7.1 Heat and energy changes

Heat and energy changes are the heart of thermochemistry. They explain how energy moves between systems and surroundings. Understanding these concepts helps us predict and control chemical reactions, from cooking to industrial processes.

Temperature measures a system's average kinetic energy, while heat is energy transfer due to temperature differences. Endothermic processes absorb heat, exothermic processes release it. These ideas form the foundation for understanding energy flow in chemical reactions.

Influence of Heat and Temperature on Energy Changes

Defining Heat and Temperature

• Heat is the transfer of thermal energy between two systems due to a temperature difference
  • Measured in units of joules (J) or calories (cal)
• Temperature is a measure of the average kinetic energy of the particles in a system
  • Measured in units of degrees Celsius (°C), Fahrenheit (°F), or Kelvin (K)
• Heat is a process, while temperature is a state variable that describes the thermal energy content of a system
• The direction of heat transfer is always from a higher temperature system to a lower temperature system (second law of thermodynamics)

Endothermic and Exothermic Processes

• Endothermic processes absorb heat from the surroundings
  • Results in an increase in the system's internal energy and a decrease in the surroundings' energy
  • Sign convention is positive (heat is absorbed, q > 0)
  • Examples include melting, vaporization, and sublimation
• Exothermic processes release heat to the surroundings
  • Results in a decrease in the system's internal energy and an increase in the surroundings' energy
  • Sign convention is negative (heat is released, q < 0)
  • Examples include freezing, condensation, and deposition

Applying the First Law of Thermodynamics

Calculating Energy Changes in a System

• The first law of thermodynamics states that the change in internal energy (ΔU) of a system is equal to the heat (q) added to the system plus the work (w) done on the system: $ΔU = q + w$
• For processes occurring at constant volume (w = 0), the change in internal energy is equal to the heat added to or removed from the system: $ΔU = q_v$
• For processes occurring at constant pressure, the change in enthalpy (ΔH) is equal to the heat added to or removed from the system: $ΔH = q_p$
• The work done by a system can be calculated using the equation $w = -PΔV$, where P is the external pressure and ΔV is the change in volume of the system

Heat Capacity and Thermal Energy Transfer

• Heat capacity is the amount of heat required to raise the temperature of a substance by one degree Celsius or Kelvin
• Specific heat capacity (c) is the amount of heat required to raise the temperature of one gram of a substance by one degree Celsius or Kelvin: $c = q / (mΔT)$, where q is the heat added, m is the mass of the substance, and ΔT is the change in temperature
  • Substances with higher heat capacities require more energy to change their temperature (water)
  • Substances with lower heat capacities require less energy to change their temperature (metals)
• The molar heat capacity (C) is the amount of heat required to raise the temperature of one mole of a substance by one degree Celsius or Kelvin: $C = q / (nΔT)$, where n is the number of moles of the substance
• The transfer of thermal energy between two systems depends on their heat capacities and the temperature difference between them, as described by the equations $q = mcΔT$ or $q = nCΔT$