OCN 623: Thermodynamic Laws & Gibbs Free Energy, Slides of Thermodynamics

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OCN 623: Thermodynamic Laws

& Gibbs Free Energy

or: How to predict chemical reactions without doing experiments

Outline

• Definitions
• Laws of Thermodynamics
  • Enthalpy and entropy
• Gibbs free energy
  • Predicting the feasibility of reactions

Intensive properties

• Do not depend on quantity or mass
• Examples:
  • Temperature
  • Pressure
  • Density
  • Refractive index Definitions

Reversible and irreversible

processes

• Reversible process occurs under equilibrium conditions e.g. gas expanding against a piston p P p = P + ∂p reversible p = P + ∆p irreversible p = P + δp: reversible p = P + Δp: irreversible Definitions

Spontaneous processes

• Occur without external assistance
• Are irreversible by definition - occur at a finite rate
• Some examples:
  • expansion of a gas from region of high pressure to low pressure
  • diffusion of a solute in a solvent Definitions

Zeroth Law of Thermodynamics

• Two systems that are in thermal equilibrium with a third system are in thermal equilibrium with each other Laws of Thermodynamics

• U is a thermodynamic function
  • dU depends only on the initial and final states of the system, not on the path taken
• q and w are NOT thermodynamic functions Internal energy of system is increased by gaining heat (q) Internal energy of system is decreased when work is done by the system dU = dq - dw Laws of Thermodynamics

Work

Work of expansion: w exp

= P∆V

where P = pressure, ∆V = change in volume At constant volume, w = 0: no work is done Laws of Thermodynamics

Electrical work

• Occurs in electrical cells e.g. Zn + Cu 2+ = Zn 2+ + Cu
• Electrical energy = I * E * t E = emf (voltage), t = time, I = current flowing It = zF (for 1 mol) z = # of electrons transferred F (Faraday) = 96,490 Coulombs/mol Electrical energy = zEF Laws of Thermodynamics

• System at constant volume (all non-gas reactions are at constant volume) P∆V = w = 0 ∆U = q - w = q v q v = heat at constant volume
• System at constant pressure (all reactions open to atmosphere) ∆U = q p - P∆V
• Ideal gas law: P∆V = ∆nRT R is the gas constant = 8.314 Joules mol - K - Therefore: ∆U = q p - ∆nRT Rearranging: q p = ∆U + ∆nRT Laws of Thermodynamics

• q p is called the enthalpy H (constant pressure)
• Change in enthalpy: ∆H = ∆U + P∆V
• In absolute terms: H = U + PV
• H is a thermodynamic property, defined in terms of thermodynamic functions: U, P and V For an infinitesimal change at constant pressure: dH = dU + PdV

Enthalpy

Laws of Thermodynamics

• ∆H

f 0 298 is the heat of formation of 1 mole of a compound from its elements at 298 o K

• ∆H < 0 = exothermic reaction
• ∆H > 0 = endothermic reaction (seen from system perspective)
• ∆H is proportional to amount of material is an extensive property Laws of Thermodynamics

• Calculate enthalpy change for unknown reaction Reaction Enthalpy (a) H 2 + 0.5 O 2 = H 2 O ∆H = -285.8 kJ mol - (b) C + O 2 = CO 2 ∆H = -393.3 kJ mol - (c) C 2 H 6 + 3.5 O 2 = 2 CO 2 + 3H 2 O ∆H = -1559.8 kJ mol
• Calculate ∆H for ethane formation 2(b) + 3(a) - (c) Canceling yields: 2C + 3H 2 = C 2 H 6 ∆H = 2(-393.3) + 3(-285.8) - (-1559.8) = -84.2 kJ mol - 2C + 2O 2

• 3H 2
• 1.5O 2

• C 2 H 6
• 3.5O 2 = 2CO 2

• 3H 2 O - 2CO 2

• 3H 2 O Laws of Thermodynamics

• But enthalpy change alone is insufficient to allow prediction of likelihood of reaction
• Entropy change is also needed Laws of Thermodynamics