Physical Chemistry Lesson of the Day – Standard Heats of Formation

The standard heat of formation, ΔHfº, of a chemical is the amount of heat absorbed or released from the formation of 1 mole of that chemical at 25 degrees Celsius and 1 bar from its elements in their standard states.  An element is in its standard state if it is in its most stable form and physical state (solid, liquid or gas) at 25 degrees Celsius and 1 bar.

For example, the standard heat of formation for carbon dioxide involves oxygen and carbon as the reactants.  Oxygen is most stable as O2 gas molecules, whereas carbon is most stable as solid graphite.  (Graphite is more stable than diamond under standard conditions.)

To phrase the definition in another way, the standard heat of formation is a special type of standard heat of reaction; the reaction is the formation of 1 mole of a chemical from its elements in their standard states under standard conditions.  The standard heat of formation is also called the standard enthalpy of formation (even though it really is a change in enthalpy).

By definition, the formation of an element from itself would yield no change in enthalpy, so the standard heat of reaction for all elements is zero.

 

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 – Heat Capacity

The heat capacity of a system is the amount of heat required to increase the temperature of the system by 1 degree.  Heat is measured in joules (J) in the SI system, and heat capacity is dependent on each substance.  To make heat capacities comparable between substances, molar heat capacity or specific heat capacity are often used.

  • Molar heat capacity is the amount of heat required to increase the temperature of 1 mole of a substance by 1 degree.
  • Specific heat capacity is the amount of heat required to increase the temperature of 1 gram of a substance by 1 degree.

For example, over the range 0 to 100 degrees Celsius (or 273.15 to 373.15 degrees Kelvin), 4.18 J of heat on average is required to increase the temperature of 1 gram of water by 1 degree Kelvin.  Thus, the average specific heat capacity of water in that temperature range is 4.18 J/(g·K).

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 – 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.