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
Endothermic Reactions
Increase in heat content of system from surroundings
∆H > 0
Exothermic Reactions
Release of heat content from system
∆H < 0
