Exothermic and Endothermic Reactions or Changes

    • Heat is released or given out to the surroundings by the materials involved, so
      the temperature rises.
    • chemical change examples involve a new substance being formed and lots of
      examples (i) to (vi) below (but they are not always exothermic –
      see endothermic below).
    • physical change examples
      e.g. condensation, freezing etc. all require the removal of energy from
      the material e.g. water, to the surroundings to produce the change in
      state (its the same as releasing heat, but it doesn’t seem like it!).
      • Note:
        Dissolving substances in water can release heat giving a warm/hot
        e.g. diluting concentrated sulphuric acid.
  • At a higher level of thinking for exothermic chemical
    The net energy change when
    the energy needed to break bonds in the reactants is less than the
    energy released when new bonds are formed in the products.
  • A burning or combustion
    usually means a very fast exothermic
    reaction where a flame is observed. It involves a highly energetic
    oxidation of ‘fuels’ where the temperature generated is so high the
    atoms give off light from the luminous flame zone  e.g.
    • (i) bunsen flame as methane gas fuel burns …
      • methane + oxygen ==> carbon dioxide + water
      • CH4(g) + 2O2(g)==> CO2(g) + 2H2O(l) 
      • This is complete combustion with a pale blue
        flame and the products cannot react any further with oxygen.
      • If the oxygen supply is limited the flame is more
        yellow and can be ‘smokey’ due to soot formation (C) and dangerous since
        carbon monoxide (CO) can be formed.
      • These are examples of incomplete combustion.
        • methane + oxygen ==> carbon monoxide + water
        • 2CH4(g) + 3O2(g)==> 2CO(g) + 4H2O(l) (carbon
          monoxide formation)
          • Most people who die in house fires are poisoned
            by carbon monoxide (and other toxic gases) in the thick smoke rather
            than from burns.
        • or
        • methane + oxygen ==> carbon (soot) +
        • CH4(g) + O2(g)==> C(s) + 2H2O(l) (soot
        • The sooty carbon particles e.g. in a candle flame,
          are heated to such a high temperature they become incandescent and
          give out yellow light, but as far as I know virtually no carbon
          monoxide is formed!
    • (ii) passing chlorine over hot aluminium metal to
      make aluminium chloride, the aluminium burns to form the chloride …
      • aluminium + chlorine ==>aluminium chloride
      • 2Al(s) + 3Cl2(g)==> 2AlCl3(s) 
    • (iii) burning magnesium ribbon with a bright white
      flame …
      • magnesium + oxygen ==> magnesium oxide
      • 2Mg(s) + O2(g)==> 2MgO(s) 
    • (i) to (iii) are all oxidation reactions, as
      are all ‘fuel’ burning reactions.
  • Continuous combustion requires the ‘fire triangle
    of heat + fuel + oxidant (oxidants like oxygen, air or other
    reactive gases like chlorine or fluorine and in rockets liquids like
    hydrogen peroxide)
    • Very fast or explosive combustion:
      • A roaring bunsen flame (of methane burning) is an
        example of fast combustion and when the air (oxygen) – methane
        (natural gas) mixture is first ignited it is a small explosion! (equation
        above). It seems contradictory, but a source of ignition is needed
        because the C-H and O=O bonds are very strong giving a high activation
        energy. However, once ignited, the heat from the flame keeps the
        burning going.
      • Another explosive example is the ‘squeaky pop test
        for hydrogen’. When a lit splint is applied there is a faint blue
        flame for a fraction of a second as the two gases explode to form
        water + heat, light and sound energy!
        • hydrogen + oxygen ==> water
        • 2H2(g) + O2(g)==> 2H2O(l)
      • In all these cases the high temperature reaction
        zone is seen as flame and an initial high energy source for
        is needed to initiate the reaction e.g. a match or an
        electrical discharge.
    • Slow or smouldering combustion:
      • In these cases no flame is seen, but a high
        temperature heat source is still required to start the reaction and
        the reaction zone is still at a high temperature e.g. the red hot slow
        burning of charcoal (mainly carbon), but the main combustion product
        is still carbon dioxide. You can only get this slow/smouldering
        combustion with solid combustible reactants.
      • Gases will tend to explode unless controlled in a
        burner and liquids will vaporise in the heat from the flame and so
        will also burn very fast with a flame.
        • carbon + oxygen ==> carbon dioxide
          • C(s) + O2(g)==> CO2(g)
          • This is an example of complete combustion.
        • BUT quite often, with limited air/oxygen supply,
          carbon monoxide is readily formed,
        • carbon + oxygen ==> carbon monoxide
          • 2C(s) + O2(g)==> 2CO(g)
          • This is an example of incomplete combustion.
    • Spontaneous combustion:
      • This is when combustion occurs without any
        application of a high energy ignition source, sometimes described as self-ignition,
        though in some cases heat is generated in some way which triggers the
      • For example, it is possible to prepare a very
        finely divided black powder form of iron(II) oxide. When the powder is
        dropped through air lots of tiny flashes of light are seen as it burns
        to form another iron oxide (probably Fe3O4). The
        reaction is triggered by heat from friction. The powder has such a
        large surface area that the friction caused by just falling in air
        produces enough heat to initiate the reaction. Powdered coal dust or
        very fine flour can behave in the same way and both have been
        responsible for serious accidents in industry.
      • Other substances can spontaneously ignite in air
        because the activation energies required are so low and the
        kinetic energy of the particles is sufficient for the reaction to
        happen without help! e.g. the highly reactive Group 1 Alkali Metal
        caesium and the silicon-hydrogen compounds called silanes(SiH4,
        Si2H6 etc. which are like organic alkanes with
        the C’s replaced by Si and far less stable).
      • Potassium and all the alkali metals below it ignite
        in water (Rb and Cs explosively) because the reaction is so exothermic
        and ignites metal vapour and the hydrogen gas produced.+.
  • BUT many exothermic reactions are not as dramatic as
    burning with a flame!
    • (iv) Respiration: the
      relatively slow ‘burning’ of carbohydrates in animals/plants, but it
      releases plenty of energy at 37oC!
      • glucose + oxygen ==> carbon dioxide + water +
      • C6H12O6(aq) + 6O2(g) ==>6CO2(g) + 6H2O(l) +
    • (v) Neutralisation: acid + alkali ==> salt
      + water
      • e.g. hydrochloric acid + sodium hydroxide ==>
        sodium chloride + water
      • which is one of the fastest reactions in water, but
        the mixture only warms up by 5 to 10oC! A bunsen flame
        reaches 1200oC in the main combustion zone!
      • More details further down.
    • (vi) Rusting in which iron slowly
      reacts with water and oxygen (from air) to form the orange-brown hydrated
      iron oxide
      we call rust.
    • Heat is absorbed or taken in by the materials involved from the surroundings,
      the system cools or has to be heated to effect the change.
    • Chemical change examples
      of endothermic reactions e.g. thermal decomposition of limestone,
      cracking oil fractions, decomposition by electrolysis etc.
      • (i) Photosynthesis:
        input of energy from sunlight needed
        • carbon dioxide + water ==> glucose + oxygen
        • 6CO2 + 6H2O + sunlight
          + 6O2
      • (ii) Making lime by heating limestone to
        over 900oC where a net input/absorption of energy is needed
        to bring about this thermal decomposition …
        • limestone ==> quicklime + carbon dioxide
        • calcium carbonate ==> calcium oxide + carbon
        • CaCO3(s) ==>CaO(s) + CO2(g) 
      • (iii) Cracking hydrocarbon molecules from
        oil to make smaller molecules, also requires this absorption of heat
        by the reactant molecules to break em’ up’, or to put it ‘poshly’,
        another example thermal decomposition e.g.
        • hexane =>ethene + butane
        • C6H14 ==>C2H4 + C4H10 
    • Physical change examples
      e.g. melting, boiling, evaporation etc. all require the input of energy
      to effect the change of state of the material.
      • Note:
        Dissolving substances in water can absorb heat giving a cool
        e.g. dissolving ammonium nitrate salt in water.