Chapter 7 Flashcards
System vs Surroundings
System: Uses up energy
Surroundings: Providing energy for function
Isothermal
No change in temperature
∆U = 0
Q = W
Adiabatic
No heat exchange
Q = 0
∆U = -W
Isobaric
No pressure change
Flat line in P-V graph
Standard Conditions
- Used in kinetics, equilibrium, and thermodynamic calculations
- 25˚C (298 K)
- 1 atm pressure
- 1 M concentration
State Function
Properties of a system at equilibrium
* Independent of path taken to achieve equilibrium
* May be dependent on each other
Examples: pressure (P), density (p), temperature (T), volume (V), enthalpy (H), internal energy (U), Gibbs free energy (G), entropy (S)
Process Functions
Define path between equilibrium states
* Include Q (heat) and W (work)
Triple Point
Combination of temperature and pressure where all three phases are at equilibrium
Critical Point
Temperature and pressure above liquid and gas phases are indistinguishable and heat of vaporization = 0
Temperature
Indirect measure of thermal content of a system that looks at average kinetic energy (KE) of particles in a sample
Heat
Thermal energy transferred between objects as a result of differences in their temperatures
Specific heat (c)
Energy required to raise temperature of 1 g of a substance by 1 ˚C
Heat Capacity (mc)
Energy required to raise any given amount of substance 1˚C
* Product of mass and specific heat
Constant-Pressure Caolrimeter
(AKA coffee cup calorimeter)
- Exposed to constant (atmospheric) pressure
- As reaction proceeds, temperature of contents measured to determine heat of reaction
Constant-Volume Calorimeter
(AKA bomb calorimeter)
- Heats of certain reactions (ex: combustion) can be measured indirectly by assessing temperature change in water bath around reaction vessel
Specific Heat of Water
1 cal/g * K
Unit: calories
Endothermic Reactions
Increase in heat content of system from surroundings
∆H > 0
Exothermic Reactions
Release of heat content from system
∆H < 0
Enthalpy of Reaction
Bonds broken - bonds formed
Phases of Matter Entropy
Solid < Liquid < Gas
Entropy (∆ S)
- Increases as sytem has more disorder/freedom of movement and energy dispersed in spontaneous system
- Entropy of universe can never be decreased spontaneously
Examples:
Freezing (decrease)
Sublimation (increase)
Dissolution (increase)
Fewer moles of gas (decrease)
Heat transferred (increase)
Deposition (largest decrease)
Gibbs Free Energy
∆ G = ∆H - T∆ S
Equilibrium
∆G = 0
Isovolumetric (AKA isochoric)
No change in volume
W = 0
∆U = Q
Standard Gibbs Free Energy form Equilibrium Constant
∆G˚rxn = -RTlnKeq
Combustion
Reaction of a hydrocarbon with O to produce CO2 and H2O
Value of e (from ln)
2.7
Gibbs Free Energy from Reaction Quotient
∆Grxn = ∆G˚rxn + RTlnQ = RTlnQ/Keq
Spontaneous Reaction
∆G < 0
If forward: Keq > Q
If backward: Keq < Q
Heat transfer no phase change
q = mc∆T
Heat transfer during phase change
q = mL
temperature is constant during phase change
Entropy Change
∆S = Qrev/T
Q rev = heat input
