Chemistry AQA GCSE -Energy Transfer

Exothermic and endothermic reactions

Chemical reactions involve energy transfers. Many chemical reactions involve the release of energy. For other chemical reactions to occur, energy must be supplied.

Energy transfer in chemical reactions

a) When chemical reactions occur, energy is transferred to or from the surroundings.

b) An exothermic reaction is one that transfers energy to the surroundings. Examples of exothermic reactions include combustion, many oxidation reactions and neutralisation.

Everyday uses of exothermic reactions include self-heating cans (eg for coffee) and hand warmers.

Do you know some examples of chemical reactions and processes that use catalysts? Do you know how to draw conclusions from given data?

c) An endothermic reaction is one that takes in energy from the surroundings. Endothermic reactions include thermal decompositions. Some sports injury packs are based upon endothermic reactions.

d) If a reversible reaction is exothermic in one direction, it is endothermic in the opposite direction. The same amount of energy is transferred in each case.

For example:

hydrated

endothermic

anhydrous

copper copper + water

sulfate sulfate

(blue)

exothermic

(white)

Calculating and explaining energy change

Knowing the amount of energy involved in chemical reactions is useful so that resources are used efficiently and economically. It is possible to measure the amount of energy experimentally or to calculate it.

Energy from reactions

a) The relative amounts of energy released when substances burn can be measured by simple calorimetry, eg by heating water in a glass or metal container. This method can be used to compare the amount of energy released by fuels and foods.

b) Energy is normally measured in joules (J).

c) The amount of energy released or absorbed by a chemical reaction in solution can be calculated from the measured temperature change of the solution when the reagents are mixed in an insulated container. This method can be used for reactions of solids with water or for neutralisation reactions.

 Learn the steps of the practical work on:

■ investigating temperature changes of neutralisations and displacement reactions, eg zinc and copper sulfate

■ investigating temperature changes when dissolving ammonium nitrate, or reacting citric acid and sodium hydrogencarbonate

■ adding ammonium nitrate to barium hydroxide

■ demonstration of the addition of concentrated sulfuric acid to sugar

■ demonstration of the reaction between iodine and aluminium after activation by a drop of water

■ demonstration of the screaming jelly baby

■ demonstration of the thermite reaction, ie aluminium mixed with iron(III) oxide

■ investigation of hand warmers, self-warming cans, sports injury packs.

Learn how to:

■ evaluate everyday uses of exothermic and endothermic reactions.

• take measurements using temperature sensors to investigate energy transfer;

  •    suggest methods to make a named soluble salt and :name the substances needed to make a named insoluble salt. 

Make sure that you learn to:

■ consider the social, economic and environmental consequences of using fuels

■ interpret simple energy level diagrams in terms of bond breaking and bond formation (including the idea of activation energy and the effect on this of catalysts)

■ evaluate the use of hydrogen to power cars compared to other fuels

:

If you sit the Higher Tier paper you should be able to:

·       calculate the energy transferred in reactions using supplied bond energies;

·       represent the effect of a catalyst on an energy level diagram.

·       compare the advantages and disadvantages of the combustion of hydrogen with the use of hydrogen fuel cells from information that is provided.

·       understand simple energy level diagrams showing the relative energies of reactants and products, the activation energy and the overall energy change, with a curved arrow to show the energy as the reaction proceeds.

·       relate these to exothermic and endothermic reactions.

TIP: You are not required to show knowledge of the details of the reactions in fuel cells.

d) Simple energy level diagrams can be used to show the relative energies of reactants and products, the activation energy and the overall energy change of a reaction.

e) During a chemical reaction:

■ energy must be supplied to break bonds

■ energy is released when bonds are formed.

f) In an exothermic reaction, the energy released from forming new bonds is greater than the energy needed to break existing bonds.

g) In an endothermic reaction, the energy needed to break existing bonds is greater than the energy released from forming new bonds.

h) Catalysts provide a different pathway for a chemical reaction that has a lower activation energy.

i) Hydrogen can be burned as a fuel in combustion engines.

hydrogen + oxygen gives water

It can also be used in fuel cells that produce electricity to power vehicles.

You need to know the steps of how you can:

■ design an investigation to compare the energy produced by different liquid fuels and different foods using a simple calorimeter

■ measure and calculate the energy change for exothermic reactions (eg react acid with Mg ribbon) and endothermic reactions (eg dissolving potassium nitrate)

■ carry out some reactions and measure the energy produced, assuming that it is only the water in the solution that is being heated and that 4.2 joules will raise the temperature of 1cm3 of water by 1°C.

What you need to know:

•  the chemical tests specified in the subject content and interpret results of any of those tests applied to solutions or mixtures of substances in different contexts.

• to comment on results and data from such analyses that are presented to them.

TIP: What you do not need to know:

You are not required to have knowledge of 

• delta H (_H) conventions and enthalpy changes, including the use of positive values for endothermic reactions and negative values for exothermic reactions.

• to interprete detailed information that uses knowledge beyond that expected at GCSE.

•  Flame colours of other metal ions.