CHEMICAL REACTION AND EQUATION

Chemical Reactions

• Definition:

A process where substances (reactants) are transformed into new substances (products) with different properties.

• Examples:

Rusting of iron, digestion of food, fermentation.

Characteristics of Chemical Reactions

• Change in state
• Change in color
• Evolution of gas
• Change in temperature
• Formation of a precipitate

Chemical Equations

• Purpose: Simplifies the representation of chemical reactions.

• Word Equation

Example:

Methane + Oxygen → Carbon dioxide + Water

Writing a Chemical Equation

• Reactants: Written on the left with a plus sign (+).
• Products: Written on the right with a plus sign (+).
• Arrow (→): Separates reactants from products; indicates the direction of the reaction.

Example

Chemical Equation:

The image shows the chemical equation for the combustion of methane. Here’s the equation:

Explanation

• Reactants:
  • CH44​: Methane.
  • O22​: Oxygen.
• Products:
  • CO22​: Carbon dioxide.
  • H22​O: Water.
• ΔΔ: Indicates the application of heat.

This reaction depicts methane burning in the presence of oxygen, producing carbon dioxide, water, and releasing heat.

Balanced Chemical Equations

• Definition: A balanced equation has equal numbers of each type of atom on both sides, following the law of conservation of mass.

• Law of Conservation of Mass: Mass cannot be created or destroyed in a chemical reaction.

Balancing a Chemical Equation

• Method: Hit and trial method is used to balance equations by adjusting coefficients.

• Steps:
  • Write the unbalanced equation.
  • List the number of atoms of each element on both sides.
  • Balance one element at a time using coefficients.
  • Recheck atom counts to ensure balance.

Example Process

• Equation: Na + H₂O → NaOH + H₂

• Balancing Steps:

• Balance hydrogen by adjusting coefficients.
• Balance sodium next.
• Verify oxygen is balanced.

Making Chemical Equations Informative

1. Physical States:
   1. Use (s) for solids, (l) for liquids, (g) for gases, (aq) for aqueous solutions.
2. Precipitate Formation:
   1. Indicated by a downward arrow (↓).
3. Gas Evolution:
   1. Indicated by an upward arrow (↑).

Additional Information in Equations

• Physical States:

Indicated by (s), (l), (g), (aq).

Explanation

• Reactants:
  • Mg(s): Solid magnesium.
  • H22​SO44​(aq): Aqueous sulfuric acid.
• Products:
  • MgSO44​(aq): Aqueous magnesium sulfate.
  • H22​(g): Hydrogen gas, indicated by the upward arrow (↑).

This reaction illustrates the formation of magnesium sulfate and the release of hydrogen gas when magnesium reacts with sulfuric acid.

• Reaction Conditions:

Temperature, pressure, catalysts noted above or below the arrow.

The image shows a chemical equation for the synthesis of methanol. Here’s the equation:

Explanation

• Reactants:

CO(g): Carbon monoxide gas.

2H₂​(g): Hydrogen gas.

Product:

CH₃OH(l): Liquid methanol.

Conditions:

The reaction occurs at 340 atm pressure.

This reaction involves the use of high pressure to synthesize methanol from carbon monoxide and hydrogen.

• Heat Change:

Evolution or absorption shown as “+Heat” or “-Heat”.

Balancing Chemical Equations

The image shows the chemical equation for the combustion of carbon. Here’s the equation:

Explanation

• Reactants:
  • C(s): Solid carbon.
  • O22​(g): Oxygen gas.
• Product:
  • CO22​(g): Carbon dioxide gas.
  •  

• Heat: The reaction releases heat, indicating it’s exothermic.

This reaction represents the burning of carbon in oxygen to produce carbon dioxide and release heat.

Types of Chemical Reactions

• Combination Reaction

Definition: Two or more reactants combine to form one product.

Examples:

• Decomposition Reaction

Definition: A single reactant breaks down into two or more products.

Types:

• Thermal Decomposition: Uses heat.

• Electrolysis: Uses electrical energy.

The image discusses types of reactions: photolysis, exothermic, endothermic, and displacement reactions. Here’s a summary:

Photolysis or Photochemical Decomposition

Definition: Uses light energy to decompose compounds.

Examples:

• Exothermic and Endothermic Reactions

Exothermic Reactions

Definition: Release heat or energy.

Examples:

Endothermic Reactions

Definition: Absorb heat, energy, or light.

Examples:

Displacement Reaction

Definition: A more reactive element displaces a less reactive one from its compound.

Examples:

Double Displacement Reaction

• Definition: Two different ions or groups of atoms in reactant molecules exchange places.

• Precipitation Reaction: Produces an insoluble substance (precipitate).

• Examples:

Oxidation and Reduction Reactions

Oxidation

Definition: Addition of oxygen or removal of hydrogen.

Examples:

Reduction

Definition: Removal of oxygen or addition of hydrogen.

Examples:

Redox Reactions

Definition: Oxidation and reduction occur simultaneously.

Examples:

These reactions demonstrate the exchange of ions and simultaneous oxidation and reduction processes.

Corrosion Definition: The gradual destruction of metals due to reactions with air, water, and chemicals.

Example:

Rusting of Iron: Formation of reddish-brown iron oxide. Other Examples: Black coating on silver, green coating on copper.

Effects:

Wasteful process causing damage to metal structures like car bodies and bridges. Leads to significant financial costs for repairs and replacements.

Prevention:

Techniques include painting, galvanizing, and electroplating to protect metal surfaces. Rancidity

Definition: Slow oxidation of fats and oils in food, leading to off-flavours and odours.

Prevention Methods:

• Store food in air-tight containers. Refrigerate cooked food.
• Use nitrogen packaging for foods like chips.
• Add antioxidants to delay oxidation.

Understanding these processes and prevention methods helps in maintaining the integrity of metals and the quality of food.