## Physical Chemistry Lesson of the Day – State Functions vs. Path Functions

Today’s lesson may seem mundane; despite its subtlety, it is actually quite important.  I needed to spend some time to learn it and digest it, and it was time well spent – these concepts are essential for understanding much of thermodynamics.  For brevity, I have not dived into the detailed mathematics of exact differentials, though I highly recommend you to learn it and review the necessary calculus.

Some thermodynamic properties of a system can be described by state variables, while others can be described by path variables.

A state variable is a variable that depends only on the final and initial states of a system and not on the path connecting these states.  Internal energy and enthalpy are examples of state functions.  For example, in a previous post on the First Law of Thermodynamics, I defined the change in internal energy, $\Delta U$, as

$\Delta U = \int_{i}^{f} dU = U_f - U_i$.

State variables can be calculated by exact differentials.

A path variable is a variable that depends on the sequence of steps that takes the system from the initial state to the final state.  This sequence of steps is called the path.  Heat and work are examples of path variables.  Path variables cannot be calculated by exact differentials.  In fact, the following quantities may seem to have plausible interpretations, but they actually do not exist:

• change in heat ($\Delta q$)
• initial heat ($q_i$)
• final heat ($q_f$)
• change in work ($\Delta w$)
• initial work ($w_i$)
• final work ($w_f$)

There is no such thing as heat or work being possessed by a system.  Heat and work can be transferred between the system and the surroundings, but the end result is an increase or decrease in internal energy; neither the system or the surroundings possesses heat or work.

A state/path variable is also often called a state/path function or a state/path quantity.

## Physical Chemistry Lesson of the Day – The Difference Between Changes in Enthalpy and Changes in Internal Energy

Let’s examine the difference between a change enthalpy and a change in internal energy.  It helps to think of the following 2 scenarios.

• If the chemical reaction releases a gas but occurs at constant volume, then there is no pressure-volume work.  The only way for energy to be transferred between the system and the surroundings is through heat.  An example of a system under constant volume is a bomb calorimeter.  In this case,

$\Delta H = \Delta U + P \Delta V = \Delta U + 0 = q - w + 0 = q - 0 + 0 = q$

This heat is denoted as $q_v$ to indicate that this is heat transferred under constant volume.  In this case, the change in enthalpy is the same as the change in internal energy.

• If the chemical reaction releases a gas and occurs at constant pressure, then energy can be transferred between the system and the surroundings through heat and/or work.  Thus,

$\Delta H = \Delta U + P \Delta V = q - w + P \Delta V = q$

This heat is denoted as $q_p$ to indicate that this is heat transferred under constant pressure.  Thus, as the gas forms inside the cylinder, the piston pushes against the constant pressure that the atmosphere exerts on it.  The total energy released by the chemical reaction allows some energy to be used for the pressure-volume work, with the remaining energy being released via heat.  (Recall that these are the 2 ways for internal energy to be changed according to the First Law of Thermodynamics.)  Thus, the difference between enthalpy and internal energy arises under constant pressure – the difference is the pressure-volume work.

Reactions under constant pressure are often illustrated by a reaction that releases a gas in cylinder with a movable piston, but they are actually quite common.  In fact, in chemistry, reactions under constant pressure are much more common than reactions under constant volume.  Chemical reactions often happen in beakers, flasks or any container open to the constant pressure of the atmosphere.

## Physical Chemistry Lesson of the Day – Enthalpy

The enthalpy of a system is the system’s internal energy plus the product of the pressure and the volume of the system.

$H = U + PV$.

Just like internal energy, the enthalpy of a system cannot be measured, but a change in enthalpy can be measured.  Suppose that the only type of work that can be performed on the system is pressure-volume work; this is a realistic assumption in many chemical reactions that occur in a beaker, a flask, or any container that is open to the constant pressure of the atmosphere.  Then, the change in enthalpy of a system is the change in internal energy plus the pressure-volume work done on the system.

$\Delta H = \Delta U + P\Delta V$.

## Physical Chemistry Lesson of the Day: Pressure-Volume Work

In chemistry, a common type of work is the expansion or compression of a gas under constant pressure.  Recall from physics that pressure is defined as force applied per unit of area.

$P = F \div A$

$P \times A = F$

Consider a chemical reaction that releases a gas as its product inside a sealed cylinder with a movable piston.

Image from Dpumroy via Wikimedia.

As the gas expands inside the cylinder, it pushes against the piston, and work is done by the system against the surroundings.  The atmospheric pressure on the cylinder remains constant while the cylinder expands, and the volume of the cylinder increases as a result.  The volume of the cylinder at any given point is the area of the piston times the length of the cylinder.  The change in volume is equal to the area of the piston times the distance along which the piston was pushed by the expanding gas.

$w = -P \times \Delta V$

$w = -P \times A \times \Delta L$

$w = -F \times \Delta L$

Note that this last line is just the definition of work under constant force in the same direction as the displacement, multiplied by the negative sign to follow the sign convention in chemistry.

## Physical Chemistry Lesson of the Day – The First Law of Thermodynamics

The change in internal energy of a system is defined to be the internal energy of a system in its final state subtracted by the internal energy of the system in its initial state.

$\Delta U = U_{final} - U_{initial}$.

However, since we cannot measure the internal energy of a system directly at any point in time, how can we calculate the change in internal energy?

The First Law of Thermodynamics states that any change in the internal energy of a system is equal to the heat absorbed the system plus any work done on the system.  Mathematically,

$\Delta U = q + w$.

Recall that I am using the sign convention in chemistry.

The value of $q$ and $w$ can be positive or negative.

• A negative $q$ denotes heat released by the system.
• A negative $w$ denotes work done by the system.

## Physical Chemistry Lesson of the Day – Basic Terminology in Thermodynamics

A system is the part of the universe of interest, and the surroundings is everything else in the universe.

The internal energy of a system is the sum of the kinetic and potential energies of all of the particles (atoms and molecules) in the system.  This cannot be measured, but changes in internal energy can be measured.

There are 2 ways in which the internal energy of a system can change: heat and work.

• Heat is the transfer of energy between 2 objects due to a temperature difference.  In chemistry, heat is commonly observed when a chemical reaction absorbs or releases energy.
• Work is force acting over a distance.  In chemistry, a common type of work is the expansion or compression of a gas.

In chemistry, it is conventional to take the system’s point of view in deciding the sign of heat and work.  Thus, if heat is entering the system or if work is done on the system, then the sign is positive.  If heat is exiting the system of if work is done by the system, then the sign is negative.