Thermodynamics

1st law of thermodynamics

energy cannot be created or destroyed - only changed from one form to another

how are system and surroundings related

delta Esystem = - delta Esurroundings

internal energy U

energy associated with random disordered motion of molecules in a system

how can U be changed

heat

work

work

w = -P deltaV = -P(Vf-Vi)

transfer of energy from system to surroundings

work = pos

surroundings are doing work on our system

i.e. system gains energy

work = neg

system doing work on surrounding

i.e. system loses energy

entropy S

a measure of how dispersed the energy is inside a closed or isolated system

higher S

more energy dispersed

lower S

energy less dispersed

second law of thermodynamics

entropy S of an isolated system always increases

deltaS>0

ex: ice cube melting

is the inital or final state of an isolated container more random

final state is less ordered and more random

system has become more disordered and entropy has increased

can we revert back to original state

no unless we do work on the system

reversible process

both system and surrounding can be reverted to original state after process has occured

irreversible process

system and surrounding cannot be reverted to original state after process has occured

how can we measure change of entropy delta S for a reversible process

delta S = qrev/T

q rev

heat supplied reversible to the system

microstate

when we look at a molecular system we want to know the number of diff ways molecules can be arranged

each individual arrangement = microstate

how is microstate and entropy related

greater no. of microstates = greater entropy

S = kB ln(W)

W

no. of microstates

kB

boltzman constant

-> 1.38 x 10^-23

how can we increase no. of microstates

increase volume

add heat

increase no. of molecules

change in entropy of a system from one state to another

delta S = kB ln(W final) - kB ln(W initial)

enthalpy H

heat content of a system at constant pressure

H = U + PV

change in enthalpy delta H

amount of heat a system loses or gains

delta H = delta U + P delta V

exothermic reaction

delta H is negative

endothermic reaction

delta H is positive

enthalpy vs internal energy

IE: sum of all molecular energies in a system

E: combo of internal energies + energy required to make space

so when heat is added enthalpy changes stays the same but if gas can expand some energy goes into work so internal energy increases less

total entropy change of the universe

ΔSuniverse=ΔSsystem+ΔSsurroundings

entropy change of surroundings

ΔSsurroundings = −ΔHsystem/T

Gibbs free energy of a system

∆𝑮𝒔𝒚𝒔𝒕𝒆𝒎 = ∆𝑯𝒔𝒚𝒔𝒕𝒆𝒎 − 𝑻∆𝑺𝒔𝒚𝒔𝒕𝒆m

neg delta G

spontaneous

pos delta G

not spontaneous

delta G = 0

equilibrium

Clausius-Clapeyron equation

OR

P2 = P1e

equation for gibbs free energy under non-standard conditions

𝚫𝑮 = 𝚫𝑮 𝟎 + 𝑹𝑻𝒍𝒏 (Q)

What is the relationship between ΔG° and K

𝚫𝑮 𝟎 = −𝑹𝑻𝒍𝒏 (K)

If products are favoured at equilibrium

K>1

delta G0 < 0

reactants are favoured at equilibrium

k < 1

delta G0 > 0

neither reactants or products are favoured

K = 1

delta G0 = 0