enthalpy change

The system and the surroundings

When dealing with chemical reactions we often talk about the system and the surroundings. The system is the reacting chemicals; the reactants and the products while the surroundings are everything else... the beaker, the room, the building, the universe!
The system can be open or closed. An open system is one where both matter (reacting chemicals) and energy can be exchanged freely with the surroundings. Carrying out a reaction in an open beaker where for example heat and gases can escape into the atmosphere is an example of a reaction taking place in an open system. A closed system is one which can exchange energy (usually heat) with the surrounding but not matter. The diagram below shows both open and closed systems. In the closed system below two gases are reacting in a sealed container which has a movable piston on top. The piston will allow energy to enter and leave the system as it moves up and down but no gases (matter) can enter or leave.

Examples and definitions for open and closed systems.

Heat, energy and enthalpy

What is enthalpy?

Energy is a word which has probably been used many times in your science lessons, but what exactly is energy? The word energy is derived from the Greek word ergon; meaning work. A simple working picture of energy might be the capacity to do work or supply heat

Energy = heat + work

One of the main forms of energy we are concerned with in chemistry is potential energy. There is stored potential energy or chemical energy within the bonds of a molecule. When a substance undergoes a chemical reaction some of this stored chemical energy maybe released as heat and light energy. When a strip of magnesium metal is held in a hot Bunsen flame or even the methane gas used in the Bunsen burner reacts with the oxygen in the air the reacting chemicals lose some of the stored potential or chemical energy in their bonds as heat and light as the products of the reaction form.

Magnesium burning in air is an example of an exothermic reaction.

Internal energy

Consider for a second a chemical system (reactants and products), it could be a flask filled with say hydrogen and oxygen or a beaker of acid reacting with magnesium ribbon, now consider the fact that all the atoms in this system are moving and so have kinetic energy. However the individual atoms consist of electrons which are also moving and will also have kinetic energy. There will also be potential energy present as a result of the separation of the protons and electrons and we also need to think about potential energy due to all the interactions of the molecules, atoms and sub-atomic particles that make up the system. The internal energy of a system (symbol E) is the sum of all the kinetic and potential energy of all the atoms and sub-atomic particles that make up the system. Not surprisingly it is not possible to calculate the internal energy (E) of a system, even with the most powerful and advanced computers available today.


From your gcse science course you may remember that in physics we used a definition of work which was: work is done when a force moves an object a distance, d.

work done = force x distance
w = F x d
Now if you consider any chemical reaction which releases a gas then this reaction will carry out work as the gas released has to push against the pressure of the atmosphere to escape and expand. This type of work is often called pressure-volume or simply PV work. As a simple example consider a reaction involving gases which takes place in a metal cylinder fitted with a movable piston, this is shown in the image below. As the reaction happens heat energy is given off and the piston is forced upwards by the gases produced. These gases are applying a force to the piston which moves; so we can say the gases are doing work on the piston. We can use the equation below to calculate the size of the force acting on the piston:
force = pressure x area of piston
F = P x A
The pressure pushing down on the stationary piston will simply be atmospheric pressure (P). Now the pressure of the gas in the cylinder will be pushing in the opposite direction to atmospheric pressure, so we will call this pressure -P, since it is being applied against atmospheric pressure. It will be the same pressure as atmospheric as the piston is being held in a stationary position. So we have:
F = -P x A
If the gases push the piston upwards a distance d, then by substituting the -P x A for force we have:
w = F x d
w = -P x A x d
Now it is apparent that the value A x d (area of piston x distance moved) is simply the change in volume (ΔV) of the gases in the cylinder, so we have:
w = -P x A x d
w = -P ΔV
So the work done by the gas is simply the pressure x change in volume. You may notice from this equation that the term ΔV will be positive, since the volume has increased, this means that the value of the work term will be negative. This simply means that the system (reacting chemicals) is losing energy.

So far we have seen that a system can exchange energy with its surrounding as heat and work:

energy = heat + work
If the reaction is carried out in a closed system where no volume change is possible then the term -P ΔV will be equal to zero and all the energy will be exchanged as heat. However most of the reactions we carry out in the lab are in containers such as beakers and test-tubes which are open to the air and carried out at constant pressure (atmospheric pressure). This means that the system is likely to exchange energy with the surroundings in the form of both heat and work. So we have:
energy transferred = heat + work
ΔE = qp + w
where qp is the heat transferred at constant pressure. However we have already seen that: W= -PΔV
So this by simply substituting for work (W) in the equation above we have:
ΔE = qp -PΔV
or by rearranging this we have:
qp = ΔE + PΔV
The heat released at constant pressure (qp) is more commonly called the enthalpy change for a reaction and it is given the symbol ΔH. The enthalpy of a system is simply:
H = E + PV

However like internal energy (E) it is not possible to calculate the enthalpy of a system. To be honest as chemists even if it were possible to calculate the enthalpy of a system, the calculated value would be of little practical use. What is much more useful is the change in enthalpy that takes place during a chemical reaction. The enthalpy change (ΔH) is simply the heat energy lost or gained by the system at constant pressure plus any work done or gained by the system. For most practical purposes the value of ΔH and ΔE vary little, especially when there is a small change in the volume in going from the reactants to the products in a chemical reaction, since little or no work will be done. The enthalpy change for a reaction is simply:

ΔH = Hproducts - Hreactants

The law of conservation of energy

You should also consider the first law of thermodynamics which is derived from the law of conservation of energy. When considering energy change. This law basically states that all the energy from when the universe was created in the big bang is still all here. You cannot make or destroy energy only change its form.

The law of conservation of energy.

For us chemists this basically means that any energy lost by a reacting system will be gained by the surroundings and in the case of an endothermic reaction any heat energy lost by the surrounding will be gained by the system.

Exothermic and endothermic reactions

In your gcse science course you will have met endothermic and exothermic reaction. The image below shows some common chemical reactions, including: burning, displacement reactions, a neutralisation reaction and metal acid/water reactions, can you guess what all these reactions have in common?

Examples of exothermic reactions.

Well all these reactions release heat energy to the surroudings. They are all exothermic reactions.

An exothermic reaction release heat energy to the surroundings. A thermometer (part of the surroundings would measure an increase in temperature during an exothermic reaction.

As we have seen above the reacting chemicals (the system) in a reaction act as a store of chemical energy. During an exothermic reaction the system (chemicals which are reacting) loses energy to the surrounding, mainly as heat. The temperature of the surrounding will increase. Remember the law of conservation of energy, this states that energy cannot be created or destroyed, only changed from one form to another. So if the reacting chemicals lose energy then the surrounding must gain the energy lost by the reacting chemicals (the system).

Endothermic Reactions

Have you ever sucked on a sherbet sweet and felt your mouth getting slightly cooler? Well this is because there is an endothermic reaction taking place in your mouth. Sour sherbet sweets contain citric acid and a base called sodium bicarbonate. In your mouth these chemical react to release carbon dioxide, which gives the fizzing sensation when you eat sherbet, but the reaction between citric acid and sodium bicarbonate is an endothermic one. It removes heat energy from the surrounding, and in this case the surroundings are your mouth! So you should feel a cooling sensation in your mouth as you enjoy your fizzy sherbet sweets.

Enthalpy profile diagrams

The image below shows an energy profile diagram for exothermic and endothermic reactions. From the diagrams below it is clear that aware that the enthalpy change (ΔH) for an exothermic reaction is always negative while the enthalpy (ΔH) for an endothermic reaction is always positive. enthalpy profile diagram for an exothermic and an endothermic reaction.

Examples of endothermic reactions

Endothermic reactions are much less common than exothermic reactions. Example of an endothermic reaction, the reaction of ammonium chloride and barium hydroxide.  Experimental set-up and observations. One of the most spectacular endothermic reactions to see is one where solid ammonium chloride is added to a beaker containing solid barium hydroxide. The beaker containing the solid barium hydroxide is sitting on top of a block of wood on which a small pool of water has been added. When the two solids are added together and stirred vigorously so much heat energy is removed from the surroundings that the pool of water under the beaker freezes and forms a solid block of ice. This ice is so thick that it will stick the beaker to the wooden block.

The temperature changes in this experiemnt are very dramatic and large. If the experiment is done in a lab at room temperature (250C) then after the experiment is over the temperature of the reacting chemicals may have dropped to -100C a temperature drop of 350C. In this experiment the chemicals (the system) have gained energy from the surrounding. As a consequence the temperature of the surroundings has dropped. The text box below is a good summary of the energy changes taking place in an endothermic reaction:

An endothermic reaction removes heat energy from the surroundings. A thermometer (part of the surroundings) would measure a decrease in temperature during an endothermic reaction.

Key points