Chapter 7- Thermodynamics Flashcards
Three types of systems
Isolated- no matter or energy exchanged
Closed- exchanges only energy
Open- exhanges both matter and energy
Isothermal vs Adiabatic vs Isobaric vs Isovolumetric
Isothermal- constant temperature
Adiabatic- exchange no heat with environment
Isobaric- constant pressure
Isovolumetric (isochoric) constant volume
State functions and examples
Describe physical properties of an equilibrium state, pathway dependent
Pressure, Density, Temperature, Volume, Enthalpy (H), internal energy (U), Gibbs free energy, Entropy (S)
Standard conditions
When do you use it?
Temperature: 298 K
Pressure: 1 atm
Concentration: 1M
When measuring Enthalpy Entropy Gibbs free energy (prime)
STP
When do you use it?
Standard temperature and pressure
T- 0 C, 273 K
P- 1 atm
Ideal gas problems
Triple point
All three phases exist in equilibrium
Critical point
When liquid and gas phases are indistinguishable
Sublimation
Solid to gas
Deposition
Gas to solid
Condensation
Gas to liquid
Vaporization
Liquid to gas
Temperature vs heat
Temperature: average kinetic energy of a substance
Heat: transfer of energy results in difference in temperature
Enthalpy
Measure of potential energy of a system found in intermolecular attractions and chemical bonds
Hess’s Law
The total change in potential energy in system is equal to the changes of potential energies of the individual steps of the process.
q=mc🔺t
mass in grams/moles?
Used in between phase changes
Mass in grams
q=ml
Mass in grams or moles?
Used when temperature is constant
Mass in mol
q=ml units
M= mass in mol
L= latent heat kj/mol (heat of fusion, heat of vaporization)
q=mc🔺t units
m= mass in grams
c= specific heat J/g x K
T=K
Specific heat of water
1 cal/g x K
How to calculate the enthalpy change of a reaction?
🔺Hrxn = H products - H reactions
Endothermic vs Exothermic 🔺H
endothermic positive delta H
exothermic negative delta H
Bond breaking absorbs or releases energy
Absorbs (endothermic)
Bond making absorbs or releases energy
Releases energy (exothermic)
Second law of thermodynamics
Energy spontaneously disperses from being localized to becoming spread out if it is not hindered from doing so. 
