Chemical Reactions and Equations

Chemical Reaction and Equation

• A chemical reaction involves two or more reactants combining to form a product or new substance.

  • Example:
    $A + B <br>ightarrow AB$

Physical vs Chemical Changes

• Physical Change:

  • Usually reversible.

  • Interchanging possible; new substance is not formed.

  • Example: Melting of ice.

• Chemical Change:

  • Irreversible; interchanging is not possible.

  • New substance is formed through a chemical reaction.

  • Example: Rusting.

Indicators of a Chemical Reaction

• Change in state

• Change in color

• Change in temperature

• Evolution of gas

• Formation of precipitate

  • A precipitate is an insoluble substance formed in a chemical reaction.

Definitions

• Reactant: A substance that undergoes a reaction to form products.

• Product: A substance that is formed after a reaction.

Skeletal Equations

• A skeletal equation shows the reactants and products but is not balanced.

  • Example:
    $Mg + O_2 <br>ightarrow MgO$

Importance of Balancing Equations

• Balancing is necessary to satisfy the Law of Conservation of Mass.

  • Mass cannot be created or destroyed in a chemical reaction.

  • Example of a balanced equation:
    $2H<em>2 + O</em>2 <br>ightarrow 2H_2O$

Symbols for States of Matter

• (g) = gaseous

• (aq) = aqueous

• (s) = solid

• (l) = liquid

Types of Chemical Reactions

1. Combination Reaction

   • Two or more reactants form a single product.

   • Example:
     $2Mg + O_2 <br>ightarrow 2MgO$

2. Decomposition Reaction

   • A single reactant breaks down into two or more products.

   • Example:
     $AB <br>ightarrow A + B$

   • Thermal Decomposition: Heat is required to break down the substance.

     • Example:
       $Zn + H<em>2SO</em>4 <br>ightarrow ZnSO<em>4 + H</em>2 + Heat$

3. Displacement Reaction

   • More reactive element displaces a less reactive element from its compound.

   • Single Displacement and Double Displacement reactions fall under this category.

4. Oxidation and Reduction

   • Oxidation: Loss of electrons or increase in oxidation state.

   • Reduction: Gain of electrons or decrease in oxidation state.

Examples of Oxidation and Reduction

• Respiration:

  • Example:
    $C<em>6H</em>{12}O<em>6 + 6O</em>2 <br>ightarrow 6CO<em>2 + 6H</em>2O + Energy$

• Combustion of Natural Gas:

  • Example:
    $CH<em>4 + 2O</em>2 <br>ightarrow CO<em>2 + 2H</em>2O + Heat$

Endothermic and Exothermic Reactions

• Endothermic Reaction: Absorbs heat.

• Exothermic Reaction: Releases heat.

Special Reaction Types

• Electrolysis: Uses electricity to decompose substances.

• Photolysis: Uses light energy to decompose substances.

Conclusion

• Understanding chemical reactions and their equations is crucial in chemistry for predicting the behavior of substances during a reaction, their classifications, and energy changes involved.