Module 8: Thermochemistry

• Thermochemistry (or chemical thermodynamics), is the study of the heat energy and work associated with chemical reactions and/or physical transformations.
• System - part of the universe we want to study (what are you observing?)

  Surroundings - the rest of the universe (normally the immediate surroundings)

  Process - a physical occurrence (usually involving energy flow)
• Open System - matter and energy can be exchanged with the surroundings

  Closed System - only energy and not matter can be exchanged with the surroundings

  Isolated System - neither energy nor matter can be exchanged with the surroundings
• Internal Energy - the combined kinetic and potential energies of atoms and molecules that make up an object or system

  Chemical Energy - energy stored in bonds and intermolecular attractions; is the energy released or absorbed during a chemical reaction
• In an endothermic change:

  • energy and heat is absorbed by the system
  • the surrounding does work on the system
  • the energy is on the reactants side (for the process to proceed)
• In an exothermic change:

  • energy and heat is released by the system
  • the system does work on the surroundings
  • the energy is on the products side
• Thermal Equilibrium - heat flows from warmer object to cooler object until the temperatures become equal
• Isothermal Process - change in phase happens in one temperature
• Heat capacity, C of a system is the quantity of heat required to change the temperature of a system by one degree. The higher the heat capacity, the more heat/energy needed to raise the temperature.
• The Law of Conservation of energy states that in interactions between a system and its surroundings, the total energy remains constant. Meaning, the heat gained by the system is lost by the surroundings, and vice versa.
• In solving problems involving thermochemistry, it is important to:

  1. Organize your given values
  2. Know the parameters or conditions
  3. Set-up the equations to be used
• Heat of reaction (qrxn) is the quantity of heat exchanged between a system and its surroundings when a chemical reaction occurs within the system at constant pressure.
• In an isolated system, qrxn leads to a temperature change within the system. In a non-isolated system, the temperature remains constant, and the heat is transferred to the surroundings.
• Calorimetry is the laboratory method for determining heat changes and relating it to energy changes. A calorimeter is used to measure quantities of heat.
• Bomb calorimeter - constant volume, measure internal energy

  Styrofoam / Coffee-cup calorimeter - constant pressure, measure enthalpy
• External pressure - system does work on the surroundings, so the sign is negative
• Expansion - positive change in volume

  Compression - negative change in volume
• A system contains only internal energy (U or E). It does not contain heat or work. These only occur when there is a change in the system.
• State functions are properties that are determined by the state (initial and final) of the system, regardless of how that condition was achieved.
• Path functions are properties or quantities whose values depend on the transition of a system from the initial to the final state (only appear when there are changes in a system).
• Enthalpy is the heat flow under conditions of constant pressure. It is an extensive property.
• During a phase change, temperature does not change with heat flow due to formation or breaking of intermolecular attractive forces.
• A formation reaction is the chemical reaction by which one mole of a compound is formed from its elements in their standard states (or the most stable form of an element at room temp, 25 degrees celcius, and pressure, 1 bar).
• Enthalpy change is directly proportional to the amount of substance in a system.
• Change in enthalpy changes sign when a process is reversed.
• Hess's Law - If a process occurs in stages or steps, the enthalpy change for the overall process is the sum of the enthalpy changes for the individual steps.
• The standard enthalpy of formation of a pure element in its reference state is zero.
• Absolute enthalpy cannot be defined, only the change in enthalpy.
• Entropy, S is a measure of the randomness or disorder of a system.
• For any substance, the solid state is more ordered than the liquid state and the liquid state is more ordered than the gas state.
• The entropy of the universe increases in a spontaneous process and remains unchanged in an equilibrium process.
• If a reaction produces more gas molecules than it consumes, ΔS° > 0.
• If the total number of gas molecules decreases, ΔS° < 0.
• If there is no net change in the total number of gas molecules, then ΔS° may be positive or negative but ΔS° will be a small number.
• The entropy of a perfect crystalline substance is zero at the absolute zero of temperature.