• EXOTHERMIC 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 solution e.g. diluting concentrated sulphuric acid.

      • At KS3-GCSE level this exothermic 'dissolving' is considered a physical change, but at AS-A2 level they may be considered a chemical change too!

• At a higher level of thinking for exothermic chemical changes: 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. See "Energy Changes" page for more details.

• A burning or combustion reaction 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

    • CH_4(g) + 2O_2(g) ==> CO_2(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.

      • 2CH_4(g) + 3O_2(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 CH_4(g) + O_2(g) ==> C_(s) + 2H₂O_(l) (soot formation)

      • 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) + 3Cl_2(g) ==> 2AlCl_3(s) 

  • (iii) burning magnesium ribbon with a bright white flame ...

    • magnesium + oxygen ==> magnesium oxide

    • 2Mg_(s) + O_2(g) ==> 2MgO_(s) 

  • (i) to (iii) are all oxidation reactions, as is the burning of all fuels.

• 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 or 2H_2(g) + O_2(g) ==> 2H₂O_(l)

    • In all these cases the high temperature reaction zone is seen as flame and an initial high energy source for ignition 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 or C_(s) + O_2(g) ==> CO_2(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 or 2C_(s) + O_2(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 reaction.

    • 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.

• See other web page for more detailed notes on energy changes and calculations.

• BUT many exothermic reactions are not as dramatic as burning with a flame! e.g.

  • (iv) Respiration: the relatively slow 'burning' of carbohydrates in animals/plants, but it releases plenty of energy at 37^oC!

    • glucose + oxygen ==> carbon dioxide + water + energy

    • C₆H₁₂O_6(aq) + 6O_2(g) ==> 6CO_2(g) + 6H₂O_(l) + energy

  • (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 10^oC! A bunsen flame reaches 1200^oC 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.

• ENDOTHERMIC CHANGES

  • 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 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 energy ==> C₆H₁₂O₆ + 6O₂ 

    • (ii) Making lime by heating limestone to over 900^oC where a net input/absorption of energy is needed to bring about this thermal decomposition ...

      • limestone ==> quicklime + carbon dioxide

      • calcium carbonate ==> calcium oxide + carbon dioxide

      • CaCO_3(s) ==> CaO_(s) + CO_2(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 solution e.g. dissolving ammonium nitrate salt in water.

      • At KS3-GCSE level this endothermic 'dissolving' is considered a physical change, but at AS-A2 level they may be considered a chemical change too!

• At a higher level of thinking for endothermic chemical changes. The net energy change when the energy needed to break bonds in the reactants is more than the energy released when new bonds are formed in the products.

• See other web page for more detailed notes on energy changes and calculations.

• Decomposition in general means to break down into small species e.g. natural organic matter decomposes with enzymes into carbon dioxide, water and nitrogen etc.

  • Fermentation is form of biological degradation, catalysed by enzymes, to break down glucose sugar into the smaller molecules of ethanol ('alcohol') and carbon dioxide ...

    • C₆H₁₂O_6(aq) ==> 2C₂H₅OH_(aq) + 2CO_2(g)

• Light can cause decomposition e.g. in photography is a sort of photo-decomposition

  • silver chloride + light ==> silver + chlorine

  • 2AgCl ==> 2Ag + Cl₂

• Thermal decomposition means to break down substances into two or more substances by heat (usually endothermic reactions) e.g.

  • The decomposition of calcium carbonate (limestone) into calcium oxide (lime) and carbon dioxide in a high temperature lime kiln.

  • CaCO_3(s) ==> CaO_(s) + CO_2(g)

    • For more details see the Extra Industrial Chemistry notes.

  • The breaking down of hydrocarbons into smaller ones using a catalyst as well as a high temperature. This reaction is also known as cracking.

    • For more details see the Oil and its useful Products notes.

  • e.g.  C₈H₁₈ ==> C₆H₁₄ + C₂H_{4 (making ethene)}

  • Other examples in reversible reactions.

(1) Addition polymers are formed by (e.g. alkene) monomers adding together and forming no other products except the polymer e.g. ethene ==> poly(ethene), phenylethene ==> poly(phenylethene), old name polystyrene.

(2) Condensation polymers are formed by one or more monomers add together, forming the polymer BUT in forming the polymer small molecules are eliminated 'between' the monomers e.g.

dicarboxylic acid + diol ==> polyester + water, diamine + dicarboxylic acid ==> nylon + water

+ small molecules eliminated

In the case of Nylon, for each 'red' monomer - 'blue' monomer, a link is formed at each end of each monomer molecule by eliminating a water molecule e.g. where [R] = 'rest of molecule' a single link formation reaction can be shown as

[R]-COOH + HO-[R] ==> [R]-CO-O-[R] + H₂O

• NEUTRALISATION usually involves mixing an acid (pH <7 if soluble) with a base or alkali (pH > 7 if soluble) which react to form a neutral salt solution of pH7 (more details  of these types of reactions).

• Two situations are common:

• (1) Water soluble bases, called alkalis and often hydroxides, are mixed with a soluble acid such as hydrochloric, citric, sulphuric or nitric acid.

  • acid + base/alkali ==> salt + water

  • e.g. sodium hydroxide + hydrochloric acid ==> sodium chloride + water

    • NaOH_(aq) + HCl_(aq) ==> NaCl_(aq) + H₂O_(l)

  • certain carbonates like sodium carbonate, are also soluble to form alkaline solutions, and they will be similarly neutralised with 'fizzing' as carbon dioxide is formed as a 3rd product

  • e.g. sodium carbonate + hydrochloric acid ==> sodium chloride + water + carbon dioxide

    • Na₂CO_3(aq) + 2HCl_(aq) ==> 2NaCl_(aq) + H₂O_(l) + CO_2(g) 

• (2) Dissolving a water insoluble base (often an oxide) in an acid

  • e.g. copper oxide + sulphuric acid ==> copper sulphate + water

    • CuO_(s) + H₂SO_4(aq) ==> CuSO_4(aq) + H₂O_(l)

  • the acid can also be neutralised with a metal or a carbonate to give a salt solution

  • metal + acid ==> salt + hydrogen (this is also a redox reaction)

  • e.g. zinc + hydrochloric acid ==> zinc chloride + hydrogen

    • Zn_(s) + 2HCl_(aq) ==> ZnCl_2(aq) + H_2(g) 

  • insoluble carbonate + acid ==> (often soluble) salt + water + carbon dioxide

  • e.g. magnesium carbonate + sulphuric acid ==> magnesium sulphate + water + carbon dioxide

    • MgCO_3(s) + H₂SO_4(aq) ==> MgSO_4(aq) + H₂O_(l) + CO_2(g) 

• A precipitation reaction is generally defined as 'the formation of an insoluble solid on mixing two solutions or a gas bubbled into a solution'.

• The silver halide salts are used in photography and can be made by precipitation on mixing solutions of two soluble salts e.g.

  • silver nitrate + potassium chloride ==> silver chloride + potassium nitrate

  • AgNO_3(aq) + KCl_(aq) ==> AgCl_(s) + KNO3(aq)

  • Ag⁺_(aq) + Cl^-_(aq) ==> AgCl_(s)

    • is the ionic equation (NO₃^- and K⁺ are spectator ions)

• Insoluble barium sulphate can be made in a similar manner ..

  • barium chloride + dil. sulphuric acid ==> barium sulphate + dil. hydrochloric acid

  • BaCl_2(aq) + H₂SO_4(aq) ==> BaSO_4(s) + 2HCl_(aq)

  • Ba²⁺_(aq) + SO₄^2-_(aq) ==> BaSO_4(s)

    • is the ionic equation (Cl^- and H⁺ are spectator ions)

Reversible Reactions

• Generally speaking most chemical reactions are irreversible, that mean they go 100% one way to the products and that's that! e.g. magnesium fizzing away in hydrochloric acid to form magnesium chloride and hydrogen. There is no feasible way of reacting hydrogen and magnesium chloride to re-form magnesium metal and hydrochloric acid!

• However a reversible reaction is a chemical reaction in which the products can be converted back to reactants under suitable conditions and there are quite a few examples of them.

• A reversible reaction is shown by the sign i.e. a half-arrow to the right (forward reaction) and a half-arrow to the left (backward reaction).

• As pointed out above, most reactions are not reversible and have the usual complete arrow only pointing to the right i.e. in the direction the reaction goes 100%.

Two examples of reversible reactions are given below: (more details for GCSE)

(a) The thermal decomposition of ammonium chloride.

On heating strongly, the white solid ammonium chloride, decomposes into a mixture of two colourless gases - ammonia and hydrogen chloride. On cooling the reaction is reversed and solid ammonium chloride reforms.

Ammonium chloride + heat ammonia + hydrogen chloride

NH₄Cl_(s) NH_3(g) + HCl_(g)

(b) The thermal decomposition of hydrated copper(II) sulphate.

• On heating the blue solid, hydrated copper(II) sulphate, steam is given off and the white solid of anhydrous copper(II) sulphate is formed (left to right reaction).

• When the white solid is cooled and water added, blue hydrated copper(II) sulphate is reformed (right to left reaction).

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

• When a reversible reaction occurs in a closed system an equilibrium is formed, in which the original reactants and products formed coexist.

• In an equilibrium there is a state of balance between the concentrations of the reactants and products.

• At equilibrium the rate at which the reactants change into products is exactly equal to the rate at which the products change back to the original reactants.

• The result is that that the concentrations of the reactants and products remain the same BUT the reactions don't stop!

• However the relative amounts of the reactants and products depend on the reaction conditions e.g. the temperature and pressure.

• You can change the conditions to favour a particular reaction direction e.g. in a limekiln the carbon dioxide is vented out so all the limestone changes to lime.

calcium carbonate (limestone) calcium oxide (lime) + carbon dioxide

nitrogen + hydrogen ammonia

• Hydration means the addition of water or combining with water.

• Dehydration means the losing or removal of water.

• Example 1: On heating the blue solid, hydrated copper(II) sulphate, steam is given off and the white solid of anhydrous copper(II) sulphate is formed (left to right reaction is a dehydration, water lost). When the white solid is cooled and water added, blue hydrated copper(II) sulphate is reformed (right to left reaction is a hydration, water gained).

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

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

    • The water in the blue crystals (the 5H₂O) is called water of crystallisation and becomes part of the crystal structure when a concentrated solution of copper sulphate is slowly evaporated to allow crystallisation e.g. in a salt preparation.

    • The left to right dehydration reaction also happens if blue hydrated copper sulphate is treated with cold/warm concentrated sulphuric acid, which is a powerful dehydrating agent (see sugar reaction further down in this section).

• Example 2: If the alcohol ethanol is heated with concentrated sulphuric acid it is dehydrated to form the alkene gas ethene plus water. The same reaction happens if ethanol vapour is passed over very hot aluminium oxide which also catalyses the dehydration of the ethanol.

  • ethanol ==> ethene + water

  • ==> + H₂O

• Example 3: If ethene gas is dissolved in concentrated sulphuric acid and diluted with water, on gentle boiling the alcohol ethanol is formed (the reverse of example 2). The same reaction occurs if ethene gas and water vapour are passed over a silica gel-phosphoric acid catalyst at 300^oC. In both cases the ethene is hydrated to form ethanol by water addition.

  • ethene + water ==> ethanol

  • + H₂O ==>

• Example 4: If white crystalline sugar is heated with concentrated sulphuric acid a spongy mass of carbon is formed. The acid dehydrates the sugar, i.e. it removes the equivalent of 'H₂O' and leaves the 'C' atoms!

  • e.g. glucose ==> carbon + water

  • C₆H₁₂O₆ ==> 6C + 6H₂O

• Displacement reactions occur when a more reactive metal displaces a less reactive metal from one of its compounds e.g. from the oxide, sulphate, chloride or nitrate.

• Displacement reactions also occur when a more reactive non-metal displaces a less reactive non-metal from one of its compounds e.g. from the metal salt.

• These experiments can be used to establish a so called 'Reactivity Series'.

Metal displacements

• e.g. the spectacular very exothermic 'thermit reaction' which shows aluminium is more reactive than iron.  When a mixture of aluminium and iron(III) oxide is ignited with a magnesium fuse (high activation energy), the mixture burns furiously with a shower of sparks to leave a red hot blob of iron and a white as of aluminium oxide ...

• iron(III) oxide + aluminium ==> aluminium oxide + iron

• Fe₂O_3(s) + 2Al_(s) ==> Al₂O_3(s) + 2Fe_(s)  

• or more reactive copper 'gently' displaces silver from silver nitrate solution to make silvery plated copper

• copper + silver nitrate ==> copper(II) nitrate + silver

• Cu_(s) + 2AgNO_3(aq) ==> 2Ag + Cu(NO₃)_2(aq)

• more details in Metal Reactivity Notes

• e.g. halogen displacement, more reactive chlorine displaces bromine from potassium bromide

• chlorine + potassium bromide ==> potassium chloride + bromine

• Cl_2(aq) + 2KBr_(aq) ==> 2KCl_(aq) + Br_2(aq) 

• or more reactive bromine displaces iodine from sodium iodide

• bromine + sodium iodide ==> sodium bromide + iodine

• Br_2(aq) + 2NaI_(aq) ==> 2NaCl_(aq) + I_2(aq)

• more details in Halogen Notes 

Galvanising

This means to coat iron or steel with a layer of zinc to stop it rusting (more details on Metal Reactivity page)

 Haber Process

• Part of the manufacture of sulphuric acid from:  (all the details)
•  sulphur ==> sulphur dioxide ==V catalyst==> sulphur trioxide* ==> sulphuric acid. (
• S_(s) + O_{2 (g)} ==> SO_{2 (g)}, * 2SO_{2 (g)} + O_{2 (g)} ==> 2SO_{3 (g)},
• followed by, indirectly, SO₃ + H₂O ==> H₂SO₄

• e.g. if you mix solutions of sodium chromate with lead nitrate you get a yellow precipitate of lead chromate and sodium nitrate is left in solution.
• sodium chromate + lead nitrate ==> lead chromate + sodium nitrate
• Na₂CrO_4(aq) + Pb(NO₃)_2(aq) ==> PbCrO_4(s) + 2NaNO_3(aq)
• This is the yellow pigment Van Gogh used in his paintings and you see it as the road markings you don't park on!

• One way of reducing pollutants from a car exhaust is to use transition metal catalysts (Pt+Rh set on a heat resistant base = the catalytic converter). A platinum/rhodium catalyst converts toxic carbon monoxide and nitrogen monoxide gases into harmless carbon dioxide and nitrogen gases.
• 2CO _(g) + 2NO _(g) == Pt/Rh ==> 2CO_{2 (g)} + N_{2 (g)} 

• Combining an organic carboxylic acid with an alcohol, produces an pleasant smelling ester (which are used in perfumes and flavourings) and water.
  • e.g. ethanoic acid + ethanol ethyl ethanoate + water
  • +   + H₂O
  • Its an equilibrium, ²/₃ rds conversion, and the reaction is catalysed by a few drops of concentrated sulphuric acid.

• Iron (or steel) corrodes more quickly than most other transition metals and readily does so in the presence of both oxygen (in air) and water to form an iron oxide. This means rusting is an oxidation reaction.
• iron + oxygen + water ==> hydrated iron(III) oxide
• 4Fe_(s) + 3O_2(g) + xH₂O_(l) ==> 2Fe₂O₃.xH₂O_(s)
• i.e. rust is an orange-brown solid of hydrated iron(III) oxide formed from the reaction with oxygen and water (the equation is not meant to be balanced and the amount of water x is variable, from dry to soggy!).
• For more details of the chemistry of rusting and its prevention go to the corrosion section on the Metal Reactivity Series page.

• A substitution reaction is where one part of a molecule is replaced by something else.
• e.g. when chlorine reacts with methane, a hydrogen atom in methane is replaced by a chlorine atom from the chlorine molecule.
• methane + chlorine ==> chloromethane + hydrogen chloride
• CH₄ + Cl₂ ==> CH₃Cl + HCl

• An addition reaction is when one molecule adds to another molecule resulting in a single product e.g.
• ethene + bromine ==> 1,2-dibromoethane
• C₂H₄ + Br₂ ==> C₂H₄Br₂
• 

Deliquescent and Hygroscopic

(overlapping terms, physical changes rather than chemical changes)

Deliquescent

• A substance is described as showing 'deliquescence' if it absorbs so much water from the air it forms a solution i.e. dissolves in the absorbed water.

• Examples: the salts calcium chloride and potassium fluoride.

• It is the extreme case of being hygroscopic (below).

• At a higher level, technically it has a lower water vapour pressure than the surrounding air, so there is a net gain in water. 

Hygroscopic

• A substance is described as hygroscopic if it absorbs water vapour from the air to form a damp or moist solid or even a solution.

• These substances can be used as drying agents (dessicants) for air or liquids.

•  e.g. the anhydrous salts calcium chloride and sodium sulphate.

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