Module 7 Lecture Notes

Types of Chemical Reactions

1. Decomposition Reactions:

   • Example: CuCO3(s) \rightarrow CuO(s) + CO2(g)

   • (green) decomposes into (black) solid and carbon dioxide gas.

2. Precipitation Reactions:

   • Example: Pb(NO3)2(aq) + 2KI(aq) \rightarrow PbI2(s) + 2KNO3(aq)

   • (colorless) + (colorless) yields (yellow) precipitate.

3. Acid-Base Reactions:

   • Example: NaOH(aq) + HCl(aq) \rightarrow NaCl(aq) + H_2O(l)

4. Complex Ion Formation:

   • Example: Cu^{2+}(aq) + 4NH3(aq) \rightarrow Cu(NH3)_4^{2+}(aq)

5. Redox (Oxidation-Reduction) Reactions:

   • Example: 2Mg(s) + O_2(g) \rightarrow 2MgO(s)

   • (white)

Chemical Reactions and Energy Changes

• Chemical reactions are accompanied by energy changes.

• Energy can be absorbed or released as:

  • Heat (e.g., acid + base)

  • Light (chemiluminescence reaction)

  • Electricity (redox reactions)

Examples of Energy Changes

1. Heat Energy Released:

   • Burning methane: CH4 + 2O2 \rightarrow CO2 + 2H2O

   • Respiration of glucose: C6H{12}O6 + 6O2 \rightarrow 6CO2 + 6H2O

2. Light Energy Absorbed:

   • Photosynthesis: 6CO2 + 6H2O \rightarrow C6H{12}O6 + 6O2

3. Electrical Energy Absorbed:

   • Electrolysis of water: 2H2O \rightarrow 2H2 + O_2

Heat of Reaction and Conservation of Energy

• Energy is not created or destroyed; it can be converted into different forms.

• Potential energy converts to kinetic energy.

Enthalpy (H)

• Heat of reaction: Energy released (e.g., burning of methane).

• Internally stored chemical energy is called Heat Content or Enthalpy (H).

• Exothermic: Heat released (surroundings get hot).

• Endothermic: Heat absorbed (surroundings get cold).

• \Delta H (change in enthalpy) is negative for exothermic reactions.

• \Delta H (change in enthalpy) is positive for endothermic reactions.

Exothermic Reaction

• Reactant -> Product + Heat

• ΔH is -ve

Endothermic Reaction

• Reactant + Heat -> Product

• ΔH is +ve

Endothermic Reaction - Example

Ba(OH)2.8H2O(s) + 2NH4NO3(s) \rightarrow Ba(NO3)2(aq) + 2NH3(g) + 10H2O(l)
Ba(OH)2.8H2O(s) + 2NH4SCN(s) \rightarrow Ba(SCN)2(aq) + 2NH3(g) + 10H2O(l)

Representing Heats of Reaction

1. As \Delta H (as a separate term).

2. As a ‘reactant’ or ‘product’ (within the reaction).

Examples

• CH4 + 2O2 \rightarrow CO2 + 2H2O \Delta H = -889kJ

• CH4 + 2O2 \rightarrow CO2 + 2H2O + 889 kJ

• 6CO2 + 6H2O \rightarrow C6H{12}O6 + 6O2 \Delta H = +2816 kJ

• 6CO2 + 6H2O + 2816 kJ \rightarrow C6H{12}O6 + O2

Rates of Reaction

• Speed of a reaction varies considerably.

  • Fast: Burning H_2, Precipitation.

  • Slow: Rusting, Decomposition of H2O2, oxidation of vitamin C.

• Important consideration: the nature of the reaction (how many and what type of bonds need to be broken).

• This determines the Activation Energy of the reaction.

Factors Affecting Rates of Chemical Reactions

1. Nature of reaction.

2. Concentration

   • Gas pressure

   • Solution concentration

   • State of sub-division of reactants

3. Temperature.

4. Catalyst.

Collision Theory

• For a reaction to occur, molecules must have sufficient energy and be in the correct orientation.

• Sufficient energy + correct orientation.

Reaction Profile Diagram

• Shows the energy changes during a chemical reaction.

• Activation Energy (E_a): Minimum energy required for reactants to react to form products.

• Transition State: Highest energy state in a chemical reaction.

Explanation of Factors Affecting Reaction Rates

• Nature of reaction: Depends on activation energy (how much bond breaking is involved).

• Concentration/gas pressure/state of subdivision: Increasing frequency of collisions.

• Temperature: Increased frequency of collisions, more high-energy collisions (more successful collisions).

• Catalyst: Provides alternative reaction pathway with lower activation energy, e.g., H2O2 decomposition is faster with MnO_2.

Effect of a Catalyst

• KMnO4 • H2O2(l) \rightarrow H2O(l) + O_2(g) + energy

Measuring Reaction Rate

• The rate at which reactants are lost.

• The rate at which products are formed.

Factors Controlling Reaction Rate

1. sufficient energy to break the bonds in the reactant molecules,

2. the correct orientation that allows bonds in the reactants to be broken and
   bonds in the products to be formed. |

Additional Explanation:

• Nature of the Reactants: If strong bonds need to be broken, the reaction is slow.

• Surface Area: Increased surface area results in more frequent collisions between reactant particles.

• Concentration: Higher concentration results in more frequent collisions between reactant particles.

• Temperature: Increased temperature results in more frequent collisions at higher energies.

• Catalysts: A catalyst lowers the activation energy required, resulting in more successful collisions.

Energy Changes During a Chemical Reaction

• Activated Complex: The molecule or ion that forms when the reactants collide.

• Energy of Activation, E_a: The minimum energy for the collision between the reactants to be successful.

• Heat of Reaction: The difference in enthalpy between the reactants and products (an exothermic reaction is shown).

• Catalysts: The activation energy with the presence of a catalyst is lower than the E_a.

Energy Profile Diagram

• E_a: Energy of Activation without a catalyst.

• E_e: Energy of Activation with a catalyst.

• \Delta H is the Heat of Reaction.