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Open, closed, isolated system s
  • open: exchange both mass and energy with surroundings
  • closed: exchange only E
  • isolated: dont exchange either
state function
pathway independent

(NOT work or heat = path functions)
  • internal energy
  • temperature
  • pressure
  • volume
  • enthalpy
  • entropy
  • gibbs energy
forms of heat
  • conduction
  • convection
  • radiation
conduction
thermal energy transfer via molecular collisions

requires direct physical contact

rate of heat flow is constant across any number of slabs between 2 heat reservoirs (but conductivity results in temperature differences)
convection
thermal E transfer via fluid movements (differences in pressure or density drive warm fluid in direction of cooler fluid)

ex: ocean and air currents
radiation
thermal E transfer via electromagnetic waves (fast as light)

object that radiates heat faster also absorbs heat faster
PV Work
at constant pressure, w=PΔV

PV work takes place when a gas expands against a force.

if volume remains constant, no work is done at all.
Heat engines
second law of thermo: heat cannot be changed completely into work in a cyclical process.

heat entering engine equals net work done by engine + heat leaving engine.

further apart temperatures of two reservoirs, more effective the conversion.
internal energy
all forms of energy on a molecular scale (ex: vibrational, rotational, translational, electronic) (doesnt include mechanical energy)

MCAT may refer to internal energy as 'heat energy' 'thermal energy' or even 'heat' !!! dont get confused

ΔU = q (if no change in volume and no work)
temperature
measurement of how fast molecules are moving or vibrating

KEave=(3/2)kT

Virtually all physical properties change with temperature.
pressure
greater the random translational KE of gas molecules per volume, the greater the pressure.
enthalpy
H ≡ U + PV

interested in change:

ΔH=ΔU + PΔV (constant P)
standard state
! dont confuse with STP

an element in its standard state at 25C is arbitrarily assigned an enthalpy value of 0 K/mol
standard enthalpy of formation
change in enthalpy for a reaction that creates 1 mole of that compound from raw elements in standard state

ΔH=q

if gas i not part of the reaction, enthalpy change is equal to heat which (in absence of work) is equal to change in energy
Hess' law
"sum of enthalpy changes for each step is equal to the total enthalpy change regardless of path chosen"

add reactions, can add enthalpies

Hess' law works because enthalpy is a state function
-ΔH
exothermic, release heat making the reaction system hot.
+ΔH
endothermic, absorb heat making the reaction system cold.
transition state
where old bonds are breaking and new bonds are forming. (peak of energy hill)

! dont confuse with intermediate which are the products of the first step of a 2-step reaction
catalyst
lowers activation energy for both forward and reverse reactions --> equilibrium is unaffected by catalyst

catalyst affects the rate not equilibrium of a rxn
(!) increasing temp. increases rates of both forward and reverse reactions due to greater E to overcome activation energy
entropy

  • state function (forward = (-) of reverse)
  • entropy and not energy dictate direction of a reaction
  • (!) entropy of universe always increases in spontaneous reaction, entropy of system may or may not increase
  • entropy is extensive (increase with amount of substance) and increases with number, volume, and temperature.
  • only ideal reactions create zero change in entropy and are reversible
  • on macroscopic scale, all real rxns are irreversible
Gibbs free energy (G)
ΔG=ΔH-TΔS

-ΔG indicates a spontaneous reaction.

*refers to changes in system
0th law
two bodies in thermal equilibrium with the same system are in thermal equilibrium with each other (temperature exists and is a state function)
1st law
energy of an isolated system is conserved for any reaction
2nd law
entropy of the universe increases for any reaction
3rd law
perfect crystal at 0K is assigned an entropy value of 0. all other substances and all temperatures have a positive entropy value.
colloid
  • larger solute particles in solution
  • particles too small to be extracted by filtration but large enough/charged to be separated by semipermeable membrane
  • ex:hemoglobin, aerosol, fog, foam, emulsion
  • colloidal suspensions will scatter light
solvation
  • break apart into cations and anions and are surrounded by oppositely charged ends of polar solvent.
  • in water process is called hydration
  • when something is hydrated it is in an aqueous phase
parts per million
mass of solute per mass of solution times one million
vapor pressure
  • pressure created by molecules in open space when liquid and gas phases are in equilibrium.
  • vapor pressure increases with temperature
  • boiling occurs when vapor pressure of liquid equals atmospheric pressure
  • melting occurs when vapor pressure of solid phase = vapor pressure of liquid
nonvolatile solute
solute with no vapor pressure
Raoult's law (for nonvolatile solutes)
P=χP

if 97% of solution is solvent, then vapor pressure will be 97% of the vapor pressure of the pure solvent.

(due to competition for surface area of liquid)
Raoult's law for volatile solutes
P = χP + χP

If 97% of solution is solvent, then vapor pressure will be 97% of the vapor pressure of the pure solvent PLUS 3% of the vapor pressure of the pure solute.
deviations from Raoult's law with nonideal solutions
look at heats of solution.

negative heat of solution form stronger bonds and lower vapor pressure. (fewer molecules are able to break free from the surface)

positive heats of solution form weaker bonds and raise vapor pressure.
saturated solution
when rate of dissolution equals rate of precipitation (concentration of salt has reached maximum)
solubility vs. solubility product
  • solubility product is a constant that only changes with temperature
  • solubility is the maximum number of moles of solute that can dissolve in solution (depends upon common ions in soln) and depends on both the temperature and the ions in solution.
solubility factors
  • partial vapor pressure of a solution component is always proportional to its mole fraction
  • partial vapor pressure is proportional to pure vapor pressure
  • solubility of a gas is proportional to its vapor partial pressure
  • as temp. increases the solubility of salts increases
  • gases behave opposite: as temp. increases, solubility of gas decreases (soda can causing gas to rise out of solution as pressure released or is heated)
  • larger gases tend to experience greater van der waals forces and be more soluble
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