Chemical+Thermodynamics

Chemical Thermodynamics

Spontaneous Processes

What’s it mean to be spontaneous?? A spontaneous reaction is one which goes to completion, and is irreversible, without continuos input of energy. Ex: combustion

Reversible vs Irreversible processes. A reversible reaction can achieve equilibrium and is not spntaeous.

An irreversible reaction such as combustion only goes in one direction and can be spontaneous.

Entropy and the Second Law of Thermodynamics Nature loves low energy stable molecules and favors reactions which produce such.

Entropy (S) Entropy is chaos, a measure of disorder.

Changes in entropy and the second law ΔS**universe** = ΔS **system -** ΔS**surroundings** with every spontaneous reaction the universe edges closer to total chaos

Entropy on the molecular level and the Third law S > 0 for all substances except a perfect crystal at 0 K

__Amount of Entropy__ s < l < < < < g s + l < aq

Entropy changes in Chemical Reactions

Standard Molar Entropies and Delta S

Example:
 * ΔS** = **ΣS**products - ΣSreactants

3NO2 + H2O > 2HNO3 + NO


 * ΔS = (502.62) - (791.26) = -288.64**

Gibbs Free Energy

First of all let’s review the changes that are favored by nature.

Signs favored by nature: "-" delta H (exothermic) "+" delta S (increase in entropy)

Enthalpy and Entropy

If a reaction is spontaneous then the energy involved may be “free” to do work. This is called the Free Energy.

"-" delta G (change in Gibbs Free Energy)

Change in Free Energy is most often calculated. (delta G)

delta G **__MUST__** **__BE__** negative for the reaction to be spontaneous and to be able to release free energy.

2 ways to calculate it:

1.) Gibbs- Helmholtz equation.

delta G = delta H - T(delta S)

If delta G in negative, then the reaction is not spontaneous If delta G is positive, it is spontaneous if delta G is 0 then the reaction is at equilibrium

2.) Standard Free energies of formation(limited to 25oC)

This information can be found in Appendix C in the back of the book. It is also in "bookie"...Mr. Williams chart

Free Energy and Temperature (see table 19.4).

delta H delta S delta G Example


 * - ||  +  ||  “-“ at all temp.  ||  Combustion, explosions, radioactive decay  ||
 * - ||  -  ||  “-“ only at “low” temps.  ||  Freezing, decomposition  ||
 * + ||  -  ||  + never “-“  ||  Not common at all  ||
 * + ||  +  ||  “-“ only at “high” temps.  ||  Melting, boiling, sublimation  ||

JS

Free Energy and the Equilibrium Constant

∆ G = ∆ Go + RT ln Q

Recall Q is just like K only not necessarily at equilibrium conditions.

Delta Go is delta G at standard conditions.

If we are at standard conditions pressures are 1 atm, concentrations are 1 M so Q = 1 and ln Q = 0. So ∆ G = ∆ Go… As you would expect.

When the concentrations and pressures are non-standard you must use the first equation to calculate ∆ G.

If Q = K. Then… the equation reverts to a very useful form:

∆ Go = - RT lnKp

Meaning:

If Kp is greater than 1, ∆ Go is negative

If Kp is less than 1, ∆ Go is positive

If Kp = 1 ∆ Go =0