Chapter 19 (Frree energy & thermodynamics)

First law of thermodynamics

Energy cannot be created or destroyed; only transferred or converted.

Internal energy (E)

The total energy of a system, including kinetic and potential energies of particles.

Change in internal energy (ΔE)

ΔE = q + w, where q is heat and w is work.

q positive (q > 0)

Heat absorbed by the system; endothermic.

q negative (q < 0)

Heat released by the system; exothermic.

w positive (w > 0)

Work done ON the system (compression).

w negative (w < 0)

Work done BY the system (expansion).

Enthalpy (H)

Heat content of a system at constant pressure.

Enthalpy change (ΔH)

ΔH = qp for processes at constant pressure.

Endothermic reaction

Absorbs heat; ΔH > 0.

Exothermic reaction

Releases heat; ΔH < 0.

Entropy (S)

Measure of disorder or randomness in a system.

Second law of thermodynamics

Spontaneous processes increase the total entropy of the universe.

Entropy change of the universe

ΔSuniv = ΔSsystem + ΔSsurroundings.

Spontaneous process

Occurs naturally; ΔSuniv > 0.

Nonspontaneous process

Requires external energy; ΔSuniv < 0.

Factors that increase entropy

Increased temperature, phase changes (solid → liquid → gas), more moles of gas, more disorder.

Entropy trend: solids vs liquids vs gases

Gases have highest entropy; solids have lowest.

Entropy and number of particles

More particles = higher entropy.

Standard entropy (S°)

Entropy of a substance at standard conditions (1 atm, 25°C).

Gibbs free energy (G)

Determines spontaneity: G = H – TS.

Gibbs free energy change (ΔG)

ΔG = ΔH – TΔS.

Spontaneous condition

ΔG < 0.

Nonspontaneous condition

ΔG > 0.

Equilibrium condition

ΔG = 0.

Temperature effect on spontaneity

If ΔH and ΔS have same sign, temperature determines spontaneity.

Spontaneity when ΔH < 0 and ΔS > 0

Always spontaneous (ΔG always negative).

Spontaneity when ΔH > 0 and ΔS < 0

Never spontaneous (ΔG always positive).

Spontaneity when ΔH < 0 and ΔS < 0

Spontaneous at low temperatures.

Spontaneity when ΔH > 0 and ΔS > 0

Spontaneous at high temperatures.

Standard free energy and equilibrium

ΔG° = –RT ln K.

Relationship between ΔG and K

If K > 1, ΔG° is negative; if K < 1, ΔG° is positive.

K and spontaneity

K > 1 favors products (spontaneous); K < 1 favors reactants.

Third law of thermodynamics

Entropy of a perfect crystal at absolute zero is 0.

ΔS surroundings formula

ΔSsurroundings = –ΔHsystem / T.

Hess’s law

Total enthalpy change is sum of individual step changes.

Standard enthalpy of formation (ΔHf°)

ΔH° when 1 mole of compound forms from its elements in their standard states.

ΔHrxn from enthalpies of formation

ΔHrxn = ΣnpΔHf°(products) – ΣnrΔHf°(reactants).

Entropy change sign for phase changes

Melting, vaporization, sublimation → ΔS > 0; freezing, condensation → ΔS < 0.

Free energy at equilibrium

ΔG = 0 and Q = K.

Relationship between ΔG and Q

ΔG = ΔG° + RT ln(Q).