Chemistry of Chemical Reactions and Enthalpy Changes

Reactivity in Chemical Reactions

Measuring Enthalpy Changes

  • Energy Transfer in Chemical Reactions

  • Total energy is conserved in chemical reactions.

  • Chemical potential energy is stored in bonds of reactants/products.

  • Reaction temperature relates to kinetic energy of atoms, ions, and molecules.

Reaction Systems

  • Energy Changes

  • All chemical reactions involve energy changes.

  • Energy may be released or absorbed during reactions.

  • Types of Systems:

  • Open System: Transfer of matter and energy possible across the boundary (e.g., beaker).

  • Closed System: No transfer of matter; energy can be transferred.

  • Isolated System: No matter or energy can enter or exit the system.

Heat vs. Temperature

  • Temperature (T):

  • A state function; changes are independent of the pathway.

  • Example: ΔT = Tfinal − Tinitial (difference in temperature measurement).

  • Heat (q):

  • Energy transferred from warmer to cooler bodies.

  • Increases average kinetic energy of particles when transferred.

  • At absolute zero (0 K), all particle motion theoretically stops; entropy is at a minimum.

Thermochemistry

  • Definition:

  • Study of heat changes during chemical reactions, often described in terms of enthalpy (ΔH).

Exothermic and Endothermic Reactions

  • Exothermic Reactions:

  • Energy is released when bonds are formed; bond making is exothermic.

  • Example: Zn(s) + CuSO4(aq) → Cu(s) + ZnSO4(aq)

    • Enthalpy change = -217 kJ mol-1, indicating a decrease in energy for products.

  • Endothermic Reactions:

  • Energy is absorbed to break bonds; bond breaking is endothermic.

  • Example: NH4NO3(s) → NH4+(aq) + NO3-(aq)

    • Enthalpy change = +25.7 kJ mol-1, indicating an increase in energy for products.

Energy Profiles

  • Definition:

  • Visual representation of enthalpy changes during a reaction.

  • Provides information on enthalpy of reactants/products and activation energy (Ea).

Standard Enthalpy Change (ΔH)

  • Defined as heat transferred at constant pressure under standard conditions, units: kJ/mol.

  • Calculation involves heat lost or gained by a substance, using specific heat capacity (c).

Specific Heat Capacity

  • Defined as heat needed to raise the temperature of 1 kg of a substance by 1 °C or 1 K.

  • Example: Specific heat capacity of ethanol = 2.44 kJ/kg·K.

  • Lower specific heat leads to higher temperature rise for a given heat input.

  • Specific heat capacity does not vary with the size of the system.

Relevant Formulas for Specific Heat Capacity

  • Amount of heat (Q) absorbed by substance: Q = mcΔT

  • Where m = mass (kg), ΔT = change in temperature (K).

  • Heat (Q) relates to enthalpy change (ΔH) depending on the number of moles of the limiting reactant.

Bomb Calorimeter

  • Used in calorimetry experiments to measure heat transfer during chemical reactions.

  • Sample burned in a bomb chamber; resulting temperature change of surrounding water measured.