Unit 3

Nuclear Reactions

Types of Nuclear Reactions:

• Fission: The nucleus of an atom breaks apart, releasing a significant amount of energy. Commonly utilized in nuclear reactors to generate power. An example equation is:

  [ \text{Uranium-235} + \text{neutrons} \rightarrow \text{Barium} + \text{Krypton} + 3 \text{neutrons} + \text{energy} ]

• Fusion: Small nuclei combine to form a larger nucleus, releasing energy. This process powers the sun and other stars. Example equation:

  [ \text{2 Hydrogen nuclei} \rightarrow \text{Helium} + \text{energy} ]

• Radioactive Decay: Unstable atoms (radioisotopes) break apart, undergoing various decay processes, typically including:

  • Alpha Decay: Emission of alpha particles, e.g., [ \text{Uranium-238} \rightarrow \text{Thorium-234} + \text{Helium-4} + \text{energy} ]

  • Beta (( \beta )) Decay: Emission of beta particles, e.g., [ \text{Carbon-14} \rightarrow \text{Nitrogen-14} + \text{beta particle} + \text{energy} ]

  • Gamma (( \gamma )) Rays: High energy electromagnetic radiation emitted during decay.

Balancing Nuclear Reactions:

• It is crucial to balance mass and atomic numbers in reactions to ensure conservation of nucleons and charge. Each side of the reaction must have the same number of protons and neutrons.

Calorimetry:

• Importance: Understanding energy transfers in chemical reactions and physical processes such as heat resistance (e.g., Nomex fiber in protective suits).

• Calorimeter: A device used to measure the heat involved in chemical or physical processes. It helps determine the enthalpy changes of reactions.

Calculations:

• Specific Heat Capacity (C): The amount of energy required to raise 1 g of a substance by 1 °C, a specific constant for each substance. Example: water = 4.185 J/(g°C).

• Energy Transfer Formula: [ E = m \cdot C \cdot \Delta T ] where ( \Delta T ) is the change in temperature.

• Enthalpy Definition: The total thermal energy contained in a substance.

• Calculating Enthalpy Changes: [ \Delta H = H_{products} - H_{reactants} ]

• Example Question: Calculate the enthalpy change when 1 mole of water is vaporized at 100°C. Given the standard enthalpy of vaporization is 40.79 kJ/mol, the calculation would be simple since it is provided directly.

Exothermic and Endothermic Reactions:

• Exothermic Reactions: Release energy (e.g., combustion of hydrocarbons).

• Endothermic Reactions: Absorb energy (e.g., photosynthesis).

• Molar Enthalpy Change: [ \Delta H_{reaction} = \frac{\Delta H}{n} ] (for specific amounts).

Types of Enthalpy Changes:

• Include combustion, vaporization, and formation reactions. Each has its specific energy change associated with the transition.

Additional Enthalpy Calculations:

• Neutralization Example: [ \Delta H = \Delta H_{reaction} + \text{other steps} ]

• Combustion Example: Calculate total enthalpy for combustion reactions, e.g., burning methane.

Reaction Energies and Bond Energy:

• Bond Energies: The amount of energy required to break or form a bond between atoms.

• Potential Energy Diagrams: Illustrate energy changes throughout a reaction, highlighting the activation energy and heat changes necessary.

• Reaction Calculations: Average bond energies can be used to estimate enthalpy changes in reactions quantitatively.

Reaction Rates:

• Chemical Kinetics: Study of the rates of chemical reactions, focusing on how concentration of reactants/products changes over time.

• Rate Measurement Methods: Include mass loss, volume of gas produced, electrical conductivity change, or color change for visual indications.

Factors Affecting Reaction Rates:

• Nature of Reactants: Reactants with fast reaction properties (e.g., metals with oxygen).

• Concentration: Higher concentrations generally lead to increased molecular collisions and thus faster reactions.

• Surface Area: More exposed surface area increases the frequency of collisions between reactants.

• Temperature: Increases in temperature raise molecular kinetic energy and collision frequency, leading to higher reaction rates.

• Catalysts: Substances that speed up reactions without being consumed in the process (e.g., enzymes in biological reactions).

Collision Theory:

• Basis: Reactions occur when molecules effectively collide. Effectiveness relies on both the frequency and efficiency of collisions.

• Activation Energy (Eₐ): The minimum energy required for a reaction to occur, which can be lowered by the presence of a catalyst.

Rate Law Equation:

• The equation used to relate the concentration of reactants to the overall rate of the reaction: [ ext{rate} = k[A]^x[B]^y ] where ( k ) is the rate constant, ( [A] ) and ( [B] ) are the concentrations of the reactants, and ( x ) and ( y ) are the reaction orders.

Reaction Mechanisms:

• Describes the step-by-step sequence of elementary steps that constitute a reaction. The slowest step in this sequence is known as the rate-determining step and controls the overall reaction rate, corresponding to the stoichiometry of the rate law equation.