Hi there! So I know the previous notes that I’ve given you in this chemistry series have been difficult. They really are, even for me. That’s why teaching it and explaining it in my own words is so helpful for me, and why I’m being persistent in continuing the material. Thankfully, thermochemistry is easier and more intuitive than Chemical Kinetics (last week’s stuff).
Thermochemistry is all about studying the heat released and absorbed during chemical reactions. Read on!
Note: The study of how fast a reaction goes is kinetics. Thermodynamics do not tell you how long it will take for a reaction to go to completion!!
We study thermochemistry because it’s useful! When we do combustion reactions for fuel, we can understand how much energy a mole of a gas can release when burned in oxygen. When we eat food, we release energy in each step of the breakdown, which is used by our body for metabolic processes.
What’s more important, we can find what conditions will optimize the yields of products.
So, thermochemistry is a subset of thermodynamics – it deals with forms of energy and their interconversions.
The system is some portion of the universe arbitrarily chosen for consideration (i.e. a beaker where the reaction is taking place).
The surroundings are everything in the universe not included in the system.
Universe = system + surroundings
This “equation” always cracks me up when I see it. What a simple, idealistic way of viewing the world!
Now, a system can be (1) open, (2) closed, or (3) isolated.
- An open system allows for the exchange of heat, matter, and work.
- A closed system allows exchange of work and heat but not matter.
- An isolated system does not allow the exchange of matter, work, or heat.
Since in thermochemistry, we’re dealing with energy, what exactly is it? It’s a rather abstract concept – unlike matter, it can’t be seen and there are different forms. Basically it’s the capacity to do work.
Potential and kinetic energy are the two main forms of energy. I see thermal energy as a subset of kinetic energy, because it’s the energy associated with the motion of atoms and molecules, and chemical as a form of potential.
The First Law of Thermodynamics (Law of Conservation of Energy)
Energy can be converted from one form into another, but it cannot be created or destroyed. The total quantity of energy in the universe is thus a constant. This observation is the law of conservation of energy, the First Law of thermodynamics.
In the form of an equation, the change in internal energy for a closed system is equal to the heat plus work:
ΔE = E final – E initial = q + w
So, heat and work are two ways of increasing the internal energy of the system.
ΔE refers to the change in internal energy, E. It’s the sum of the kinetic and potential energies of the individual particles of a system.
Heat (q) is the energy transfer resulting from a temperature difference.
Work is the action of a force through a distance (force × distance). Here, w is the work done on the system.
Sign convention: Heat absorbed by a system is positive; heat released by a system is negative. If work is done on a system, work is positive; if work is done by the system, work is negative.
So, based on the first law of thermodynamics, in an isolated system, ΔE should be zero.
The thermodynamic property (state function) is a property whose change depends only on the initial and final states of the system.
Examples of state functions include energy, enthalpy, entropy, Gibbs free energy, pressure, volume, and temperature.
Examples of properties that are NOT state functions include heat and work.
All right. The heat of reaction is the heat released or absorbed during a chemical reaction.
This can be either exothermic – it gives off heat to the surroundings – or endothermic – it absorbs heat from the surroundings.
Note: During endothermic reactions (especially in isolated systems), the temperature of the system decreases because the reaction uses the energy abstracted from the reactants to form products.
These chemical reactions are conducted in either constant volume (isochoric) or constant pressure (isobaric) conditions.
Most (99%) of the reactions we do for educational purposes are isobaric – the pressure doesn’t change, while the volume of the reacting system usually changes. To do a isochoric reaction, most commonly we use bomb calorimeters, which allows the reaction to take place in a sealed container with rigid walls that do not expand or contract.
For the two conditions, the amount of heat absorbed or given out is different for a reaction, so we wish to understand the difference in heat change.
Pressure-Volume Work (PV-Work)
Need a song to get you through the rest of the post? Here you go 🙂
PV work is mechanical (not electrical, for example) and is the work associated with the volume change of a system.
w = -PΔV
where P is the constant external pressure of the reacting system and
ΔV = V final – V initial
given that the external pressure is constant (isobaric reaction).
In general, the change in internal energy of a closed system is
ΔE = q + w = q + w PV + w non-PV
Since by definition, a constant volume process has a “w PV = 0”, assuming that all work is PV work, for an isochoric process:
q = ΔE
Therefore, we can measure the change in internal energy by measuring the heat absorbed or given out in a constant volume chemical reaction (assuming all work is PV-Work)
Because q + w = ΔE, the heat of the reaction for a constant pressure process is:
q = ΔE + PΔV
So we can describe a constant pressure process, it’s convenient to have another state function enthalpy, H:
H = E + PV
So enthalpy isn’t easily visualized – it’s a mathematical construct. ΔH is the enthalpy of the products minus the enthalpy of the reactants.
For any reaction:
ΔH = ΔE + Δ(PV)
For a constant pressure reaction:
ΔH = ΔE + PΔV
The heat of a reaction at constant pressure equals the change in enthalpy of the reacting system, if the work involves only pressure-volume work, so that
q = ΔE
ΔH is positive for an endothermic reaction and negative for an exothermic reaction.
Finally, a thermochemical equation is a chemical equation that includes the enthalpy change. For example:
2 H2(g) + O2(g) → 2 H2O(l) ΔH = -571.6 kJ
Thermodynamic Standard States
We don’t want to tabulate heats of reaction for all temperatures and pressures, so we have defined thermodynamic standard states.
The standard state of a substance is its pure form at 1 bar pressure and a temperature of 25 C.
The standard enthalpy change is the enthalpy change for reactants in their standard state going to products in their standard state. We use a superscript zero to denote standard state conditions. These are usually given in a table.
There are different types of standard enthalpies of reactions.
- Enthalpy of Combustion – the standard enthalpy change accompanying the combustion of 1 mole of a substance in oxygen.
- Enthalpy of Formation – the standard enthalpy change accompanying the formation of 1 mole of pure substance from its elements in their most stable forms, with all substances in their standard states.
The standard enthalpies of formation of all elements in their most stable form are zero.
The enthalpy change for any chemical reaction is the same, regardless of the path by which the reaction occurs. (This law is a consequence of the fact that enthalpy is a state function.)
That’s saying that there is more than one pathway for a reaction (via a different intermediate, for example).
The overall reaction enthalpy is the sum of the reaction enthalpies of the individual reactions into which a reaction may be divided.
This is sooo good! Because otherwise, we wouldn’t be able to calculate enthalpy changes that are difficult to impossible to measure!
Now, there are two methods for finding ΔH.
Here’s an example 0f (1) the algebraic method:
(2) Alternatively, we can write Hess’s law as:
so that the enthalpy change of a reaction is equal to the sum of the enthalpies of formation of the products minus the sum of the enthalpies of formation of the reactants.
Bond Energy (Bond Enthalpy)
The bond dissociation enthalpy is defined as the standard reaction enthalpy for the process in which the A-B bond is broken homolytically:
A-B(g) -> A(g) + B(g)
where A and B could be atoms or groups of atoms.
Bond enthalpies are derived from the heats of formation of the species involved in the definition of bond enthalpy or from spectroscopic measurements.
These bond energies are always positive!
When bond enthalpies are given in a table, this is their mean value averaged from many different molecules. This is useful because they allow us to make estimates of enthalpy changes.
Calorimetry is the experimental method of determining heats of reactions.
The heat capacity (C) of an object is the amount of heat required to raise the temperature of the substance by 1 K. The SI unit of heat capacity is J/K.
The specific heat (c) is the heat capacity per unit mass. The SI unit is J/kg/K, and you can find these values in tables.
q = mcΔT
If we know the specific heat capacity and the mass of a substance, then the change in the sample’s temperature will tell us the amount of heat (q) that has been absorbed or released in a process.
Great job! Hope you enjoyed the music, and hope you enjoyed learning about thermochemistry. As always, if you have questions, feedback, or anything you want me to know, leave it in the comments! I always check 🙂