NEW site: GCSE Revision Notes on REACTIVITY SERIES of METALS for IGCSE-KS4 at Doc Brown's Chemistry Clinic. reactions reaction equation equations of potassium sodium calcium magnesium aluminium zinc iron lead copper with air oxygen water dilute/concentrated hydrochloric sulphuric/sulfuric nitric acid/acids rusting and corrosion causes and prevention courses docbreactivity

docbreactivity updated April 25th 2008

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KS4 Science GCSE-IGCSE chemistry revision notes on

The Metal Reactivity Series, the Corrosion of Metals like iron and Redox Reactions

The page also includes an introduction to displacement and oxidation-reduction (redox) reactions

Associated GCSE notes pages on-site: The Periodic Table * Group 1 Alkali Metals * Metal extraction * Transition metals * Alloys-uses of metals * Electrochemistry * Rates of Reactions (e.g. metal-acid) and EMAIL query?comment

reactivity For a summary of the metals chemical reactions with air/oxygen, acids and oxides/salts (displacement), including word equations and balanced symbol equations, all in the context of the reactivity series just click on its name from this alphabetical order list ... aluminium .. caesium .. calcium .. copper .. francium .. gold .. iron .. lead .. lithium .. magnesium .. platinum .. potassium .. rubidium .. silver .. sodium .. tin .. zinc (but the notes are in reactivity order) and other sub-sections on this page: METAL CORROSION- RUSTING * DISPLACEMENT REACTIONS and OXIDATION - REDUCTION explained

Easy KS3 science multiple choice quiz start on metal reactivity and KS3 word-fills

GCSE m/c QUIZ: Foundation or Higher & GCSE reactivity word-fill or Rusting word-fill

The higher the metal in the series, the more reactive it is i.e. the faster, more vigorous and more exothermic the reaction.
At a more theoretical level, the more reactive a metal, the greater tendency it has to form a positive ion in the context of a chemical reaction (e.g. Na ==> Na⁺ or Fe ==> Fe²⁺).
This also implies that the reverse reaction becomes more difficult i.e. the more reactive a metal, the more difficult it is to extract from its ore and the metal is also more susceptible to corrosion with oxygen and water.
The reactivity series can be established by observation of the reaction of metals with water, oxygen or acids (and also from simple cell experiments).
DISPLACEMENT REACTIONS:
- A metal in the series, can displace any metal below it in the series, from the less reactive metal's oxide, chloride or sulphate or other compound.

e.g. on heating the mixture of a metal and another metal oxide, such as magnesium powder and black copper(II) oxide, a very exothermic reaction occurs in a shower of sparks and white magnesium oxide is formed with brown bits of copper:
- magnesium + copper oxide ==> magnesium oxide + copper
- Mg_(s) + CuO_(s) ==> MgO_(s) + Cu_(s)
- The more reactive magnesium displaces the less reactive copper as it does in the 2nd example below.
or adding a metal to a salt solution of another metal e.g. adding magnesium to blue copper(II) sulphate solution, the blue colour fades as colourless magnesium sulphate is formed and brown bits of copper metal form a precipitate:
- magnesium + copper sulphate ==> magnesium sulphate + copper
- Mg_(s) + CuSO_4(aq) ==> MgSO_4(aq) + Cu_(s)
- The electron transfer redox theory behind displacement reactions is explained later.
If no reaction happens, then it means the added metal is less reactive than the metal in the oxide or sulphate etc.
See also the Thermit reaction.

Some general word equations where the metal does react:

(a) metal + cold water ==> metal hydroxide + hydrogen (metals above aluminium)
(b) heated metal + steam ==> metal oxide + hydrogen (for metals above tin?)
(c) metal + acid ==> metal salt + hydrogen

if the metal is at least as reactive as lead (see reactivity series list above left)
and hydrochloric acid makes a metal chloride salt,
and sulphuric acid makes a metal sulphate salt,
reactions with nitric acid are complex, the nitrate salt, is formed BUT the gas is rarely hydrogen, and more often an oxide of nitrogen (not usually studied at GCSE level these days).
- Oxides of nitrogen note: NO is nitrogen(II) oxide [old names nitrogen monoxide or nitric oxide] and NO₂ is nitrogen(IV oxide [old name nitrogen dioxide]
The electron transfer redox theory behind metal-acid reactions is explained later.

Within the general Reactivity or Activity Series of Metals there are some Periodic Table Trends …

Down Group 1 (I) the "Alkali Metals" the activity increases Cs > Rb > K > Na > Li
Down Group 2 (II) the activity increases e.g. Ca > Mg
On the same period, the Group 1 metal is more reactive than the group 2 metal, and the group 2 metal is more reactive than the Group 3 metal, and all three are more reactive than the "Transition Metals". e.g. Na > Mg > Al (on Period 3) and K > Ca > Ga > Fe/Cu/Zn etc. (on Period 4)

The reactivity of a metal has an important bearing on the method by which a metal is extracted from its ore. Since prehistoric times, as technology has improved more and more, all metals can now be extracted and comments on when the metals were first isolated and used are added in the table below. If the metal is above carbon, it cannot be extracted by carbon reduction and must be usually extracted by electrolysis.
Two non-metals, carbon and hydrogen, are included in the table for comparison, and are important chemical reference points concerning the method of metal extraction and reactivity towards acids

Metals above carbon cannot usually be extracted by carbon or carbon monoxide reduction and are usually extracted by electrolysis
Metals below carbon in the series can be extracted by heating the oxide with carbon or carbon monoxide.
Metals below hydrogen will not displace hydrogen from acids and can be extracted by heating the oxide in hydrogen.

Notes on the corrosion of metals and the prevention of iron rusting are dealt with at the end of the page.
The theory of OXIDATION and REDUCTION and their application to REDOX reactions are also dealt with at the end of the page.
A brief note on some of the uses of reactive metals and their compounds relating to this page:
- Many of the metals give bright flame colours when burnt in air, and the same colours are seen when a compound of the metal is heated strongly in a bunsen flame e.g. calcium/lithium give red, sodium yellow. So their compounds are used in fireworks and magnesium powder burns brightly with a brilliant white flame is also used in fireworks and flares.
- The displacement of a less reactive metal from its compound by a more reactive metal is used to extract metals e.g. chromium is obtained from a chromium oxide by a thermit type reaction using more reactive aluminium and titanium is released from titanium chloride by heating it with highly reactive sodium or magnesium metal.
  - see notes on Extraction of Metals and Making metals more useful on other web pages.
- Rusting is prevented by coating iron and steel with a more reactive metal like zinc which is preferentially/sacrificially corroded away and sparing the iron/steel. (see corrosion notes on this page)

METAL in decreasing reactivity order (and where it is in the Periodic Table)	Reactivity and Reactions The compounds formed in the reactions are white insoluble solids, (s), or soluble colourless solutions, (aq), unless otherwise stated. Some modern systematic names and 'old names' are given in square brackets [], though these are usually only needed by advanced level students.
francium Fr Group 1 Alkali Metal	Theoretically Francium, in the Group 1 Alkali Metals, is the most reactive of any metal and therefore the most explosive metal when in contact with water, however, it is also very radioactive and so the experiment is highly unlikely to be performed!
caesium Cs Group 1 Alkali Metal	Caesium is so reactive, that when a lump is freshly cut, although you see at first the typical silvery metallic lustre of the pure metal, it rapidly tarnishes-oxidises at room temperature by reaction with the oxygen in air. It forms successively the oxide, the hydroxide from water vapour in the air, and then the carbonate from carbon dioxide in the air. That's why if an 'old' lump is picked out from the bottle where it is stored under oil (because of its reactivity), it is encrusted with a white layer of these compounds. Burns vigorously with a blue flame when heated in air/oxygen to form the white powder caesium oxide. caesium + oxygen ==> caesium oxide 4Cs_(s) + O_2(g) ==> 2Cs₂O_(s) also forms caesium peroxide, Cs₂O₂ and caesium superoxide, CsO₂ Because it is extremely reactive, it reacts and explodes violently with cold water forming the alkali caesium hydroxide and flammable-explosive hydrogen gas. caesium + water ==> caesium hydroxide + hydrogen 2Cs_(s) + 2H₂O_(l) ==> 2CsOH_(aq) + H_2(g) Caesium was first extracted in 1860 by electrolysis of the molten chloride CsCl.
rubidium Rb Group 1 Alkali Metal	Rubidium is so reactive, that when a lump is freshly cut, although you see at first the typical silvery metallic lustre of the pure metal, it rapidly tarnishes-oxidises at room temperature by reaction with the oxygen in air. It forms successively the oxide, the hydroxide from water vapour in the air, and then the carbonate from carbon dioxide in the air. That's why if an 'old' lump is picked out from the bottle where it is stored under oil (because of its reactivity), it is encrusted with a white layer of these compounds. Burns vigorously with a red flame when heated in air/oxygen to form the white powder rubidium oxide. rubidium + oxygen ==> rubidium oxide 4Rb_(s) + O_2(g) ==> 2Rb₂O_(s) also forms rubidium peroxide, Rb₂O₂ and rubidium superoxide, RbO₂ Extremely reactive, can ignite in air, it reacts and explodes violently with cold water forming the alkali rubidium hydroxide and flammable-explosive hydrogen gas. rubidium + water ==> rubidium hydroxide + hydrogen 2Rb_(s) + 2H₂O_(l) ==> 2RbOH_(aq) + H_2(g) Rubidium was first extracted in 1861 by electrolysis of the molten chloride RbCl.
potassium K Group 1 Alkali Metal	Potassium is so reactive, that when a lump is freshly cut, although you see at first the typical silvery metallic lustre of the pure metal, it rapidly tarnishes-oxidises at room temperature by reaction with the oxygen in air. It forms successively the oxide, the hydroxide from water vapour in the air, and then the carbonate from carbon dioxide in the air. That's why if an 'old' lump is picked out from the bottle where it is stored under oil (because of its reactivity), it is encrusted with a white layer of these compounds. Potassium burns vigorously with a lilac flame when heated in air/oxygen to form the white powder potassium oxide. potassium + oxygen ==> potassium oxide 4K_(s) + O_2(g) ==> 2K₂O_(s) also forms potassium peroxide, K₂O₂ and potassium superoxide, KO₂ Potassium is very reactive with water - the reaction is the same as for sodium (full description below) BUT it is faster and more exothermic AND so the hydrogen is ignited to give a purple or lilac flame. The hydrogen flame is coloured by the excitation of potassium atoms in the very hot flame (e.g. as in the flame test for potassium, yellow for sodium in the next section). The very rapid reaction with cold water forms the alkali potassium hydroxide and flammable-explosive hydrogen gas. potassium + water ==> potassium hydroxide + hydrogen 2K_(s) + 2H₂O_(l) ==> 2KOH_(aq) + H_2(g) Potassium was first extracted in 1807 by electrolysis of the molten chloride KCl.
sodium Na Group 1 Alkali Metal	Sodium is so reactive, that when a lump is freshly cut, although you see at first the typical silvery metallic lustre of the pure metal, it rapidly tarnishes-oxidises at room temperature by reaction with the oxygen in air. It forms successively the oxide, the hydroxide from water vapour in the air, and then the carbonate from carbon dioxide in the air. That's why if an 'old' lump is picked out from the bottle where it is stored under oil (because of its reactivity), it is encrusted with a white layer of these compounds. Sodium burns vigorously with a yellow flame when heated in air/oxygen to form the white powder sodium oxide. sodium + oxygen ==> sodium oxide 4Na_(s) + O_2(g) ==> 2Na₂O_(s) also forms some sodium peroxide, Na₂O₂ Sodium is very reactive with water: the sodium melts to a silvery ball and fizzes as it spins over the water. The rapid exothermic reaction produces a colourless gas which gives a squeaky pop! with a lit splint (hydrogen). Universal indicator will turn from green to purple/violet as the strong alkali sodium hydroxide is formed. The initially sodium floats because it is less dense than water. sodium + water ==> sodium hydroxide + hydrogen 2Na_(s) + 2H₂O_(l) ==> 2NaOH_(aq) + H_2(g) Sodium was first extracted in 1807 by electrolysis of the molten chloride NaCl. Extraction details (see alphabetical list at top of this other page)
lithium Li Group 1 Alkali Metal	Lithium is so reactive, that when a lump is freshly cut, although you see at first the typical silvery metallic lustre of the pure metal, it rapidly tarnishes-oxidises at room temperature by reaction with the oxygen in air. It forms successively the oxide, the hydroxide from water vapour in the air, and then the carbonate from carbon dioxide in the air. That's why if an 'old' lump is picked out from the bottle where it is stored under oil (because of its reactivity), it is encrusted with a white layer of these compounds. Lithium burns vigorously with a red flame when heated in air/oxygen to form the white powder lithium oxide. lithium + oxygen ==> lithium oxide 4Li_(s) + O_2(g) ==> 2Li₂O_(s) Quite a fast reaction with cold water forming the alkali lithium hydroxide and hydrogen gas. For full description see sodium above, but the reaction is not as fast. lithium + water ==> lithium hydroxide + hydrogen 2Li_(s) + 2H₂O_(l) ==> 2LiOH_(aq) + H_2(g) Lithium was first extracted in 1821 by electrolysis of the molten chloride LiCl.
calcium Ca Group 2 Alkaline Earth Metal	Calcium burns quite fast with a brick red flame when strongly heated in air/oxygen to form the white powder calcium oxide. calcium + oxygen ==> calcium oxide 2Ca_(s) + O_2(g) ==> 2CaO_(s) Quite reactive with cold water forming the moderately soluble alkali calcium hydroxide and hydrogen gas. A white milky precipitate can develop as calcium hydroxide is only slightly soluble in water. calcium + water ==> calcium hydroxide + hydrogen Ca_(s) + 2H₂O_(l) ==> Ca(OH)_2(aq/s) + H_2(g) Very reactive with dilute hydrochloric acid forming the colourless soluble salt calcium chloride and hydrogen gas. calcium + hydrochloric acid ==> calcium chloride + hydrogen Ca_(s) + 2HCl_(aq) ==> CaCl_2(aq) + H_2(g) Not very reactive with dilute sulphuric acid because the colourless calcium sulphate formed is not very soluble and coats the metal inhibiting the reaction, so not many bubbles of hydrogen. calcium + sulphuric acid ==> calcium sulphate + hydrogen Ca_(s) + H₂SO_4(aq) ==> CaSO_4(aq/s) + H_2(g) Calcium was first extracted in 1808 by electrolysis of the molten chloride CaCl₂.
magnesium Mg Group 2 Alkaline Earth Metal	Magnesium burns vigorously with a bright white flame when strongly heated in air/oxygen to form a white powder of magnesium oxide. magnesium + oxygen ==> magnesium oxide 2Mg_(s) + O_2(g) ==> 2MgO_(s) Slow reaction with water forming the slightly soluble alkali magnesium hydroxide and hydrogen gas, a bit faster in boiling water. magnesium + water ==> magnesium hydroxide + hydrogen Mg_(s) + 2H₂O_(l) ==> Mg(OH)_2(aq/s) + H_2(g) If heated in steam, the magnesium will burn with a bright white flame and the white powder magnesium oxide is formed and hydrogen gas. magnesium + water ==> magnesium oxide + hydrogen Mg_(s) + H₂O_(g) ==> MgO_(s) + H_2(g) In fact it is so reactive, it will even burn in carbon dioxide, the products being white magnesium oxide powder and black specks of elemental carbon! magnesium + carbon dioxide ==> magnesium oxide + carbon 2Mg_(s) + CO_2(g) ==> 2MgO_(s) + C_(s) Very reactive with dilute hydrochloric acid forming the colourless soluble salt magnesium chloride and hydrogen gas. magnesium + hydrochloric acid ==> magnesium chloride + hydrogen Mg_(s) + 2HCl_(aq) ==> MgCl_2(aq) + H_2(g) Very reactive with dilute sulphuric acid forming colourless soluble magnesium sulphate and hydrogen. magnesium + sulphuric acid ==> magnesium sulphate + hydrogen Mg_(s) + H₂SO_4(aq) ==> MgSO_4(aq) + H_2(g) Magnesium nitrate Mg(NO₃)₂ and hydrogen are formed with very dilute nitric acid. However another reaction occurs simultaneously, particularly in more concentrated nitric acid, in which the gas nitrogen monoxide (nitrogen(II) oxide, NO) is formed instead of hydrogen. The colourless nitrogen monoxide rapidly combines with oxygen in air to give the dangerous irritating gas nitrogen dioxide (nitrogen(IV) oxide, NO₂). (i) magnesium + nitric acid ==> magnesium nitrate + hydrogen Mg_(s) + 2HNO_3(aq) ==> Mg(NO₃)_2(aq) + H_2(g) which competes with the reaction ... (ii) magnesium + nitric acid ==> magnesium nitrate + water + nitrogen(II) oxide [nitric oxide] 3Mg_(s) + 8HNO_3(aq) ==> 3Mg(NO₃)_2(aq) + 4H₂O_(l) + 2NO_(g) and followed rapidly by ... (iii) nitrogen(II) oxide + oxygen ==> nitrogen(IV) oxide 2NO_(g) + O_2(g) ==> 2NO_2(g) However with concentrated nitric acid, nitrogen(IV) oxide [nitrogen dioxide] is formed directly. (iv) magnesium + nitric acid ==> magnesium nitrate + water + nitrogen(IV) oxide 3Mg_(s) + 4HNO_3(aq) ==> Mg(NO₃)_2(aq) + 2H₂O_(l) + 2NO_2(g) So, whatever concentration of nitric acid is used, you get a colourless solution of magnesium nitrate AND nasty brown fumes of nitrogen dioxide. Nitric acid is a strong oxidising agent and it is also NOT a reaction on which to base magnesium's position in the 'metal reactivity series' because of the complications. Reactive magnesium gives lots of displacement reactions with the oxides and salts of less reactive metals e.g. (i) After heating a mixture of grey magnesium powder and black copper(II) oxide, the mixture burns exothermically to give white magnesium oxide and pinky-brown bits of copper magnesium + copper(II) oxide ==> magnesium oxide + copper Mg_(s) + CuO_(s) ==> MgO_(s) + Cu_(s) (ii) Adding magnesium powder to copper(II) sulphate solution, remove the blue colour of the copper(II) salt, leaving a colourless solution of magnesium sulphate and a pinky-brown deposit of copper. magnesium + copper sulphate ==> magnesium sulphate + copper Mg_(s) + CuSO_4(aq) ==> MgSO_4(aq) + Cu_(s) Magnesium was first extracted in 1808 by electrolysis of the molten chloride MgCl₂.
aluminium Al Group 3 Metal	The surface of aluminium goes white when strongly heated in air/oxygen to form white solid aluminium oxide. Theoretically its quite a reactive metal but an oxide layer is readily formed even at room temperature and this has quite an inhibiting effect on its reactivity. Even when scratched, the oxide layer rapidly reforms, which is why it appears to be less reactive than its position in the reactivity series of metals would predict but the oxide layer is so thin it is transparent, so aluminium surfaces look metallic and not a white matt surface. aluminium + oxygen ==> aluminium oxide 4Al_(s) + 3O_2(g) ==> 2Al₂O_3(s) Under 'normal circumstances' in the school laboratory aluminium has virtually no reaction with water, not even when heated in steam due to a protective aluminium oxide layer of Al₂O₃. (see above) The metal chromium behaves chemically in the same way, forming a protective layer of chromium(III) oxide, Cr₂O₃, and hence its anti-corrosion properties when used in stainless steels and chromium plating. Although this again illustrates the 'under-reactivity' of aluminium, the Thermit Reaction shows its rightful place in the reactivity series of metals. The following is NOT needed for pre-university GCSE-AS-A2 etc. chemistry students as far as I'm aware, but maybe of interest to some students, because it illustrates what happens if you dig a little deeper into what appears to be a simple experimental situation! (1) If the surface of aluminium is treated with less reactive metal salt, it is still possible to get displacement reaction. Check this out by leaving a piece of aluminium foil in copper(II) sulphate solution and a patchy pink colour of copper metal slowly appears over many hours/days?. However, as a student teacher back in 1975, I did the experiment with a mercury salt (highly nerve toxic and now use banned in UK schools) and found all of the aluminium foil reacted when left in water overnight. The next morning, after the hydrogen had 'departed', there was nothing left but a soggy mass of hydrated aluminium hydroxide! The aluminium-mercury 'couple' enables the aluminium to displace the hydrogen from water even at room temperature. You get a similar 'speeding up' effect when copper(II) sulphate solution is added to a zinc-dilute sulphuric acid mixture. However, they are not as fast and exciting as the Thermit Reaction described below! which is legal for teachers to do with suitable health and safety precautions like using a transparent safety barrier and goggles and sending the class to the back of the room! (2) I am informed that water will react with molten aluminium (email from jg?) because in the bulk of the liquid there is no oxygen. Thinking about, it does make sense if it is theoretically a reactive metal. Any traces of oxygen would be removed by the liquid aluminium forming Al₂O₃, leaving most of it un-oxidised. The reaction can then take place, and is very exothermically violent, forming the oxide/hydroxide and the flammable-explosive hydrogen gas. This is an important chemical health and safety issue encountered when dealing with metal extraction and foundry metal processes in industry well away from the relative 'small scale safety' of limited school industrial chemistry! The Thermit reaction: However the true reactivity of aluminium can be spectacularly seen when its grey powder is mixed with brown iron(III) oxide powder. When the mixture is ignited with a magnesium fuse (needed because of the very high activation energy!), it burns very exothermically in a shower of sparks to leave a red hot blob of molten=>solid iron and white aluminium oxide powder. Note the high temperature of the magnesium fuse flame is so high, the oxide layer (to the delight of all pupils) fails to inhibit the displacement reaction! yippee! (see above) aluminium + iron(III) oxide ==> iron + aluminium oxide aluminium + iron(III) oxide ==> aluminium oxide + iron 2Al_(s) + Fe₂O_3(s) ==> Al₂O_3(s) + 2Fe_(s) This is a typical displacement reaction by a more reactive metal displacing a less reactive metal from one of its compounds. Slow reaction with dilute hydrochloric acid to form the colourless soluble salt aluminium chloride and hydrogen gas. (see above) aluminium + hydrochloric acid ==> aluminium chloride + hydrogen 2Al_(s) + 6HCl_(aq) ==> 2AlCl_3(aq) + 3H_2(g) The reaction with dilute sulphuric acid is very slow to form colourless aluminium sulphate and hydrogen. (see above) aluminium + sulphuric acid ==> aluminium sulphate + hydrogen 2Al_(s) + 3H₂SO_4(aq) ==> Al₂(SO₄)_3(aq) + 3H_2(g) If the surface of aluminium is treated with less reactive metal salt, it is still possible to get a displacement reaction. Check this out by leaving a piece of aluminium foil in copper(II) sulphate solution and a patchy pink colour of copper metal slowly appears over many hours/days? Aluminium was first extracted in 1825 by electrolysis of its molten oxide Al₂O₃ (bauxite ore). Extraction details (see alphabetical list at top of this other page
(Carbon C, a non-metal)	Elements higher than carbon i.e. aluminium or more reactive, must be extracted by electrolysis (or displacing it with an even more reactive metal). Metals below it, i.e. zinc or a less reactive, can be extracted by reducing the hot metal oxide with carbon.
zinc Zn At the end of the 1st block-series of Transition Metals	The surface of zinc goes white-yellow when strongly heated in air/oxygen to form zinc oxide (curiously ZnO is white when cold and yellow when hot due to an electron level effect). zinc + oxygen ==> zinc oxide 2Zn_(s) + O_2(g) ==> 2ZnO_(s) No reaction with cold water. When the metal is heated strongly in steam zinc oxide and hydrogen are formed. zinc + water ==> zinc oxide + hydrogen Zn_(s) + H₂O_(g) ==> ZnO_(s) + H_2(g) Quite reactive with dilute hydrochloric acid forming the colourless soluble salt zinc chloride and hydrogen gas. zinc + hydrochloric acid ==> zinc chloride + hydrogen Zn_(s) + 2HCl_(aq) ==> ZnCl_2(aq) + H_2(g) Quite reactive with dilute sulphuric acid forming the colourless soluble salt zinc sulphate and hydrogen gas. zinc + sulphuric acid ==> zinc sulphate + hydrogen Zn_(s) + H₂SO_4(aq) ==> ZnSO_4(aq) + H_2(g) (this reaction is catalysed by adding a trace of copper sulphate solution which form a deposit on the zinc surface) Very little, if any? hydrogen is formed with dilute nitric acid, though zinc nitrate is. This is because another reaction does occur in which the gas nitrogen monoxide (nitrogen(II) oxide, NO) is formed instead of hydrogen. The colourless nitrogen monoxide rapidly combines with oxygen in air to give the dangerous irritating brown gas nitrogen dioxide (nitrogen(IV) oxide, NO₂). (i) zinc + nitric acid ==> zinc nitrate + hydrogen Zn_(s) + 2HNO_3(aq) ==> Zn(NO₃)_2(aq) + H_2(g) which can occur in very dilute nitric acid but has to compete with the reaction ... (ii) zinc + nitric acid ==> zinc nitrate + water + nitrogen(II) oxide [nitric oxide] 3Zn_(s) + 8HNO_3(aq) ==> 3Zn(NO₃)_2(aq) + 4H₂O_(l) + 2NO_(g) and (ii) is rapidly followed rapidly by ... (iii) nitrogen(II) oxide + oxygen ==> nitrogen(IV) oxide 2NO_(g) + O_2(g) ==> 2NO_2(g) [nitric oxide ==> nitrogen dioxide] However with concentrated nitric acid, nitrogen dioxide is formed directly. (iv) zinc + nitric acid ==> zinc nitrate + water + nitrogen(IV) oxide 3Zn_(s) + 4HNO_3(aq) ==> Zn(NO₃)_2(aq) + 2H₂O_(l) + 2NO_2(g) So, whatever concentration of nitric acid is used, you get a solution of zinc nitrate AND nasty brown fumes of nitrogen dioxide. Nitric acid is a strong oxidising agent and it is also NOT a reaction on which to base magnesium's position in the 'metal reactivity series' because of the complications. Adding zinc granules to copper(II) sulphate solution, removes the blue colour of the copper(II) salt, leaving a colourless solution of zinc sulphate and a pinky-brown deposit of copper. zinc + copper sulphate ==> zinc sulphate + copper Zn_(s) + CuSO_4(aq) ==> ZnSO_4(aq) + Cu_(s) This is a typical displacement reaction by a more reactive metal displacing a less reactive metal from one of its compounds. Zinc can be extracted by reducing the hot metal oxide on heating with carbon zinc oxide + carbon ==> zinc + carbon dioxide 2ZnO_(s) + C_(s) ==> 2Zn_(s) + CO_2(g) A zinc coating (galvanising) is used to protect iron from rusting. The more reactive zinc oxidises 1^st. Blocks of zinc attached to steel are used as 'sacrificial corrosion'. Zinc was known and used in India and China before 1500 so it must have been extracted like copper or iron by carbon reduction of the oxide, sulphide or carbonate. Extraction details (see alphabetical list at top of this other page
iron Fe In the 1st block-series of Transition Metals	The surface of iron goes dark grey-black when strongly heated in air/oxygen to form a tri-iron tetroxide. When steel wool is heated in a bunsen flame it burns with a shower of sparks - large surface area - increased rate of reaction - so even moderately reactive iron has its moments! iron + oxygen ==> iron oxide [iron tetroxide, diiron(III)iron(II) oxide] 3Fe_(s) + 2O_2(g) ==> Fe₃O_4(s) No reaction with cold water (rusting is a joint reaction with oxygen). When the metal is heated in steam an iron oxide (unusual formula) and hydrogen are formed. This oxide is 'technically' diiron(III)iron(II) oxide but its sometimes called 'tri-iron tetroxide'. iron + water (steam) ==> iron tetroxide + hydrogen 3Fe_(s) + 4H₂O_(g) ==> Fe₃O_4(s) + 4H_2(g) Moderate reaction with dilute hydrochloric acid forming the soluble pale green salt iron(II) chloride and hydrogen gas. iron + hydrochloric acid ==> iron(II) chloride + hydrogen Fe_(s) + 2HCl_(aq) ==> FeCl_2(aq) + H_2(g) It does not form iron(III) chloride, FeCl3, in this reaction, but it does form this other iron chloride compound when iron is heated in a stream of chlorine gas (see salt preparation by direct synthesis note). Slow reaction with dilute sulphuric acid forming the soluble pale green salt iron(II) sulphate and hydrogen gas. iron + sulphuric acid ==> iron(II) sulphate + hydrogen Fe_(s) + H₂SO_4(aq) ==> FeSO_4(aq) + H_2(g) Iron can be extracted by reducing the hot metal oxide on heating with carbon monoxide formed from carbon in the blast furnace e.g. iron(III) oxide + carbon monoxide ==> iron + carbon dioxide Fe₂O_3(s) + 3CO_(g) ==> 2Fe_(s) + 3CO_2(g) iron tetroxide + carbon monoxide ==> iron + carbon dioxide Fe₃O_4(s) + 4CO_(g) ==> 3Fe_(s) + 4CO_2(g) For the past 2500 years. iron has been extracted from pre-historic times using charcoal (C). Known in Anglo-Saxon as 'iron' and in Roman times in Latin as 'ferrum' hence the Fe symbol! Extraction details - see alphabetical list at top of this other page
tin Sn A Group 4 metal	Slow reaction when heated in air to form white tin(IV) oxide or tin dioxide tin + oxygen ==> tin oxide [tin dioxide, tin(IV oxide] Sn_(s) + O_2(g) ==> SnO_2(s) No reaction with cold water or when heated in steam. Very slow reaction with dilute hydrochloric acid forming the slightly soluble tin(II) chloride and hydrogen gas. tin + hydrochloric acid ==> tin(II) chloride + hydrogen Sn_(s) + 2HCl_(aq) ==> SnCl_2(aq) + H_2(g) Very slow reaction with dilute sulphuric acid forming the colourless slightly soluble tin(II) sulphate and hydrogen gas. tin + sulphuric acid ==> tin(II) sulphate + hydrogen Sn_(s) + H₂SO_4(aq) ==> SnSO_4(aq) + H_2(g) Tin can be extracted from its oxide by heating with carbon. Tin has been known from pre-historic times. Known in Anglo-Saxon as 'tin' and in Latin - 'stannum' hence the symbol Sn!
lead Pb A Group 4 metal	Slow reaction when heated in air to form red/yellow lead(II) oxide and tri-lead tetroxide lead + oxygen ==> lead(II) oxide [lead monoxide] 2Pb_(s) + O_2(g) ==> 2PbO_(s) and 3Pb_(s) + 2O_2(g) ==> Pb₃O_4(s) No reaction with cold water or when heated in steam. Very slow and effectively no reaction with dilute hydrochloric acid or dilute sulphuric acid. Lead can be extracted from its oxide by heating with carbon. Probably used from pre-historic times and known in Anglo-Saxon as 'lead' and in Latin 'plumbum' hence the symbol Pb!
Hydrogen H non-metal	Non of the metals below hydrogen can react with acids to form hydrogen gas. They are least easily corroded metals and partly accounts for their value and uses in jewellery, electrical contacts etc.
copper Cu In the 1st block-series of Transition Metals	Surface blackens when strongly heated in air/oxygen to form copper(II) oxide. copper + oxygen ==> copper oxide [copper(II) oxide] 2Cu_(s) + O_2(g) ==> 2CuO_(s) No reaction with cold water or when heated in steam. No reaction with dilute hydrochloric acid or dilute sulphuric acid. Copper can be extracted by reducing the hot black metal oxide on heating with carbon copper(II) oxide + carbon ==> copper + carbon dioxide 2CuO_(s) + C_(s) ==> 2Cu_(s) + CO_2(g) Extraction and purification details (see alphabetical list at top) Although copper doesn't readily react with dilute hydrochloric acid and dilute sulphuric acid (low in reactivity series), if heated with nasty oily concentrated sulphuric acid you make nasty pungent irritating sulphur dioxide gas and white anhydrous copper(II) sulphate, but this is NOT a reaction on which to base its place in the metal reactivity series and hydrogen gas isn't produced. copper + sulphuric acid ==> copper(II) sulphate + sulphur dioxide + water Cu_(s) + 2H₂SO_4(l) ==> CuSO_4(s) + SO_2(g) + H₂O_(l) Hydrogen is NOT formed with dilute nitric acid, though copper(II) nitrate is. This is because another reaction does occur in which the gas nitrogen monoxide (nitrogen(II) oxide, NO) is formed instead of hydrogen. The colourless nitrogen monoxide rapidly combines with oxygen in air to give the dangerous irritating brown gas nitrogen dioxide (nitrogen(IV) oxide, NO₂). (i) copper + nitric acid ==> copper(II) nitrate + water + nitrogen(II) oxide [nitric oxide] 3Cu_(s) + 8HNO_3(aq) ==> 3Cu(NO₃)_2(aq) + 4H₂O_(l) + 2NO_(g) and (i) is rapidly followed rapidly by ... (ii) nitrogen(II) oxide + oxygen ==> nitrogen(IV) oxide [nitric oxide + oxygen ==> nitrogen dioxide] 2NO_(g) + O_2(g) ==> 2NO_2(g) However with concentrated nitric acid, nitrogen(IV) oxide [nitrogen dioxide] is formed directly. (iii) copper + nitric acid ==> copper(II) nitrate + water + nitrogen(IV) oxide Cu_(s) + 4HNO_3(aq) ==> Cu(NO₃)_2(aq) + 2H₂O_(l) + 2NO_2(g) So, whatever concentration of nitric acid is used, you get a blue solution of copper(II) nitrate AND nasty brown fumes of nitrogen dioxide. Nitric acid is a strong oxidising agent and it is also NOT a reaction on which to base magnesium's position in the 'metal reactivity series' because of the complications. The elemental metal can be found 'native' and was probably first used over 6000 years ago in Turkey by literally beating it out of rocks and into shape (malleable at room temperature!) - no high temperature technology used or available. It has been extracted by carbon reduction of a copper mineral for at least 3000 years. Latin 'cuprum' meaning Cyprus?, anyway that's why its symbol is Cu! Extraction and purification details - see alphabetical list at top of this other age.
silver Ag a transition metal (2nd series)	No reaction when heated in air. No reaction with cold water or when heated in steam. No reaction with dilute hydrochloric acid or dilute sulphuric acid. Silver can be extracted by BUT can be found 'native' as the element because it is so unreactive. It has been used from pre-historic times in jewellery for 4000 years at least. In Anglo-Saxon it was 'siolfur' meaning 'silver in nature' and in Latin 'Argentum' hence its symbol Ag.
gold Au a transition metal (3rd series)	No reaction when heated in air No reaction with cold water or when heated in steam. No reaction with dilute hydrochloric acid or dilute sulphuric acid. Gold can be readily extracted from its ores easily by reduction BUT it is usually found 'native' as the element because it is so unreactive and has been used from pre-historic times in jewellery for at least 6000 years. Known in Anglo-Saxon as 'gold'. Gold is rather a soft metal and is 'hardened' by alloying with other metals - pure gold is 24 carat - 22, 18, 15, 12 and 9 carat gold are legalised, meaning it has that carat number/24 as parts of gold as a measure of its purity and value!
platinum Pt a transition metal (3rd series)	No reaction when heated in air. No reaction with cold water or when heated in steam. No reaction with dilute hydrochloric acid or dilute sulphuric acid. It seems ironic that despite its apparent lack of 'reactivity' it is a very potent catalyst e.g. catalytic converter in cars. Spanish 'platina' meant 'silvery in nature'. Like gold, it is a very rare metal but was known by pre-Columbian South American Indians and brought to Europe in about 1750. It is used in expensive jewellery, laboratory ware (e.g. inert crucible container) and catalytic converters in car exhausts.

METAL CORROSION and the RUSTING of IRON

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. You can do simple experiments to show that BOTH oxygen and water are needed. Put an iron nail into (1) boiled water in a sealed tube; (2) a tube of air and a drying agent; (3) an open test tube with water. Rusting appears overnight with (3) only.
Rusting is speeded up in the presence of salt or acid solutions because of an increased concentration of ions. Corrosion is a redox process involving redox electron transfer and ion movement. The rusting metal behaves like a simple cell and more ions enable the current, and hence the electron transfer, to occur more readily.
Rusting is overall:
- 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 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!).
- The reaction proceeds via (i) iron(II) hydroxide Fe(OH)₂ which is (ii) oxidised further to the hydrated Fe₂O₃, or if very soggy, it amounts to the formation of iron(III) hydroxide!
  - The reactions can be summarised in terms of hydroxide formation e.g.
  - (i) iron + water + oxygen ==> iron(II) hydroxide
  - 2Fe_(s) + 2H₂O_(l) + O_{2 (g)} ==> 2Fe(OH)_{2 (s)}
  - (ii) iron(II) hydroxide + water + oxygen ==> iron(II) hydroxide
  - 4Fe(OH)_{2 (s)} + 2H₂O_(l) + O_{2 (g)} ==> 4Fe(OH)_{3 (s)}
- Rusting is an oxidation because it involves iron gaining oxygen (Fe ==> Fe₂O₃) or atoms of iron losing electrons (Fe - 3e^- ==> Fe³⁺).
- See more examples of oxidation and reduction below.
The rusting of iron is a major problem in its use as a structural material.
- Preventing rusting adds cost to manufacturing things.
- Corroded components or structures weakened, adding further costs in rust treatment or replacement.
Iron and steel (alloy of iron) are most easily protected by paint which provides a barrier between the metal and air/water. Moving parts on machines can be protected by a water repellent oil or grease layer.

reactivity An experiment to investigate sacrificial corrosion

This 'rusting' corrosion can be prevented by connecting iron to a more reactive metal (e.g. zinc or magnesium). This is referred to as sacrificial protection or sacrificial corrosion, because the more reactive protecting metal is preferentially oxidised away, leaving the protected metal intact. The picture illustrates what might be seen after a few days.* Iron or steel can also be protected by mixing in other metals (e.g. chromium) to make non-rusting alloys called stainless steel. The chromium, like aluminium, forms a protective oxide layer.
* Theoretically, any iron ions formed by oxidation would be reduced by electrons from the oxidation of the more reactive 'sacrificed' metal.
Coating iron or steel with a thin zinc layer is called 'galvanising'. The layer is produced by electrolytic deposition by making the iron/steel the negative cathode or by dipping the iron/steel object in molten zinc (more details on the Extra Industrial Chemistry page). The zinc preferentially corrodes or oxidises to form a zinc oxide layer that doesn't flake off like iron oxide rust does. Also, if the surface is scratched, the exposed zinc again corrodes before the iron and continues to protect it.
Steel tin cans are protected by relatively unreacted tin and works well as long as the thin tin layer is complete and acts as an inert barrier between the steel and oxygen (air)/water HOWEVER, if a less reactive metal is connected to the iron, it then the iron rusts preferentially (try scratching a 'tin' can and leave out in the rain and note the corrosion by the scratch!)

Aluminium does not oxidise (corrode) as quickly as its reactivity would suggest. Once a thin oxide layer of Al₂O₃ has formed on the surface, it forms a barrier to oxygen and water and so prevents further corrosion of the aluminium.
Aluminium is a useful structural metal. It can be made harder, stronger and stiffer by mixing it with small amounts of other metals (e.g. magnesium) to make alloys.

Copper and Lead are both used in roofing situations because neither is very reactive and the compounds formed do not flake away as easily as rust does from iron. Lead corrodes to a white lead oxide or carbonate and copper corrodes to form a basic green carbonate (combination of the hydroxide Cu(OH)₂ and carbonate CuCO₃ e.g. seen as green roof on buildings).
Both metals have been used for piping but these days lead is considered too toxic and copper is usually used as the stronger, but equally unreactive alloy with zinc, brass. Now of course, most piping is flowing in the plastic direction which doesn't corrode at all!

The Group 1 Alkali Metals rapidly corrode in air and need to be stored under oil.
Apart from their structural weakness they would hardly used for any outside purpose!

OXIDATION and REDUCTION - REDOX REACTIONS
OXIDATION - definition and examples	REDUCTION - definition and examples
(a) The gain or addition of oxygen by an atom, molecule or ion e.g. ... (1) S ==> SO₂ [burning sulphur - oxidised] (2) CH₄ ==> CO₂ + H₂O [burning methane to water and carbon dioxide, C and H gain O] (3) NO ==> NO₂ [nitrogen monoxide oxidised to nitrogen dioxide] (4) SO₃ ==> SO₃ [oxidising the sulphur dioxide to sulphur trioxide in the Contact Process for making sulphuric acid]	(b) The loss or removal of oxygen from a compound etc. e.g. ... (1) CuO ==> Cu [loss of oxygen from copper(II) oxide to form copper atoms] (2) Fe₂O₃ ==> Fe [iron(III) oxide reduced to iron in blast furnace] (3) NO ==> N₂ [nitrogen monoxide reduced to nitrogen, catalytic converter in car exhaust] (4) SO₃ ==> SO₂ [sulphur trioxide reduced to sulphur dioxide]
(c) The loss or removal of electrons from an atom, ion or molecule e.g. (1) Fe ==> Fe²⁺ + 2e^- [iron atom loses 2 electrons to form the iron(II) ion, start of rusting chemistry] (2) Fe²⁺ ==> Fe³⁺ + e^- [the iron(II) ion loses 1 electron to form the iron(III) ion] (3) 2Cl^- ==> Cl₂ + 2e^- [the loss of electrons by chloride ions to form chlorine molecules]	(d) The gain or addition of electrons by an atom, ion or molecule e.g. ... (1) Cu²⁺ + 2e^- ==> Cu [the copper(II) ion gains 2 electrons to form neutral copper atoms, electroplating or displacement reaction) (2) Fe³⁺ + e^-==> Fe²⁺ [the iron(III) ion gains an electron and is reduced to the iron(II) ion] (3) 2H⁺ + 2e^- ==> H₂ [hydrogen ions gain electrons to form neutral hydrogen molecules, electrolysis of acids or metal-acid reaction]
(e) An oxidising agent is the species that gives the oxygen or removes the electrons	(f) A reducing agent is the species that removes the oxygen or acts as the electron donor
REDOX REACTIONS - in a reaction overall, oxidation and reduction must go together
(g) Redox reaction analysis based on the oxygen definitions
(1) copper(II) oxide + hydrogen ==> copper + water CuO_(s) + H_2(g) ==> Cu_(s) + H₂O_(g) copper oxide reduced to copper, hydrogen is oxidised to water hydrogen is the reducing agent (removes O from CuO) copper oxide is the oxidising agent (donates O to hydrogen) (2) iron(III) oxide + carbon monoxide ==> iron + carbon dioxide Fe₂O_3(s) + 3CO_(g) ==> 2Fe_(l) + 3CO_2(g) the iron(III) oxide is reduced to iron, the carbon monoxide is oxidised to carbon dioxide CO is the reducing agent (O remover from Fe₂O₃) the Fe₂O₃ is the oxidising agent (O donator to CO)] (3) nitrogen monoxide + carbon monoxide ==> nitrogen + carbon dioxide 2NO_(g) + 2CO_(g) ==> N_2(g) + 2CO_2(g) nitrogen monoxide is reduced to nitrogen carbon monoxide is oxidised to carbon dioxide CO is the reducing agent and NO is the oxidising agent (4) iron(III) oxide + aluminium ==> aluminium oxide + iron (the Thermit reaction) Fe₂O_3(s) + 2Al_(s) ==> Al₂O_3(s) + 2Fe_(s) iron(III) oxide is reduced and is the oxidising agent aluminium is oxidised and is the reducing agent
(h) Redox reaction analysis based on the electron definitions
(1) magnesium + iron(II) sulphate ==> magnesium sulphate + iron Mg_(s) + FeSO_4(aq) ==> MgSO_4(aq) + Fe_(s) this is the 'ordinary molecular' equation for a typical metal displacement reaction, but this does not really show what happens in terms of atoms, ions and electrons, so we use ionic equations like the one shown below. The sulphate ion SO₄^2-_(aq) is called a spectator ion, because it doesn't change in the reaction and can be omitted from the ionic equation. No electrons show up in the full equations because electrons lost by x = electrons gained by y!! magnesium + iron(II) ion ==> magnesium ion + iron Mg_(s) + Fe²⁺_(aq) ==> Mg²⁺_(aq) + Fe_(s) the magnesium atom loses 2 electrons (oxidation) to form the magnesium ion, the iron(II) ion gains 2 electrons (reduced) to form iron atoms. Mg is the reducing agent (electron donor) and the Fe²⁺ is the oxidising agent (electron remover or acceptor) Displacement reactions involving metals and metal ions are electron transfer reactions. (2) zinc + hydrochloric acid ==> zinc chloride + hydrogen Zn_(s) + 2HCl_(aq) ==> ZnCl_2(aq) + H_2(g) the chloride ion Cl^- is the spectator ion zinc + hydrogen ion ==> zinc ion + hydrogen Zn_(s) + 2H⁺_(aq) ==> Zn²⁺_(aq) + H_2(g) Zinc atoms are oxidised to zinc ions by electron loss, so zinc is the reducing agent (electron donor) hydrogen ions are the oxidising agent (gaining the electrons) and are reduced to form hydrogen molecules (3) copper + silver nitrate ==> silver + copper(II) nitrate Cu_(s) + 2AgNO_3(aq) ==> 2Ag + Cu(NO₃)_2(aq) the nitrate ion NO₃^- is the spectator ion copper + silver ion ==> silver + copper(II) ion Cu_(s) + 2Ag⁺_(aq) ==> 2Ag_(s) + Cu²⁺_(aq) copper atoms are oxidised by the silver ion by electron loss electrons are transferred from the copper atoms to the silver ions, which are reduced the silver ions are the oxidising agent and the copper atoms are the reducing agent (4) iron(II) chloride + chlorine ==> iron(III) chloride (5) halogen (more reactive) + halide salt (of less reactive halogen) ==> halide salt (of more reactive halogen) + halogen (less reactive) X_2(aq) + 2KY_(aq) ==> 2KX_(aq) + Y_2(aq) X_2(aq) + 2Y^-_(aq) ==> 2X^-_(aq) + Y_2(aq) where halogen X is more reactive than halogen Y, F > Cl > Br > I X is the oxidising agent (electron acceptor) KY is the reducing agent (electron donor) See GCSE Group 7 The Halogens - displacement reaction notes (6) Electrode reactions in electrolysis are electron transfer redox changes at the negative cathode positive ions are attracted: metal ions are reduced to the metal by electron gain: Mⁿ⁺ + ne^- ==> M n = the numerical charge of the ion and the number of electrons transferred or 2H⁺_(aq) + 2e^- ==> H_2(g) (for the discharge of hydrogen) at the positive anode negative ions are attracted: negative non-metal ions are oxidised by electron loss e.g. for oxide ions: 2O^2- - 4e^- ==> O₂ or 2O^2- ==> O₂ + 4e^- for hydroxide ion: 4OH^- - 4e^- ==> O₂ + 2H₂O or 4OH^- ==> O₂ + 2H₂O + 4e^- for halide ions (X = F, Cl, Br, I): 2X^- - 2e^- ==> X₂ or 2X^- ==> X₂ + 2e^-
Miscellaneous Extra Redox Notes Redox changes can often be observed as significant colour changes e.g. iron + copper(II) sulphate ==> iron(II) sulphate + copper Fe_(s) + CuSO_4(aq) ==> FeSO_4(aq) + Cu_(s) iron + copper(II) ion ==> iron(II) ion + copper Fe_(s) + Cu²⁺_(aq) ==> Fe²⁺_(aq) + Cu_(s) Sulphate, SO₄^2-_(aq), is colourless BUT a blue to pale green colour change is observed in the solution as the blue copper(II) ion is replaced by the pale green iron(II) ion as well as the pink-dark precipitate of copper metal. Potassium manganate(VII) is a powerful oxidising agent and an intense purple colour in water due to the MnO₄^- ion. In acidified solution it changes to an almost colourless* manganese(II) ion, Mn²⁺ when it oxidises something (* which actually is a very pale pink transition metal ion). Potassium dichromate(VI) is another strong oxidising agent and is orange due to the dichromate(VI) ion, Cr₂O₇^2- ion. When it oxidises something it changes to the green chromium(III) ion, Cr³⁺. Potassium iodide is a colourless salt dissolving in water to form a colourless solution. If it is oxidised e.g. with chlorine a yellow==>orange==>brown colour develops as iodine is formed from the colourless iodide ion. The use of Roman Numerals in names: This indicates what is called the oxidation state of an atom in a molecule or ion. It is easy to follow for simple metal ions because it equals the charge on the ion e.g. the oxidation state of copper in the copper(II) ion is referred to as +2 the more electrons removed from the atom or ion by oxidation, the higher its oxidation state e.g. Fe²⁺ - e^- ==> Fe³⁺, gives iron the oxidation state of +3 in the iron(III) ion (via a suitable oxidising agent). but for more complex ions things are not so simple. in manganate(VII) ion, the Mn is in the +7 oxidation state in dichromate(VI) ion, the Cr is in the +6 oxidation state This topic is dealt with more thoroughly at AS-A2 advanced level chemistry. ks4 science examinations e-gcse-igcse chemistry revision * ks4 science examinations-gcse-igcse chemistry revision * ks4 science examinations-gcse-igcse chemistry revision * ks4 science examinations-gcse-igcse chemistry revision * ks4 science examinations-gcse-igcse chemistry revision * ks4 science examinations-gcse-igcse chemistry revision * SITE PURPOSE EDUCATION - online learning or 'self-private-tuition' using revision notes, quizzes, practice tests involving GCSE Science CHEMISTRY in the areas of REVISING only the CHEMISTRY-Earth Science-Radioactivity at Doc Brown's Chemistry Clinic via HOMEPAGE in secondary school/schools, 6th form college/colleges, academy/academies or home self-study. Hopefully it will encourage interest and understanding of Chemistry, Earth Science and Radioactivity in any country of the world, though the site is written entirely in English. The website is designed to help and unofficially support students/teachers revise-learn/teach the chemistry for modular or co-ordinated examination science courses from UK QCA based AQA, OCR (Oxford and Cambridge) Twenty First (21st) Century and Gateway Science, Edexcel 360Science , Nuffield, Salters, Cambridge International (CIE), London International, WJEC, CCEA exams etc. Also, national award assessments-examinations for GCSE-IGCSE-KS4-O level-BTEC-NVQ applied, additional and chemistry national science courses. Also covers, mainly via quizzes the UK National KS3 SATs Science-biology/chemistry/physics (SAT revision levels 3-5 or 5-7) and covers much of the revising, learning and teaching chemistry examinations for the national curriculum for secondary schools and colleges. The site does not support the content of England, Wales or Northern Ireland primary science KS1 or KS2. The notes should also provide some background theory for a coursework assignment or project. BUT please note that my on-line revision notes and quizzes are no substitute for good classroom teaching-lecturing and thorough studying of your own notes and textbooks, practicing past papers and a copy of the syllabus which are readily downloaded from the examination board sites, but I hope here and there they will lend a tutoring hand on some topic, unit, module etc. For final revision you have to be intellectually honest about what you don't know or follow, YOU have to take the stuff to pieces, analyse what you do/do not understand and reconstruct it so it all makes sense in the end. There is no other way, there are no magic secrets on how to revise and learn, its mainly down to hard work and just good old fashioned study and employing teach-yourself strategies without the need for extra tutors and tutoring lessons. I also think there is too much hit and miss revision using past papers (which I do NOT supply) and not enough systematic revision. I also hope it will help teachers in planning lessons and developing schemes of work for science-chemistry. There are no lesson plans on the site but there are plenty of quizzes to incorporate into classroom activities whether photocopied or on electronic whiteboard projector for use as self-tuition-assessment purposes and a variety of teaching and learning styles and the images may be used in Microsoft Word documents and powerpoint projections. The site seems to be used by a large number of home study tutors, particularly the revision notes. An individual tutor may print out the notes for science-chemistry learning teaching-tuition purposes and for background material for assignments and projects. I have no interest or time in producing WORD.doc or xxxx.pdf files of the notes at the moment. Neither have I time to write up many practical laboratory experiments ('lab'-'labs') at the moment, but the notes contain lots of background information of chemical reactions in terms of observations-balanced equations-reactants-products-theory etc. I also find it difficult to recommend specific exam websites or syllabus textbooks, it depends exactly on what you need, what you have time for, and there are so many of them to choose from and I do not supply past examination papers for classes. The