking.
• Covalent compounds react slower because their strong bonds must be broken before new ones form.

Example:
 The reaction between sodium chloride and silver nitrate is fast (ionic), while the reaction between hydrogen and oxygen is slow (covalent).

(ii) Concentration of Reactants (or Pressure for Gases)

Increasing the concentration of reactants increases the number of particles per unit volume. This causes more collisions per second, leading to high effective collisions and so a faster reaction rate.
For gaseous reactions, increasing pressure has a similar effect — the gas molecules are forced closer together, increasing collision frequency. 
An increase or decrease in the concentration of the reactants will result in a corresponding increase or decrease in effective collisions of the reactants and hence the reaction rate.

Example:
 Dilute hydrochloric acid reacts slowly with magnesium, but concentrated acid reacts faster because there are more acid particles available for collision.

(iii) Temperature

Raising the temperature increases the kinetic energy of the reacting particles. This means the particles move faster, collide more frequently, and with greater energy. As a result, more collisions become effective, and the rate of reaction increases.

Example:
 The decomposition of hydrogen peroxide occurs slowly at room temperature but becomes much faster when heated.

(iv) Catalyst: A catalyst is a substance that alters the rate of a chemical reaction without being chemically changed or used up in the process. Catalysts work by providing an alternative pathway with lower activation energy, allowing more particles to collide effectively. A positive catalyst increases the rate of reaction by lowering the activation energy of the reaction whereas the one which increases the activation energy is known as the negative catalyst or inhibitor.

Example:

• Manganese(IV) oxide (MnO₂) speeds up the decomposition of potassium chlorate (KClO₃).
• Iron (Fe) is used as a catalyst in the Haber process for producing ammonia.

Note: Some catalysts are specific — they only work for certain reactions.

Properties of a catalyst:

• It does not start a reaction.
• It remains unchanged in chemical nature and mass at the end of the reaction.
• It is specific in its action.
• It has no effect on the equilibrium of a reversible reaction.

(v) Surface Area of Reactants

When a solid reactant is broken into smaller pieces or powdered, its surface area increases. This allows more particles to be exposed for collision with other reactants, leading to a faster reaction.

Example:
 Powdered zinc reacts faster with acid than a large lump of zinc because it has a greater surface area.

(vi) Presence of Light (in some reactions)

Some reactions are influenced by light, known as photochemical reactions. Light provides the energy required for bond breaking and effective collisions.

Example:
 The decomposition of silver chloride and photosynthesis in plants are reactions that occur only in the presence of light.

CHEMICAL EQUILIBRUM

INTRODUCTION:
Equilibrium can be defined as the state where there is no observable or detectable change in the properties of a system with respect to time. Examples include: 
 - A balanced seesaw
 - A saturated solution of NaCl

STATIC AND DYNAMIC EQUILIBRUM

Static Equilibrium is attained when a system in  equilibrium is in a state of rest or stationary state. E.g. a balanced see saw.

Dynamic Equilibrium occurs during a physical or chemical change that is reversible. Dynamic equilibrium involving a physical change is referred to as physical equilibrium while that involving a chemical change is a chemical equilibrium. A system is said to be in a state of dynamic equilibrium when both the forward and backward reactions take place at the same rate. Example:

  N2O4(g) ↔ 2NO2

Chemical Equilibrium is a state in a reversible reaction where the rate of the forward reaction equals the rate of the backward reaction, and the concentrations of the reactants and products remain constant over time. It does not mean the reactions have stopped, but that they occur at the same rate in both directions.

For example:
N2(g)  +  3H2(g)    ⇌  2NH3(g)
 In this reaction, nitrogen and hydrogen react to form ammonia. At equilibrium, the rate at which ammonia is formed equals the rate at which it decomposes back into nitrogen and hydrogen.
        NH4Cl(s)     ⇌     NH3(g)  +  HCl(g)

Characteristics of Chemical Equilibrium

• It occurs in a closed system.
• The rate of the forward and backward reactions are equal.
• The concentrations of reactants and products remain constant.
• It is dynamic, meaning the reactions continue to occur, but with no net change.

Le Chatelier’s Principle
 Le Chatelier’s Principle states that if a system at equilibrium and one of the factors involved in equilibrium (such as temperature, pressure, or concentration) is altered, the equilibrium position will shift to neutralize the effect of the change and restore balance.

Applications of Le Chatelier’s Principle:

• Change in Concentration:
   If the concentration of reactants is increased, the equilibrium shifts to the right (favoring the forward reaction) to produce more products while increase in concentration of products shifts the equilibrium toward the reactants (backward direction).
   Example: In the reaction N2+3H2⇌2NH3​, increasing N2 or H2​ shifts the equilibrium to the right.
  
   
• Change in Temperature:
   If the reaction is exothermic, increasing the temperature shifts the equilibrium to the left (favoring the reverse reaction), while decreasing temperature shifts it to the right.
   Example: In the Haber process (exothermic), higher temperature reduces ammonia yield.
  
   
• Change in Pressure:
   Increasing pressure favors the side with fewer gas molecules, while decreasing pressure favors the side with more gas molecules.
   Example: In N2+3H2⇌2NH3​, increasing pressure favors ammonia formation (fewer molecules on the product side).

HABER PROCESS

Haber Process is an industrial process for the manufacturer of ammonia. Ammonia, NH3, is an important chemical for the manufacture of fertilizers. During the process, the following equilibrium is established.

  N2(g)   +       3H2        ⇌        2NH3(g) H = -96KJmol-1

          1mole       3moles           2moles 

A Flow Scheme of Haber Process

Choice of Temperature: 
Since the forward reaction is exothermic, Le Chatelier’s principle shows that the greatest yield of ammonia, NH3, would be obtained at low temperature. However, if the temperature is too low, the time taken to reach equilibrium would be much. A temperature of 5000C gives a reasonable yield of ammonia.

Choice of Pressure: 
Since forward reaction leads to a decrease in the number of moles (pressure), Le Chatelier’s principle shows that high yield of ammonia would be obtained at high pressure. The pressure of 25atm gives a satisfactory yield.

Choice of Catalyst:
Iron is used as a catalyst in the Haber process. Iron is found to be effective in speeding up the rate at which equilibrium is reached. It is used in the form of small pellets which increased the surface area available for contact.

EQUILIBRIUM CONSTANT (K) AND LAW OF MASS OF ACTION
The law of mass action states that at a constant temperature, the rate of reaction is proportional to the active masses of each of the reactants. The active mass of the reactants is the concentration of the substance raised to the power of coefficient (number of moles) of that substance in the equation of reaction under study.

Note: Active mass = molar concentration x activity coefficient.

In most reactions, the active masses of the reactant may be taken as their molar concentrations. Example:  aA  +     bB  → products

Thus,   r   ∝    [A]a      and   r   ∝    [B]b , where [ ] represents concentration in mol/dm3.

r   ∝    [A]a  x   [B]b

r    =   k[A]a[B]b    , where k is called equilibrium constants.

Applying the law to the reversible reaction represented by the equation below:

aA(aq)    +      bB(aq) ⇌ cC(aq)     +   dD(aq)
Kc   = [C]c[D]d
          [A]a[B]b where Kc is called equilibrium constant. 

Therefore, for the general gaseous reaction: 

aA(g)    +      bB(g) ⇌ cC(g)     +   dD(g)

Kp   = PC    x     PD
            PA       x   PB

Where Kp is the equilibrium constant for the gaseous reaction, PA, PB, PC and PD are partial pressures of the gases A, B, C and D respectively.

Note: For a given reaction, concentration of solids are ignored. A Kc or Kp value greater than one tells you that product formation is favoured at equilibrium while a Kc or Kp value less than one tells you that reactant formation is favoured at equilibrium. A system in equilibrium has Kc or Kp value equal to one.