Physical Pharmacy I Lecture 9

Study Notes

Thermodynamic equilibrium is a state where system properties don't change over time.
A reversible process is one where, if undone, the original equilibrium is restored. Crucially, no energy is lost (e.g., due to friction) during the process.
A reversible process is an ideal process; it is practically impossible to achieve in real-world scenarios.
A reversible process yields the maximum amount of work. Any irreversible process will produce less work due to energy losses, like friction.
Examples of reversible processes: Increasing pressure infinitesimally slowly on a gas to convert all input work into internal energy.

When an energy is permanently lost during a process instead of converted to internal energy, then the process is irreversible.
Converting mechanical work into frictional heat is an example of an irreversible process; there's no way to reverse this and get the heat back.
All real processes are irreversible.

Some processes happen spontaneously, while others don't. For example, objects fall spontaneously, but lifting requires external energy.
A spontaneous process occurs naturally, without outside intervention. Examples: diamond turning into graphite, heat flowing from a hot object to a cold object.
A non-spontaneous process does not occur naturally and requires external input; throwing objects up is an example.
All spontaneous processes are irreversible.

The first law of thermodynamics states that energy is conserved when converted from one form to another.
Historically, a negative enthalpy change (ΔH) was thought to be a guarantee of spontaneity. This is incorrect.
While many natural reactions do have negative enthalpy change melting of ice is an example of a reaction where a negative enthalpy change is not a guarantee of spontaneity; the melting of ice at room temperature requires an input of energy and the change in enthalpy is positive.

The inability to predict the direction of purely energy-based processes means another state function is needed.
The second law uses a state function called entropy change (ΔS).
Entropy represents the probability of a process occurring and the tendency of a system to reach energy equilibrium.

In spontaneous processes, thermal energy spreads out. For example, gas molecules disperse throughout a container. This is a high energy distribution (large entropy).
Entropy is a measure of energy dispersion and sharing within a system.
As thermal energy spreading increases, entropy increases.

The entropy change (ΔS) in an isothermal reversible process is equal to the heat change (Q_rev) divided by the absolute temperature (T). (ΔS = Q_rev/T).
The units of entropy are energy per degree Kelvin (e.g., J/K or cal/K).
Entropy change only depends on a system's starting and ending states, not how it gets there, making it a thermodynamic state function.

The total entropy change for the entire system (including surroundings) of a spontaneous process is always positive. (ΔS_total> 0)
For a system at equilibrium, the total entropy change is zero. (ΔS_total = 0)
For a reversible process, the total entropy change is zero.
For an irreversible process, the total entropy change is greater than zero.

The entropy of a perfect crystal at absolute zero (0 K) is zero because the crystals have maximum order at absolute zero.
The third law deals with an ideal state (0K); practically impossible to achieve.

The third law helps establish a scale for measuring absolute entropy (S) making it possible to calculate absolute entropies of pure substances.

The absolute entropy of a substance at a given temperature is the sum of all the entropy it would acquire on warming from absolute zero to the specific temperature.

Standard molar entropy (S°) is the entropy of one mole of a substance at standard conditions. (Units: J K^-1 mol^-1).
Standard entropies can be calculated or measured; they are always greater than zero.
For similar substances, gas entropy > liquid entropy > solid entropy.
An increase in molecular weight often leads to an increase in standard molar entropy. Examples given, including CH₄, C₂H₆, and C₃H₈.

Chemical changes occur due to two factors:
Minimizing energy (ΔH)
Maximizing entropy (ΔS)
Gibbs free energy (ΔG) is a state function that combines the first and second laws to determine the direction of a chemical change.
ΔG = ΔH – TΔS (where T is temperature)

If ΔG is negative, the process is spontaneous.
If ΔG is zero, the system is at equilibrium.
If ΔG is positive, the process is non-spontaneous.
A highly negative ΔH and a highly positive ΔS favor a spontaneous reaction because ΔG would be a large negative value.

Phase change (ice to liquid water) calculation showing ice to water is spontaneous using enthalpy and entropy values at 25°C and 1 atm by using the ΔG equation.

Phase change (liquid water to ice) calculation showing ice formation at -10°C and 1 atm is spontaneous by using the ΔG equation.

Thermodynamic concepts explain drug diffusion across biological membranes as spontaneous processes.
Enthalpy is used to understand drug dissolution in solvents.
Entropy is a measure of energy unavailable for work, and higher entropy generally indicates the system is closer to equilibrium.