IB Chemistry HL 100113

Entropy, $S$, is a measure of the dispersal or distribution of matter and/or energy in a system. The more ways the energy can be distributed, the higher the entropy. Under the same conditions, the entropy of a gas is greater than that of a liquid, which in turn is greater than that of a solid.

Predict whether a physical or chemical change will result in an increase or decrease in entropy of a system.

Calculate standard entropy changes, $ΔS^{⦵}$, from standard entropy values, $S^{⦵}$.

Standard entropy values are given in the data booklet.

Change in Gibbs energy, $ΔG$, relates the energy that can be obtained from a chemical reaction to the change in enthalpy, $ΔH$, change in entropy, $ΔS$, and absolute temperature, $T$.

Apply the equation $ΔG^{⦵} = ΔH^{⦵} − TΔS^{⦵}$ to calculate unknown values of these terms.

Thermodynamic data values are given in the data booklet.

Note the units: $ΔH$ kJ mol^-1; $ΔS$ J K^-1 mol^-1; $ΔG$ kJ mol^-1.

At constant pressure, a change is spontaneous if the change in Gibbs energy, $ΔG$, is negative.

Interpret the sign of $ΔG$ calculated from thermodynamic data.

Determine the temperature at which a reaction becomes spontaneous.

$ΔG$ takes into account the direct entropy change resulting from the transformation of the chemicals and the indirect entropy change of the surroundings resulting from the transfer of heat energy.

As a reaction approaches equilibrium, $ΔG$ becomes less negative and finally reaches zero.

Perform calculations using the equation $ΔG = ΔG^{⦵} + RT \ln{Q}$ and its application to a system at equilibrium $ΔG^{⦵} = −RT \ln{K}$.

The equations are given in the data booklet.