Enthalpy Change for Exothermic and Endothermic Reactions Chemistry Tutorial

Key Concepts

• The enthalpy of a chemical system refers to the "heat content" of the system. ¹
• Enthalpy is given the symbol H
• Enthalpy change refers to the amount of heat released or absorbed when a chemical reaction occurs at constant pressure.
• Enthalpy change is given the symbol² ΔH
• Enthalpy change for a chemical reaction (ΔH) is defined as the enthalpy of the products (H_products) minus the enthalpy of the reactants (H_reactants)

  ΔH = H_(products) - H_(reactants)

• An exothermic reaction is a chemical reaction or physical change in which heat is released.

  ⚛ Heat is a product of the reaction.

  ⚛ Temperature of reaction mixture increases.

  ⚛ H_products < H_reactants

  ⚛ ΔH = H_products - H_reactants = a negative number

• An endothermic reaction is a chemical reaction or physical change in which heat is absorbed.

  ⚛ Heat is a reactant in the reaction.

  ⚛ Temperature of reaction mixture decreases.

  ⚛ H_products > H_reactants

  ⚛ ΔH = H_products - H_reactants = a positive number

• Enthalpy change, ΔH, values are given in units of energy unit per mole.

  ⚛ The S.I. units³ for enthalpy change are joules per mole, J mol^-1 ( or J/mol)

  ⚛ Most commonly, enthalpy change is given in units of kilojoules per mol, kJ mol^-1 (or kJ/mol) ⁴

• ΔH is specified per mole of substance as in the balanced chemical equation for the reaction

  (See Enthalpy Change Calculations for Chemical Reactions tutorial.)

• Enthalpy changes are measured under standard laboratory conditions.

  (See Heat of Reactions and Heat of Combustion tutorials.)

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Enthalpy Change for Exothermic Reactions

An exothermic reaction is defined as a chemical reaction or a physical change that releases heat.
When a chemical reaction gives off heat the temperature of the system increases.
When hydrochloric acid is added to water the temperature of the water increases because heat is being given off as the hydrochloric acid dissolves in the water. This is an example of an exothermic reaction.
Similarly, if you dissolve sodium hydroxide pellets in water the temperature of the water will increase as heat is given off as the sodium hydroxide dissolves in the water. This is also an example of an exothermic reaction.
A common, everday example of exothermic reactions are the combustion of fuels. When we combust (or burn) a fuel, heat is released. We can use this heat to warm ourselves, like when we burn coal or wood in a fireplace. We can use this heat to cook our food, like combusting (burning) charcoal in a barbeque. We can use this heat to do work for us, like petrol (gasoline) is used in a car.

Let's take the combustion (burning) of coal or charcoal as an example of an exothermic reaction.

When coal (charcoal), C_(s), combusts in excess oxygen, O_2(g), carbon dioxide, CO_2(g), is produced.
The balanced chemical equation for this reaction is given below:

We know that this reaction gives off (releases or produces) heat because we can feel the "heat" from the fire warm us.
"Heat" is therefore a product of the reaction, so we could include it on the "products" side of the chemical reaction:

But the First Law of Thermodynamics tells us that energy is conserved, that is, energy can neither be created nor destroyed.
So we where did this heat energy come from?

In order for energy to be conserved, the energy present on the reactant side of the chemical equation must be the same as the energy present on the product side of the equation:

The "energy" of each species is its enthalpy, H, so we can write:

The enthalpy of the reactants, H(C_(s)) + H(O_2(g)), is greater than the enthalpy of the product, H(CO_2(g)), by an amount equal to the heat released by the chemical reaction (the "heat" in the equation).

Which we could represent schematically⁵ as:

The "heat released" is therefore the change in enthalpy when reactants react to produce products:

So we can replace "heat" in our balanced chemical equation with enthalpy change (ΔH):

So far, so good.
However, the enthalpy change term, ΔH, is not usually incorporated into the chemical equation, it is usually shown as a separate term:

C_(s) + O_2(g) → CO_2(g)     ΔH_reaction = ? kJ mol^-1

and the enthalpy change for a reaction, ΔH_reaction, is defined as:

ΔH_reaction = H_products - H_reactants

and we saw above that the enthalpy of the product, H(CO_2(g)), is less than the enthalpy of the reactants, H(O_(s)) + H(O_2(g)), :

H(CO_2(g)) < [H(C_(c)) + H(O_2(g))]

so the enthalpy change, ΔH, for the reaction will have a negative value:

ΔH = H(CO_2(g)) - [H(C_(c)) + H(O_2(g))] = a negative number

We can look up a table of values to find that the combustion of 1 mole of carbon releases 393.5 kJ of heat energy.
So we can write ΔH = -393.5 kJ mol^-1
Note, the value of ΔH is negative (-) because the reaction is exothermic (releases heat).

We can now represent the combustion of this carbon in either of two ways:

Enthalpy Change for Endothermic Reactions

A chemical reaction or physical change is endothermic if it absorbs energy from its surroundings.
As heat is absorbed by the reaction the temperature of the reaction mixture will decrease.

Endothermic reactions can be very useful.
When you place an ice cube in your drink to keep it cool, this is an example of an endothermic reaction. Heat from the cool drink is absorbed by the ice cube, which keeps the drink cool while the absorbed heat is used to melt the ice cube. The melting of the ice cube is an example of an endothermic reaction, heat is being absorbed from the surroundings.
If you've ever had an injury while playing sport you may have used a disposable (or single-use) cold pack. In this case you break a barrier in the cold pack which allows ammonium nitrate, NH₄NO_3(s), to dissolve in water. It absorbs heat from the surroundings while it dissolves, so that the temperature of the reaction mixture (the cold pack) decreases and it feels cold. This is also an example of an endothermic reaction because heat is being absorbed from the surroundings to enable the ammonium nitrate to dissolve in the water.
The thermal decomposition of some compounds like calcium carbonate and calcium hydroxide requires heat to be absorbed to decompose (break apart) the reactant compound. These are also examples of endothermic reactions.

Calcium hydroxide, Ca(OH)_2(s), is an important industrial base. It can be heated so that it decomposes into water, H₂O_(g), and calcium oxide, CaO_(s), which is used in the manufacture of cement.
The balanced chemical equation for the thermal decomposition of calcium hydroxide is given below:

Ca(OH)_2(s) → CaO_(s) + H₂O_(g)

We know that the reaction only occurs if it is heated, so heat must be a reactant:

The First Law of Thermodynamics tells us that energy must be conserved during a chemical reaction, that is, energy can neither be created nor destroyed, therefore:

and since we know that the "energy of reactants" is referred to as their enthalpy, and the "energy of products" is referred to as their enthalpy, we can write:

The enthalpy of the products, H(CaO_(s)) + H(H₂O_(g)), must be greater than the enthalpy of the reactant, H(Ca(OH)_2(s)), by an amount equal to the "heat" term.

Which we could represent schematically as:

"Heat absorbed" by the reaction is the enthalpy change for the reaction, ΔH, that is

heat absorbed = H_products - H_reactants = ΔH

so we can replace "heat" in our balanced chemical reaction with ΔH :

Ca(OH)_2(s) + ΔH → CaO_(s) + H₂O_(g)

Now, because the enthalpy of products is greater than the enthalpy of reactants in our reaction:

H_products > H_reactants

the value of ΔH will be positive:

ΔH = H_products - H_reactants = a positive number

If we look up a table of values we would find that the decomposition of 1 mole of calcium hydroxide absorbs 66 kJ of heat energy, that is, ΔH = +66 kJ mol^-1 (or ΔH = +66 kJ/mol)

We could include this value as a reactant in the chemical equation as shown below:

Ca(OH)_2(s) + 66 kJ mol^-1 → CaO_(s) + H₂O_(g)

Or we can separate the enthalpy change term from the chemical reaction as shown below:

Ca(OH)_2(s) → CaO_(s) + H₂O_(g)     ΔH = +66 kJ mol^-1

Summary of Enthalpy Changes for Chemical Reactions

The table below summarises the differences between an exothermic and an endothermic reaction :

Worked Example of Identifying Exothermic and Endothermic Reactions using Enthalpy Change

Question: Thermite can be used for welding in places where neither electricity nor welding gases are available because the thermite reaction evolves enough heat, about 852 kJ per mole of ferric oxide, to generate temperatures of about 3000°C which is hot enough to melt iron.
In the thermite reaction, iron(III) oxide (ferric oxide, Fe₂O_3(s)), reacts with powdered aluminium (Al_(s)) to produce liquid iron (Fe_(l)) and aluminium oxide (Al₂O₃) according to the following balanced chemical equation:

2Al_(s) + Fe₂O_3(s) → 2Fe_(l) + Al₂O_3(s)

What is the value of ΔH for the thermite reaction?

Solution:

(Based on the StoPGoPS approach to problem solving.)

1. What is the question asking you to do?

   Determine the value of ΔH
   ΔH = ? kJ mol^-1

2. What data (information) have you been given in the question?

   Extract the data from the question:

   (i) balanced chemical equation:

   2Al_(s) + Fe₂O_3(s) → 2Fe_(l) + Al₂O_3(s)

   (ii) the amount of heat produced = 852 kJ per mole of Fe₂O_3(s)

3. What is the relationship between what you know and what you need to find out?

   Exothermic reaction: heat is produced, ΔH is negative

   Endothermic reaction: heat is absorbed, ΔH is positive

4. Decide if reaction is exothermic or endothermic and therefore determine the value of ΔH:

   (i) 852 kJ of energy is produced per mole of Fe₂O_3(s), therefore reaction is exothermic.

   (ii) For an exothermic reaction, ΔH has a negative value, ΔH = -852 kJ mol^-1

5. Is your answer plausible?

   Consider the First Law of Thermodynamics, that is, energy is conserved during a chemical reaction:

   energy of reactants = energy of products

   Since the temperature of the reaction mixture increases to about 3000°C, heat is a product of the reaction, so

   H_reactants = H_products + heat

   This shows us that the enthalpy of the products is less than the enthalpy of the reactants:

   H_products < H_reactants

   So when we calculate ΔH for the reaction:

   ΔH = H_products - H_reactants

   ΔH will be a negative number, ΔH = -? kJ mol^-1

   We have been told that 852 kJ of energy is produced per mole of Fe₂O_3(s), so ΔH = - 852 kJ mol^-1

   Since this is the same as the answer we got above, we are reasonably confident that our answer is plausible.

6. State your solution to the problem "ΔH for the thermite reaction":

   ΔH = -852 kJ mol^-1

Footnotes