Spontaneous Process
- Processes that are spontaneous in one direction are nonspontaneous in the reverse direction.
- Anything above 0°C is spontaneous for ice to melt.
- Below 0°C is spontaneous for the reverse process.
- Ex. A nail left outside will rust meaning it's reacting with oxygen from the air to form iron oxide; the rusting process is spontaneous and the reverse process is nonspontaneous. (See picture in slideshow)
- Spontaneity depends on the temperature (T):
- If T>0°C ice melts spontaneously to liquid.
- If T<0°C reverse process, water freezing to ice.
- If T=0°C the two states are in equilibrium.
Laws of Thermodynamics
First Law of Thermodynamics:
- Energy can be moved between a system and its surrounding, but the total energy stays the same.
- ΔE = q + w
- ΔE: change in the internal energy of a system.
- q: absorbed heat by system from surroundings.
- w: work done on system by surroundings.
- The total amount of energy lost in a system is the same as the total amount of energy gained by its surrounding.
- Entropy increases in a spontaneous process.
- Reversible process: ΔS(univ) = ΔS(system) + ΔS(surroundings) = 0
- Irreversible process: ΔS(univ) = ΔS(system) + ΔS(surroundings) > 0
- Entropy cannot be a negative!!
- Absolute zero: S(0 K) = 0
- 0° Kelvin = -273°C
- One microstate: S = k ln W = k ln 1 = 0
- The order of the phases of a given substance: S(solid) < S(liquid) < S(gas)
Entropy Change
- When entropy increases:
- Liquids and solids form gases.
- Liquids or solutions formed from solids.
- Number of gas molecules and moles increases.
- Standard Molar Entropy: 298 K (room temp.)
- ΔS: Entropy Change
- System depends on the initial and final states of system: ΔS = S(final) - S(initial)
- Isothermal Process: ΔS is equal to the heat that's transferred if reversible then divided by the temperature: ΔS = q(rev)/T
- q(rev) = ΔH(fusion)
- Ex. Enthalpy of fusion for H2O is ΔH(fusion) = 6.01 kJ/mol. Calculate the ΔS(fusion) for melting one mole of ice at 273 K.
- To solve first convert 6.01 kJ/mol to J/mol (divide by 1000)
- After converting do multiplication (so multiply 1 mol to 6.01 x 10^3 J/mol)
- Then divide that by the temperature and get your answer
- Entropy Change in Universe: ΔS(universe) = ΔS(system) + ΔS(surroundings)
- To be spontaneous the ΔS(universe) > 0
- ΔS(surroundings) = -ΔH(system)/T
- ΔS(universe) = ΔS(system) + -ΔH(system)/T (A reaction occurring at constant temp. & pressure.)
- -TΔS(univ) = ΔH(sys) - TΔS(sys) (multiply both sides by (-T) to get rid of fraction)
Gibbs Free Energy
- Gibbs free energy of state: G = H - TS
- G: Gibbs free energy
- H: Enthalpy
- T: Temperature
- S: Entropy
- Standard Gibbs Free Energy Change: ΔG = ΔH - TΔS
- If ΔG < 0, the reaction is spontaneous in a forward direction.
- If ΔG = 0, the reaction is at equilibrium.
- If ΔG > 0, the reaction in forward direction is nonspontaneous
Standard Free Energy Changes:
Free Energy & Equilibrium Constant
- ΔG = ΔG° + RT ln Q
- ΔG = 0, Q = K; 0 = ΔG° + RT ln K
- ΔG°= -RT ln K (subtract ΔG° to both sides then divide both sides by -1.)
- R: Ideal gas constant, 8.314 J/(mol)(K)
- T: Absolute Temperature
- Q: Reaction quotient