EXOTHERMIC REACTIONS and ENDOTHERMIC REACTIONS

PART A Exothermic and Endothermic Energy Changes - Chemical Energetics Introduction -  energy transfers in physical state changes AND chemical reactions

 Sub-index for this page

1(a) Introduction to exothermic and endothermic changes and the law of conservation of energy

1(b) Why is it important to know about energy changes in chemical reactions?

1(c) Heat Changes in Chemical Reactions - reaction profile diagrams - examples of exothermic and endothermic chemical reactions

1(d) Heat changes in physical changes of state

1(e) More advanced 'delta' notation and terminology in thermochemistry

2. Reversible Reactions and energy changes

Index of GCSE level energetics notes

GCSE Chemistry Revision notes

Energy is conserved in chemical reactions.

One way of stating the 'law of Conservation of Energy' is to say the amount of energy in the universe at the end of a chemical reaction is the same as before the reaction took place.

All chemicals have there own unique chemical energy store - by virtue of their unique atomic/molecular/ionic chemical structure.

Different chemicals, with different structures, store different amounts of chemical potential energy.

In the graphs describing energy level changes you should think of the vertical y axis marked ENERGY as the chemical potential energies of the substance.

The greater the ENERGY value, the greater the potential to effect chemical change, but it might not always go in the direction you think e.g. by heating reactants you can make them go 'uphill' - an analogy with physics and gravitational potential energy or heating a material to a higher temperature to increase its thermal energy store - in both these cases you are increasing the potential energy of the material in some way.

If a reaction transfers energy to the surroundings the product molecules must have less energy than the reactants, by the amount transferred - exothermic energy change.

Conversely, if a reaction absorbs energy from the surroundings, they must have less energy, and the products must have more energy - endothermic energy change.

The law of conservation of energy in chemistry parallels the law of conservation of mass.

Both laws allow us to make theoretical calculations and predictions.

AND REMEMBER - no mass is lost and no energy is lost - it must all add up in the end.

An exothermic chemical reaction transfers energy to the surroundings, usually given out in the form of heat energy, so raising the temperature of the surroundings.

Therefore the products store less energy than the reactants and the surroundings have more energy.

The excess energy in an exothermic reaction is transferred to the surroundings increasing its thermal energy store.

Exothermic reactions include combustion of fuels, many oxidation reactions, acid-alkali neutralisation reactions, reactive metals with water, moderately reactive metals with strong acids.

Exothermic reactions are used in heat energy sources - burning fuels, self-heating cans and hand warmers.

An endothermic chemical reaction absorbs energy from the surrounding, usually in the form of heat energy, so cooling the surroundings, but sometimes the system is heated to provide the heat energy and a high enough temperature to promote the reaction.

This means the products store more energy than the reactants and the surroundings have less energy.

In an endothermic reaction, the extra energy taken in from the surroundings (or any heat source) decreases a thermal energy store.

Endothermic reactions include thermal decomposition of compounds e.g. carbonates, the reaction between citric acid and sodium hydrogencarbonate, sports injury packs to produce cooling effects.

1(b) Why is it important to know about energy changes in chemical reactions?

Its important to know how much energy fuels release on combustion i.e. their calorific value.

Its important to know the energy released on burning petrol. diesel, coal or any other fossil fuel and alternative fuels like hydrogen or biofuels (biomass fuels).

The same sort of data is important in knowing how much energy is released on metabolising foods such as fats and carbohydrates.

Accurate energy change data is important in managing chemical processes in industry.

Exothermic reactions may provide their own heat if the process is carried out at high temperatures, energy transfer data provides some of the information needed.

Conversely, excess heat from an exothermic reaction may have to be removed using heat exchangers to avoid 'overheating' and excessive reaction rates that could be dangerous. If gases are involved, lack of control could lead to a build up of pressure resulting in an explosion.

Endothermic chemical processes often need a high temperature to promote the absorption of heat energy, otherwise the reaction rate might economically far too slow. The amount of energy needed can be calculated from energy transfer data.

See sections 6. Calorimeter methods of determining energy changes and EXAMPLES of experiments you can do

and 7. Energy transfer calculations from calorimeter results

• When chemical reactions occur, as well as the formation of the products - the chemical change, there is also a heat energy change which can often be detected as a temperature change.
• This means the products have a different energy content than the original reactants (see the reaction profile diagrams below).

Examples of EXOTHERMIC REACTIONS

• If the products contain less energy than the reactants, heat is released or given out to the surroundings and the change is called an exothermic reaction (exothermic energy transfer, exothermic energy change of the system).
  • ENERGY = chemical potential energy of reactants and products
  • This is illustrated by the simple energy level diagram above for an exothermic reaction.
    • The vertical axis is energy content of the reactants and products.
    • The horizontal axis represents the course of the reaction or the progress of reaction.
    • The products have less energy (lower level) than the original reactants (higher energy level) and the difference between the levels is the heat energy released to the surroundings.
    • So in this case the difference in heights of the energy levels tells you how much energy is released in an exothermic reaction.
    • See also reaction profiles in section B where the activation energy has been added to the reaction profiles.
  • The temperature of the system will be observed to rise in an exothermic change.
  • So an exothermic reaction is one which gives out energy to the surroundings, usually in the form of heat energy, hence the rise in temperature.
  • Examples of exothermic reactions:
    • The burning or combustion of hydrocarbon fuels (see Oil Products) e.g. petrol or candle wax, these are very exothermic reactions.
      • The exothermic burning-combustion of fossil fuels is very important source of energy.
      • methane (natural gas) + oxygen ==> carbon dioxide + water (+ heat energy)
        • CH₄ + 2O₂ ==> CO₂ + 2H₂O
    • The burning of magnesium, reaction of magnesium with acids, or the reaction of sodium with water (see Metal Reactivity Series)
      • 2Mg + O₂ ==> 2MgO (+ heat energy)
    • Using hydrogen as a fuel in hydrogen-oxygen fuel cells (see Electrochemistry).
    • All these combustion reactions are oxidations.
    • Explosions are caused by VERY fast exothermic reactions producing very fast large expanding volumes of gases.
    • Metal displacement reactions are also exothermic. If you add iron filings to copper sulfate solution there is quite a temperature rise.
      • iron  +  copper sulfate  ===>  iron sulfate  +  copper
      • Fe(s)  +  CuSO₄(aq)  ==>  FeSO₄(aq)  +  Cu(s)
    • The neutralisation of acids with alkalis (see Acids, Bases and salts) e.g.
      • sodium hydroxide  +  hydrochloric acid  ==>  sodium chloride  +  water
      • NaOH + HCl ==> NaCl + H₂O (+ heat energy)
      • Its the same for the neutralisation reactions between potassium hydroxide and sulfuric and nitric acids etc.
  • Other uses of exothermic reactions:
    • Hand warmers contain chemicals that when mixed together give out heat.
    • Self-heating cans of coffee, soup or hot chocolate have chemicals contained in the base of the container that when mixed generate enough energy to heat the contents of the can.
    • The Thermit reaction between aluminium powder and iron(III) oxide is VERY exothermic and when the mixture is ignited with a lit magnesium 'fuse' it goes off like a firework.
      • aluminium  +  iron(III) oxide  ===>  aluminium oxide  +  iron
      • 2Al  +  Fe₂O₃  ===>  Al₂O₃  +  2Fe
      • This is another example of a displacement reaction where a more reactive metal displaces a less reactive metal from one of its compounds.
      • This kind of reaction is used to extract certain metals from their purified ores.

See sections 6. Calorimeter methods of determining energy changes and EXAMPLES of experiments you can do

and 7. Energy transfer calculations from calorimeter results

Examples of ENDOTHERMIC REACTIONS

• If the products contain more energy than the reactants, heat is taken in or absorbed from the surroundings and the change is called an endothermic reaction (endothermic energy transfer, endothermic energy change of the system).
  • ENERGY = chemical potential energy of reactants and products
    • This is, again, illustrated by the simple energy level diagram above for an endothermic reaction.
    • The vertical axis is energy content of the reactants and products.
    • The horizontal axis represents the course of the reaction or the progress of reaction.
    • The products have more energy (higher level) than the original reactants (lower energy level) and the difference between the levels is the heat energy absorbed from the surroundings.
    • So in this case the difference in heights of the energy levels tells you how much energy is absorbed in an endothermic reaction.
    • See also reaction profiles in section B where the activation energy has been added to the reaction profiles.
  • If the change can take place spontaneously, the temperature of the reacting system will fall but, as is more likely, the reactants must be heated to speed up the reaction and provide the absorbed heat.
  • So an endothermic reaction is one which absorbs energy from the surroundings, usually in the form of heat, hence the observed fall in temperature in some reaction OR you heat the reaction mixture to supply the heat energy required to effect the chemical change.
  • Examples of endothermic reactions
    • The reaction between citric acid and sodium hydrogencarbonate is endothermic.
    • the thermal decomposition of limestone (see Industrial Chemistry)
      • calcium carbonate (limestone) ==> calcium oxide (lime) + carbon dioxide
        • CaCO₃ (+ heat energy) ==> CaO + CO₂
        • This only happens at temperatures above 900^oC.
    • the cracking of oil fractions (see Oil products)
      • e.g. octane (+ heat energy) ==> hexane + ethene
        • C₈H₁₈ ==> C₆H₁₄ + C₂H₄
        • Again this needs a very high temperature AND a catalyst too.
    • These are two very important endothermic reactions used in the chemical industry.
    • A few simple experiments to illustrate endothermic reactions
      • Dissolving ammonium nitrate in water doesn't need heating, the salt spontaneously dissolves and the temperature of the water/solution immediately falls as energy from the surroundings is absorbed, in fact from the water itself.
      • Adding ammonium chloride to barium hydroxide solution produces ammonia gas and quite a cooling effect below 0^oC!
        • ammonium chloride  +  barium hydroxide  ==> barium chloride  +  water  +  ammonia
        • 2NH₄Cl  +  Ba(OH)₂  ==>  BaCl₂  +  2H₂O  +  2NH₃
      • Adding sodium hydrogencarbonate to citric acid solution also produces a cooling effect.
  • Other uses of endothermic reactions:
    • Some sports injury packs have a mixture of chemicals (or maybe a salt and water) that when mixed undergo an endothermic energy change, thereby absorbing heat from the surroundings, cooling some poor bruised limb! Rather more convenient and less messy than packs of ice!
      • Conversely hand warmer packs use an exothermic reaction between chemicals that mix on activating the pack.
    • One of the most important endothermic reactions, for which most of animal life depends is photosynthesis.
    • The energy from sunlight is absorbed as water and carbon dioxide are converted to glucose and oxygen.
      • 6H₂O + 6CO₂ + sunlight energy ==> C₆H₁₂O₆ + 6O₂
      • See GCSE biology notes on photosynthesis
    • However, on 'burning' the glucose/carbohydrates in our bodies, the 'stored' sunlight energy is released to keep us warm and drive all the chemical processes in our cells, so the opposite reaction is exothermic!
      • C₆H₁₂O₆ + 6O₂  ==>  6H₂O_(l) + 6CO₂  + heat/chemical energy
      • See GCSE biology notes on respiration
• There are brief descriptions of other examples of exothermic and endothermic reactions on the "Types of Reaction" page.
• The difference between the energy levels of the reactants and products gives the overall energy change for the reaction (the activation energies are NOT shown on the diagrams below, but see section 3.).

See sections 6. Calorimeter methods of determining energy changes and EXAMPLES of experiments you can do

and 7. Energy transfer calculations from calorimeter results

• If the direction of a reversible reaction is changed, the energy change is also reversed.

• For a reversible reaction, the energy released in the exothermic reaction is numerically equal to the heat absorbed in the reverse reaction.

• For example: the thermal decomposition of hydrated copper(II) sulphate is a very good example to observe in the school laboratory, even though it is not practical to measure the actual energy changes involved.

• blue hydrated copper(II) sulphate + heat    white anhydrous copper(II) sulphate + water

• CuSO₄.5H₂O_(s)    CuSO_4(s) + 5H₂O_(g)

• On heating the blue solid, hydrated copper(II) sulphate, steam is given off and the white solid of anhydrous copper(II) sulphate is formed and left as the residue.
  • This is a thermal decomposition and is endothermic as heat is absorbed (taken in) from the surroundings to drive off the water.
  • The energy is needed to break down the crystal structure and drive off the water.
• When the white solid is cooled and a few drops of water added, blue hydrated copper(II) sulphate is reformed and heat energy is given out to the surroundings, the mixture hots up!
  • The reverse reaction is exothermic as heat is given out.
  • i.e. on adding water to white anhydrous copper(II) sulphate the mixture heats up as the blue crystals reform.
  • Water molecules recombine with the copper ion releasing energy when the new bonds are formed.
• This illustrate the general rule for ANY reversible reaction - change the direction of chemical change you also change whether heat is released to the surroundings (exothermic) or heat is absorbed from the surroundings (endothermic).
  • So, reversing a chemical change also reverses the energy change - but this remains numerically the same.

• For more on reversible reactions and chemical equilibrium see GCSE notes OR Advanced A Level Notes.

• LINKS to associated webpages

  • Foundation tier-easier multiple choice GCSE QUIZ on exothermic/endothermic reactions etc.

  • Higher tier-harder multiple choice GCSE QUIZ on exothermic/endothermic reactions etc.

  • GCSE/IGCSE/O level notes on Oil-Fuel burning

  • GCSE/IGCSE/O level Types of Chemical Reaction Notes

  • GCSE/IGCSE/O Level Rates of Reaction Revision Notes