Calorimetry is all about measuring heat transfer during chemical reactions or physical processes. It's based on the idea that energy is conserved, so the heat released or absorbed by a system equals the heat absorbed or released by its surroundings.

To measure heat transfer, we use calorimeters. These devices monitor temperature changes in a known amount of substance, usually water. By calculating the heat change, we can determine the energy involved in reactions, which is crucial for understanding thermochemistry.

Calorimetry

Heat transfer measurement in calorimetry

• Calorimetry measures heat transfer during chemical reactions or physical processes based on conservation of energy principle
• Heat released or absorbed by system equals heat absorbed or released by surroundings
• Monitors temperature change in known amount of substance (water) in calorimeter
  • Substance undergoing temperature change called calorimeter's contents
• Temperature change of calorimeter's contents proportional to heat transferred during reaction or process
  • Proportionality constant is heat capacity of calorimeter's contents
• Must consider heat capacity of calorimeter itself in calculations
  • Known as calorimeter constant or heat capacity of calorimeter

Calculations with calorimetry data

• Calculate heat change using equation: $q = mc\Delta T$
  • $q$: heat energy transferred (J)
  • $m$: mass of substance (g)
  • $c$: specific heat capacity of substance (J/g°C)
  • $\Delta T$: change in temperature (°C)
• Specific heat capacity: amount of heat required to raise temperature of 1 g of substance by 1°C
  • Unique to each substance, found in reference tables
• When using calorimeter, heat change equation becomes: $q = (mc\Delta T){\text{contents}} + (C\Delta T){\text{calorimeter}}$
  • $(mc\Delta T)_{\text{contents}}$: heat change of calorimeter's contents
  • $(C\Delta T)_{\text{calorimeter}}$: heat change of calorimeter itself
  • $C$: calorimeter constant (J/°C)
• In calorimetry experiment, heat lost by system equals heat gained by surroundings (calorimeter and contents), assuming no heat lost to environment
  • Represented by equation: $q_{\text{system}} + q_{\text{surroundings}} = 0$
  • This principle is based on the law of energy conservation in thermochemistry

Types of calorimeters and applications

• Two main types: constant-pressure calorimeters and bomb calorimeters
• Constant-pressure calorimeters (coffee-cup calorimeters) operate at constant pressure (usually atmospheric)
  • Measure heat changes in reactions not involving gases or producing gases at low pressures
  • Examples: measuring heat of solution, neutralization, combustion of food or fuel samples
  • Simple and inexpensive but less precise than bomb calorimeters
• Bomb calorimeters operate at constant volume, measure heat changes in combustion reactions
  • Reaction takes place inside sealed, pressurized "bomb" filled with oxygen
  • Bomb submerged in known amount of water inside insulated container
  • More precise than constant-pressure calorimeters, can measure heat of combustion of solids, liquids, and gases
  • More expensive and complex to use than constant-pressure calorimeters
  • Designed to maintain an adiabatic process, minimizing heat exchange with the surroundings

Thermal Equilibrium and Energy Transfer

• Calorimetry relies on the principle of thermal equilibrium between the system and surroundings
• Heat flows from higher temperature regions to lower temperature regions until thermal equilibrium is reached
• The total energy of the system and surroundings remains constant, demonstrating energy conservation
• Accurate measurements in calorimetry depend on achieving and maintaining thermal equilibrium