● Calorimetry: Measurement of heat that’s exchanged or transferred in a chemical system
● Critical Temperature: Highest Temperature at which a liquid phase can form
● Dynamic Equilibrium: when evaporation and condensation occur at the same rate
● Endothermic Reaction: Energy of the system increases so the energy must be taken up by absorption from the surroundings
● Energy: The ability to do work (Metric Unit: Joule or J). There are many forms of energy, and it can be transferred from one form to another
● Exothermic Reaction: The energy of the system decreases, and energy is released into the surroundings
● Heat: Transfer of energy between two objects with different temperatures (in Joules)
● State Function: A property that describes an object’s current condition (doesn’t depend on how it got there). Total energy is a state function but work is not a state function because it depends on the pathway, as well as heat.
● Heating Curve: Temperature of system vs. amount of heat added. Temperature stops at the points where substance changes state
● Spontaneous Reactions: Reaction will occur within a certain set of conditions. If a reaction can increase in entropy, it will likely be spontaneous
● System: The change being studied
● Work: Force exerted over a distance (Joules)
● Heat energy change that takes place when reactants go to products
● Changes in Enthalpy occur during chemical reactions
● A change in enthalpy can be from many variables but temperature is one of the most important
● To standardize enthalpies of reaction, data is presented for reactions in which both reactants and products have the standard thermodynamic temperature of 25˚C or 298.15 K.
○ Using a list of standard enthalpies of formation, the enthalpy change of any reaction for which there is data available can be calculated:
ΔHreaction = ΔHf0products - ΔHf0reactants
○ ΔHreaction is in kJ or Joules (moles cancel out)
○ Ex. Calculate the enthalpy change for the following reaction.
■ HCl(g) + NH3(g) → NH4Cl(s)
● ΔHreaction = ΔHf0products - ΔHf0reactants
● ΔHf0product = (1 mol)(-314.4 kJ/mol) = -314.4 kJ
● ΔHf0 reactants = [(1 mol)(-92.3 kJ/mol)+(1 mol)(-45.9 kJ/mol)] = -138.2 kJ
● ΔH_reaction = (-314.4 kJ) – (-138.2 kJ)
● (exothermic reaction), -176.2 kJ
● Calorimetry can be used to determine enthalpies of vaporization, fusion, reaction, and heat capacities
● When something is heated up, it takes different amounts of time to heat it up depending on what it's made up of.
○ Some materials require more energy to raise their temperature
○ Amount of energy needed to raise the temperature of 1 gram of an object by 1 degree celsius
○ A high specific heat means that the object requires a lot of energy to change temperature
○ Water has a high specific heat, metals have low specific heat
○ Used to figure out how much energy is required
○ Q = Energy
○ m = Mass
○ c = Specific Heat Capacity
○ Δt = Change in Temperature
○ Temperature doesn’t change so use, q = ΔHfusion * moles (melting)
■ ΔHfusion is the amount of heat required to melt 1 mol of solid
■ Q = ΔHvaporization * moles (Boiling)
■ ΔHvaporization is the amount of heat required to boil 1 mol of liquid
○ Vaporization is liquid to gas, Condensation is gas to liquid
○ Melting is solid to liquid, Freezing is liquid to solid
○ Deposition is gas to solid, Sublimation is solid to gas
○ If two bodies are in thermal equilibrium with some third body, than they are also in equilibrium
○ If A = B, B = C, then A = C
○ Energy can neither be created nor destroyed. It can only change forms. In any process, the total energy of the universe remains the same.
○ ΔU = q + w
■ ΔU is the total change in internal energy of a system
■ q is the heat exchanged between the system and its surroundings
■ w is the work done by or on the system
○ Work is also equal to negative external pressure on the system, called pressure-volume work
■ w = -p(Δ)V
■ P is the external pressure on the system
■ V is the change in volume
○ Any work or system that goes into or out of a system changes the internal energy
■ Energy is never created nor destroyed, the change in internal energy always is 0
■ If energy was absorbed into a system, than the energy was released into the surroundings
● ΔUsystem = -ΔUsurroundings
● ΔUsystem is the total internal energy
● ΔUsurroundings is the total energy of the surroundings
○ Heat energy can’t be transferred from a body at a lower temperature to a body with higher temperature without the addition of energy
○ ΔSuniverse = ΔSsystem + ΔSsurroundings > 0
○ Gibbs’ Free Energy
○ ΔG = ΔH - T(Δ)S
■ ΔH is the heat change for a reaction
■ ΔG is the measure of change of a system’s free energy in which a reaction takes place at constant temperature and pressure
○ When entropy decreases and enthalpy increases the free energy change, ΔG is positive and not spontaneous. Temperature comes into focus only when both the enthalpy and entropy increase or decrease
○ Reaction is not spontaneous when both enthalpy and entropy are positive with a low temperature but is spontaneous when there is high temperatures
○ Reaction is spontaneous when both enthalpy and entropy are negative at low temperature but not spontaneous when in higher temperatures
○ As temperature reaches absolute zero, the entropy of a system reaches a constant minimum
● Main idea behind the second and third laws of thermodynamics
● Measure of disorder or randomness in a system
○ S = Kln(V)
■ K = Boltzmann Constant or 1.38 * 10-23 J/K
■ V = Volume
● Can also be defined as change when energy is transferred at a constant temperature
○ ΔS= Q/T
■ ΔS= Change in Entropy
■ Q = energy
■ T = Constant Temperature