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EXOTHERMIC and ENDOTHERMIC REACTIONS

Doc Brown's Chemistry KS4 GCSE, IGCSE, O level & A level Revision Notes

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

EXOTHERMIC REACTIONS and ENDOTHERMIC REACTIONS

Why is it important to know how much energy is transferred in an exothermic or endothermic reaction? Why are there energy changes when a chemical reaction takes place? Why do reactions give out heat energy to the surroundings (exothermic reaction) and other reactions absorb heat energy (endothermic reactions). Do physical state changes involve energy changes? Examples of exothermic energy changes and endothermic energy changes in chemical reactions are described and explained. Also note that if a chemical reaction is reversible, one chemical change is exothermic and the other reverse reaction is endothermic and the energy changes are numerically equal. Uses of exothermic reactions and uses of endothermic reactions are described. These revision notes on chemical energy changes, energy transfers in exothermic reactions and endothermic reactions  should prove useful for the new AQA, Edexcel and OCR GCSE (9–1) chemistry science courses.


Sub-index for ENERGY CHANGES: 1. Heat changes in chemical/physical changes - exothermic and endothermic  *  2. Reversible reactions and energy changes (this page)  *  3. Activation energy and reaction profiles  *  4. Catalysts and activation energy  *  5. Introduction to bond energy/enthalpy calculations  *  6. Calorimeter methods of determining energy changes and examples of experiments  *  7. Energy transfer calculations from calorimeter results

See also Advanced Level Energetics–Thermochemistry – Enthalpies of Reaction, Formation & Combustion

1. Heat changes - EXOTHERMIC and ENDOTHERMIC

INTRODUCTION

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. If a reaction transfers energy to the surroundings the product molecules must have less energy than the reactants, by the amount transferred. Conversely, if a reaction absorbs energy from the surroundings, they must have less energy, and the products must have more energy. 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.

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 have less energy than the reactants and the surroundings have more energy.

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 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 have more energy than the reactants and the surroundings have less energy.

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.

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.


 

1a. Heat Changes in Chemical Reactions

  • 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).
  • 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).
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    • This is illustrated by the simple energy level diagram above for an exothermic reaction.
      • The products have less energy than the original reactants (lower energy level) and the difference comes out as heat energy released to the surroundings.
      • 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
    • 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.
  • 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).
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    • This is illustrated by the simple energy level diagram above for an endothermic reaction.
      • The products have more energy than the original reactants (higher energy level) and the difference comes out as heat energy absorbed from the surroundings.
      • 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
    • 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 900oC.
      • 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 0oC!
          • 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
      • 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
  • 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.).
  • At a more advanced level the heat change is called the enthalpy change is denoted by delta H, ΔH.
    • ΔH is negative (-ve) for exothermic reactions i.e. heat energy is given out and lost from the system to the surroundings which warm up.
    • ΔH is positive (+ve) for endothermic reactions i.e. heat energy is gained by the system and taken in from the surroundings which cool down OR, as is more likely, the system is heated to provide the energy needed to effect the change.
    • See later on for the bond energy arguments.

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  • Extra NOTE: In some exothermic changes, no heat is not released e.g. in batteries and fuel cells, where the energy is released as electrical energy.


 

1b. Heat changes in physical changes of state
  • Changes of physical state i.e. gas <==> liquid <==> solid are also accompanied by energy changes.

  • To melt a solid, or boil/evaporate a liquid, heat energy must be absorbed or taken in from the surroundings, so these are endothermic energy changes (ΔH +ve). The system is heated to effect these changes.

  • To condense a gas, or freeze a solid, heat energy must be removed or given out to the surroundings, so these are exothermic energy changes (ΔH -ve). The system is cooled to effect these changes.

  • PLEASE NOTE that much of section 1b. is for advanced level students NOT GCSE/IGCSE/O level students.

A comparison of energy needed to melt or boil different types of substance

  • The heat energy change (ΔH, enthalpy change) involved in a state change can be expressed in kJ/mol of substance for a fair comparison.

    • In the table below

    • ΔHmelt is the energy needed to melt 1 mole of the substance (formula mass in g) and is known as the enthalpy of fusion.

    • ΔHvap is the energy needed to vaporise by evaporation or boiling 1 mole of the substance (formula mass in g) and is known as the enthalpy of vaporization.

  • The energy required to boil or evaporate a substance is usually much more than that required to melt the solid.

    • This is because in a liquid the particles are still quite close together with attractive forces holding together the liquid, but in a gas the particles of the structure must be completely separated with virtually no attraction between them.

  • The stronger the forces between the individual molecules, atoms or ions, the more energy is needed to melt or boil the substance.

    • As this is shown by the varying energy requirements to melt or boil a substance.

  • For simple covalent molecules, the energy absorbed by the material is relatively small to melt or vaporise the substance and the bigger the molecule the greater the inter-molecular forces.

    • These forces are weak compared to those that hold the atoms together in the molecule itself.

  • For strongly bonded 3D networks e.g. (i) an ionically bonded lattice of ions, (ii) a covalently bonded lattice of atoms or (iii) a metal lattice of ions and free outer electrons, the structures are much stronger in a continuous way throughout the structure and consequently much greater energies are required to melt or vaporise the material. See structure and bonding notes

    • Note

    • Enthalpy of fusion ΔHmelt is also known as the 'latent heat of melting'.

    • Enthalpy of vaporisation ΔHvap is also known as the 'latent heat of vapourisation'.

    • In fact all of the energy changes associated with ANY change of state are known as LATENT HEATs.

    • In other words ALL changes of state require an energy change, either latent heat energy is absorbed eg for melting and boiling, or latent heat energy is given out or removed eg freezing or condensing.

Substance formula Type of bonding, structure and attractive forces operating Melting point K (Kelvin) = oC + 273 Enthalpy of fusion ΔHmelt Boiling point K (Kelvin) = oC + 273 Enthalpy of vaporisation ΔHvap
methane CH4 small covalent molecule - very weak intermolecular forces 91K/-182oC 0.94kJ/mol 112K/-161oC 8.2kJ/mol
ethanol  ('alcohol') C2H5OH larger covalent molecule than methane, greater, but still weak intermolecular forces 156K/-117oC 4.6kJ/mol 352K/79oC 43.5kJ/mol
sodium chloride Na+Cl- ionic lattice, very strong 3D ionic bonding due to attraction between (+) and (-) ions 1074K/801oC 29kJ/mol 1740K/1467oC 171kJ/mol
iron Fe strong 3D bonding by attraction of metal ions (+) with free outer electrons (-) 1808K/1535oC 15.4kJ/mol 3023K/2750oC 351kJ/mol
silicon dioxide (silica) SiO2 giant covalent structure, strong continuous 3D bond network 1883K/1610oC 46.4kJ/mol 2503K/2230oC 439kJ/mol

See Gases, Liquids and Solids notes & Structure and bonding notes for more details on structure and physical properties.


 

2. Reversible Reactions and energy changes

  • 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)    CuSO4(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.
    • 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.
  • For more on reversible reactions and chemical equilibrium see GCSE notes OR Advanced Level Notes.


Sub-index for ENERGY CHANGES: 1. Heat changes in chemical/physical changes - exothermic and endothermic  *  2. Reversible reactions and energy changes  *  3. Activation energy and reaction profiles  *  4. Catalysts and activation energy  *  5. Introduction to bond energy/enthalpy calculations  *  6. Calorimeter methods of determining energy changes  *  7. Energy transfer calculations from calorimeter results


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