Energy Changes in Chemical Reactions

When chemical reactions take place they are often accompanied by energy changes.

Chemical reactions most frequently occur in open vesels. That is, they take place at constant pressure. Enthalpy refers to energy at constant pressure (volume may vary).

The term 'enthalpy' comes from the Greek word enthalpo, meaning warming up.

Enthalpy, H

You are given a sample of methane, CH4. How much energy does its molecules contain? The first thing you want to know is the amount of methane present. We are chemists, so the answer is 1 mole (16 g). Energy is measured in joules, J, so you begin by thinking where to start measuring from. There seems to be no starting point; can methane molecules ever have no energy contained within them? Indeed, it is impossible to know the total amount of energy stored in these molecules. What ever its value, the total amount of energy in a given amount of a substance (sometimes called the Heat energy content) is known as the enthalpy, denoted H.

Methane is a fuel, so how do we get energy from it? The answer is to react it with oxygen.

CH4(g) + 2O2(g)    ®    CO2(g) + 2H2O(l)

The above chemical equation shows that 2 moles (64 g) of oxygen molecules are required to burn 1 mole of methane. Again, it is impossible to know the total enthalpy (heat energy content) of the oxygen. Likewise, we can't know the total heat energy content of 1 mole of CO2 and 2 moles of H2O (the products).

Now imagine that we could find 'molar enthalpy values' for elements and compounds in a chemical data book. This would allow us to work out the amount of energy given out when methane reacts with oxygen to form carbon dioxide and water, that is, an overall change in enthalpy, DH, when the above reaction takes place. The following equations represent such a calculation.

DH = (HCO2 + 2HH2O) - (HCH4 + 2HO2)

In general,

DH = SHproducts - SHreactants

But remember, this is theoretical; it is not possible to determine the absolute value of the enthalpy of a chemical element or compound. However, DH values for chemical reactions can be obtained. They can be measured experimentally, or calculated using Hess's Law (see later), or worked out in other ways.

Exothermic and Endothermic reactions

When chemical reactions take place they are often accompanied by heat changes. The system (the reactants which form products) may give out heat to the surroundings, causing them to warm up. In this case the reactants have more stored energy (greater total enthalpy) than the products. Such chemical reactions are said to be exothermic. The system may take heat from the surroundings, causing them to cool down. In this case the reactants have less stored energy (less total enthalpy) than the products. Such chemical reactions are said to be endothermic.

Exothermic reactions give out energy to the surroundings.
Endothermic reactions take energy from the surroundings.

Most reactions take place at constant pressure...

It is possible to measure changes in heat energy that accompany chemical reactions. Most reactions take place in vessels that are open to the atmosphere, that is, they take place at constant pressure (volume may vary). The special name given to a change in heat energy content measured at constant pressure is enthalpy change. This change in enthalpy is denoted by DH. The value of DH (often expressed in kJ, or kJ mol-1 when appropriate) is given a negative sign for exothermic reactions and a positive sign for endothermic reactions, indicating whether the system loses or gains energy as a result of the reaction.

The value of DH is given a negative sign for an exothermic reaction.
The value of DH is given a positive sign for an endothermic reaction.

Precise thermochemical measurements are made in a closed vessel of fixed volume, such as a calorimeter. For a reaction involving a change in volume of gases there is a small but real difference in the measured heat change. You can read more about this here.

Enthalpy Level Diagrams...

Enthalpy level diagrams can be used to illustrate overall exothermic and endothermic changes. They show the difference in total enthalpy of the reactants and products for a reaction:

For an exothermic reaction the total enthalpy of the products is less than that of reactants. For an endothermic reaction the total enthalpy of the reactants is less than that of the products. For each, the difference in these total enthalpies is equal to the overall enthalpy of the reaction, DH.

Temperature and pressure matter...

As well as the amounts of substances reacting (molar amounts are taken), the precise value of DH depends on both the temperature and pressure at which it is measured. For this reason DH values are expressed at standard conditions (normally 298 K and 1 atm.). A standard enthalpy change is written as DH298. You can read more about temperature and pressure here.

Thermochemical data for chemical reactions can be found in chemical data books.

CH4(g) + 2O2(g) ® CO2(g) + 2 H2O(l)            DH298 = - 890.3 kJ mol-1
H2(g) + I2(s) ® HI(g)            DH298 = + 26.5 kJ mol-1

Note the following changes:

CO2(g) + 2 H2O(l) ® CH4(g) + 2O2(g)            DH298 = + 890.3 kJ
H2(g) + I2(s) ® 2HI(g)            DH298 = + 53.0 kJ

Units are important...

Also note how the units are expressed in these thermochemical equations. kJ mol-1 is used when 1 mole of a substance is burned and when 1 mole of compound is formed. Refer to the Enthalpy Definitions section below for more explanation.

Activation Enthalpy, Ea

Methane and oxygen do not react spontaneously when mixed. An input of energy, such as a flame, is required to get the reaction started, after which its exothermic nature will sustain it. With regard to the collision theory of reaction rates, molecules react only if in a collision they possess between them energy equal to or greater than a certain critical value. This is called the activation enthalpy, Ea. An enthalpy profile diagram illustrates this:

See also Activation Enthalpy.

Enthalpy Definitions

There are several enthalpy definitions. The combustion of elements and compounds, and the formation of compounds from their elements, are important here.

Standard Molar Enthalpy of Combustion, DHc,298. The heat evolved when 1 mole of an element or compound completely burns in oxygen, measured at 298 K and 1 atm. pressure.

Standard Molar Enthalpy of Formation, DHf,298. The heat change when 1 mole of a compound forms from its elements in their standard states, measured at 298 K and 1 atm. pressure. From the definition, the enthalpy of formation of an element at 298 K and 1 atm. is zero.

The definitions refer to measurements under standard conditions, but these are sometimes omitted.

Here is a very useful equation that follows from the definition for enthalpy of formation. It can be used to calculate overall enthalpy changes when the enthalpies of formation of the reactants and products are given. It is

DH = SHf (products) - SHf (reactants)

Thermochemical equations can be written to satisfy these enthalpy definitions. Always include state symbols, for example,

C(s) + O2(g) ® CO2(g)            DHc,298 = -393.5 kJ mol-1 (or DHf,298 = -393.5 kJ mol-1)
2C(s) + 3H2(g) + O2(g) ® C2H5OH(l)            DHf,298 = -277.1 kJ mol-1

Thermochemical Calculations and Hess's Law

A calorimeter is used to measure heat changes directly (with adjustments made for standard conditions at constant pressure). However, this is not possible for some reactions. Their enthalpy changes can be calculated with the application of Hess's Law.

Hess's Law says that the overall enthalpy change accompanying a chemical reaction is independent of the route taken in going from reactants to products (provided that in each case the same initial and final states of temperature and pressure apply to the reactants and products).
Hess's Law (G.H. Hess, 1840) of 'constant heat summation' follows from the First Law of Thermodynamics, which can be stated very simply: Energy can neither be created nor destroyed.

Thermodynamics in the study of transforming energy from one form to another.

Specific heat capacity is used in some thermochemical calculations. It is the amount of energy required to raise the mass of 1 gram of a substance by 1 C.

q = specific heat capacity x mass x DT.

The specific heat capacity of water is 4.184 J g-1 K-1 . This means that 4.184 J of energy are required to raise each gram of water by each C.

Heat capacity can also be used in calculations. It is the amount of energy required to raise the whole mass of a body by 1 C.

q = heat capacity x DT.

Why do energy changes accompany chemical changes?

The answer involves the chemical bonds that hold atoms together in molecules. In chemical reactions the rearrangement of atoms involves breaking chemical bonds in reactant molecules and forming new bonds in product molecules. The atoms are themselves not created or destroyed, but are simply rearranged. Breaking chemical bonds requires energy; bond breaking is an endothermic process. Conversely, when chemical bonds form, energy is given out; bond formation is an exothermic process.

The energy needed to break a bond and the energy given out when a bond forms are definite and characteristic for each bond. The energy of the product molecules may therefore be greater or smaller than the energy of the reactant molecules. Thus a chemical reaction is accompanied by an overall enthalpy change. This is the enthalpy of reaction, DH.

DH for a reaction can be calculated from bond enthalpies. However, this value will be different from that calculated in other ways (for example, as above by applying Hess's Law) or given in chemical data books. Two reasons for this are:

For more detail about molar bond enthalpy and the calculation of DH for a chemical reaction using bond enthalpies see:

Bond Breaking and Bond Making in Chemical Reactions

Further links include...

Construct a Hess Cycle
4 ways to do an Enthalpy Calculation.
Bond Enthalpy
Construct a Bond Enthalpy Cycle