1. System. surroundings. universe.
thermodynamic state of a system.
The number of moles and identity of each substance.
The physical states of each substance.
The temperature of the system.
The pressure of the system.
Thermochemical standard state conditions
The thermochemical standard T = K.
The thermochemical standard P = atm.
Be careful not to confuse these values with STP.
Thermochemical standard states of matter
For pure substances in their liquid or solid phase the standard state is the pure liquid or solid.
For gases the standard state is the gas at 1.00 atm of pressure. For gaseous mixtures the partial pressure must be 1.00 atm.
For aqueous solutions the standard state is 1.00 M concentration.
State Functions are independent of pathway:
T (temperature), P (pressure), V (volume), E (change in energy), H (change in enthalpy – the transfer of heat), and S (entropy)
Examples of non-state functions are:
n (moles), q (heat), w (work)
Entropy
Enthalpy
Gibbs Free Energy
Heat, Latent Heat, Sensible Heat
Energy
Internal energy
Kinetic Energy
Potential Energy
Endothermic
Exothermic
Thermodynamics
Thermal Equilibrium
System
Surroundings
Law of Conservation of Energy
Law of Conservation of Mass
Heat Capacity, Molar Heat Capacity
Specific Heat Capacity
First Law of Thermodynamics
State Function
Standard state temperature
Standard state pressure
Standard states matter
Hess’s Law
Closed System
Open System
Isolated System
Heat of Fusion
Heat of Vaporization

2. Specific Heat Capacity
How much energy is transferred due to Temperature difference?
The heat (q) “lost” or “gained” is related to
a) sample mass
b) change in T and
c) specific heat capacity
Specific heat capacity =
heat lost or gained by substance (J)
(mass,
g) (T change, K)

3. q cP = m(Tf –Ti) ∆H = Hfinal - Hinitial
The stoichiometric coefficients in thermochemical equations must be interpreted as numbers of moles. 1 mol of C5H12 reacts with 8 mol of O2 to produce 5 mol of CO2, 6 mol of H2O, and releasing 3523 kJ is referred to as one mole of reactions.
∆Horxn =  ∆Hfo (prod) -  ∆Hfo (react)
Specific heat capacity (J/(g∙K) =
heat lost or gained by system (Joules)
mass(grams) DT (Kelvins)
Variable
System 1
System 2 Cp Tf Ti m q
m(Tf –Ti) q
cP =

4. heat transfer out
(exothermic), -q
heat transfer in
(endothermic), +q
SYSTEM
∆E = q + w
w transfer in
(+w)
w transfer out
(-w)

5. There are two basic ideas of importance for thermodynamic systems Law of Conservation of mass and the Law of Conservation of Energy. The second leads to 3 thermodynamic laws.
First Law – the energy of the Universe is constant Chemical systems tend toward a state of minimum potential energy. The first law is also known as the Law of Conservation of Energy. Energy is neither created nor destroyed in chemical reactions and physical changes.
Esys +Esurr = Euniverse Esys = KEsys + PEsys
KE – kinetic energy: translational, rotational, vibrational
PE – energy stored in bonds
The second law of thermodynamics states, “In spontaneous changes the universe tends towards a state of greater disorder.” Chemical systems tend toward a state of maximum disorder. The entropy of universe must increase.
Fundamentally, the system must be capable of doing useful work on surroundings for a spontaneous process to occur. In general for a substance in its three states of matter:
Sgas > Sliquid > Ssolid
When:
S > 0 disorder increases (which favors spontaneity).
S < 0 disorder decreases (does not favor spontaneity).
The Third Law of Thermodynamics states, “The entropy of a pure, perfect, crystalline solid at 0 K is zero.”

6. DG = DH-TDS at constant T and P
The relationship describes the spontaneity of a system.
The relationship is a new state function, G, the Gibbs Free Energy.
Sign conventions for G.
G > 0 reaction is nonspontaneous
G = 0 system is at equilibrium
G < 0 reaction is spontaneous
heat transfer out
(exothermic), -q
heat transfer in
(endothermic), +q
SYSTEM
∆E = q + w
w transfer in
(+w)
Compression of system
w transfer out
(-w)
Expansion of system

7. Forward reaction Free energy has the relationship G = H -TS.
Because 0 ≤ H ≥ 0 and 0 ≤ S ≥ 0,
there are four possibilities for G.
Forward reaction
H S G spontaneity
< 0 > 0 < at all T’s.
< 0 < 0 T dependent at low T’s.
> 0 > 0 T dependent at high T’s.
> 0 < 0 > Nonspontaneous
at all T’s.

8. Kinetics Four factors that affect the rate of reaction
nature of reactant
concentration
temperature
presence of a catalyst
Rate of reaction
Rate constant
Order of reactant
Overall order of reaction
General rate expression
Initial rate
Instantaneous rate
Average rate
Integrated rate laws ([ ] with respect to time)
If zero order [A]0 - [A] = ak t
first order
second order
Differential rate laws ([ ]with respect to rate)
Rate = k [A]m[B]n[C]p
If zero order Rate = k[A]0 = k
first order Rate = k[A]1 = k[A]
second order Rate = k[A]2
Change in rate is independent of the change in concentration of a zero order reactant
Change is rate is directly proportional to the change in concentration of a 1st order reactant
Change in rate is directly proportional to the square of the change in concentration of a 2nd order reactant
1/[A] vs t
2nd order
1/[A] t
0 order
[A] vs t
[A] t
ln [A] vs t
1st order
ln [A] t

9. All equations are shown for a stoichiometric coefficient of 1
All equations are shown for a stoichiometric coefficient of 1. for all others use the term akt, or ak in place of kt and k below.
Five factors that affect the rate of reaction:
nature of reactant
Concentration
Temperature
presence of a catalyst
solvent effects
Rate of reaction
Rate constant
Order of reactant
Overall order of reaction
Substitute ½ [A]o for [A]
In integrated rate law equation
to get ½ life equation