CHEM2: Thermodynamics

• Thermodynamics - Study of energy and its transformations.
• Energy - The capacity to do work or transfer heat. Units are Joule and calorie.
• 1 cal = 4.184 J
• 2 types of energy:

  1. Kinetic Energy
  2. Potential Energy
• Kinetic Energy - Energy of motion.
• Potential Energy - Energy an object possesses by virtue of its position.
• Thermochemistry - Study of heat change in chemical reactions.
• Heat - Transfer of thermal energy between two bodies that are at different temperatures.
• Temperature - A measure of the thermal energy.
• Temperature is not the same as thermal energy.
• System - Part of the universe we are interested in (this is where the chemical reaction takes place). Ex: water molecules - water, ice
• Surroundings - The rest of the universe. Ex: everything else - cup, air, etc.
• 3 types of system:

  1. Open System - Ex: mass and energy
  2. Closed System - Ex: energy only
  3. Isolated System - Ex: nothing
• Potential energy is a state function, while work is a path function.
• State Functions - Properties that depend only on the initial and final states of the system and not on how the change is accomplished to reach the state. Path independent. Ex: pressure, volume, temperature
• Path Functions - Properties or quantities whose values depend on the transition of a system from the initial state to the final state. Ex: work and heat
• Internal Energy - Symbol is U. Total energy of a system (kinetic + potential). State function.
• First Law of Thermodynamics - Energy cannot be created nor destroyed. Any energy lost from a system must be transferred to the surroundings (and vice versa). Total change in system's internal energy is the sum of the energy transferred as heat and/or work.
• Formula for total change in system's internal energy: ΔU = q + w
• +q = Gained by the system.
• -q - Lost by the system.
• +w - Work done on the system.
• -w - Work done by the system.
• Endothermic - Absorb energy. Breaking bonds. Heat enters in the system. Ex: change of state (solid > liquid > gas), melting, evaporation, sublimation
• Exothermic - Release energy. Forming bonds. Heat exits from the system. Ex: change of state (gas > liquid > solid), condensation, freezing, deposition
• Enthalpy - Symbol is H. Heat transferred between the system and surroundings carried out under constant pressure. State function.
• ΔH - Change in enthalpy. "Enthalpy of reaction." H (products) - H (reactants)
• Enthalpy is an extensive property.
• When we reverse a reaction, we change the sign of ΔH.
• Change in enthalpy depends on state.
• Hess's Law - If a reaction is carried out in a number of steps, ΔH for the overall reaction is the sum of ΔH for each individual step.
• Enthalpy of formation - Enthalpy change for the formation of 1 mol of compound from its constituent elements.
• Standard Enthalpy - The enthalpy measured when everything is in its standard state.
• Standard enthalpy of formation - Change in enthalpy if 1 mol of compound is formed from its element with all substances in their standard states.
• Standard enthalpy of formation of the most stable form of an element is zero.