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GCSE Chemistry Notes: The Reactivity Series of Metals

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Reactivity series of metals   Reactivity series of metals

Reactivity series of metalsReactivity series of metalsDoc Brown's Chemistry KS4 Science GCSE chemistry/IGCSE Chemistry Revision Notes

1. The Metal Reactivity Series - Metal Reaction Notes - equations and explanations

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Sub-index for this page

(a) Introduction to the reactivity series of metals

(b) A summary diagram of important ideas to do with the reactivity series of metals

(c) Reactivity of the metals and their reactions

aluminium .. caesium .. calcium .. copper .. francium .. gold .. iron .. lead .. lithium

magnesium .. platinum .. potassium .. rubidium .. silver .. sodium .. tin .. zinc

(d) Other associated pages and quizzes

SEE ALSO

2. RUSTING & Introducing REDOX reactions

3. Metal Reactivity Series Experiments-Observations-theory and ...

Metal Extraction Fe, Cu, Al etc.  *  Transition Metals

Other notes on using metals eg Al & Ti  *  Metal Structure - bonding

How to use the reactivity series to predict the voltage of a simple cell

EQUATION NOTE: The equations are often written three times: (i) word equation, (ii) balanced symbol equation without state symbols, and, (iii) with the state symbols (g), (l), (s) or (aq) to give the complete balanced symbol equation.

Keywords and phrases: The ideas behind the 'Reactivity Series of Metals' is introduced and what happens to a metal atom when it reacts. The experimental evidence for establishing the reactivity order for metals is described in terms of metal displacement reactions and the reactions of metals with oxygen (i.e. heating or burning in air), reaction with cold water and hydrochloric acid, sulfuric/sulfuric acid and nitric acid, but the theory and chemical equation details are on this page. The implications of a metal's reactivity and method of extraction are also mentioned, but the methods of extraction are on separate web pages. These revision notes on the reactivity series of metals, theory of reactivity order and balanced symbol equations should prove useful for the new AQA, Edexcel and OCR GCSE (9–1) chemistry science courses.

(a) Introduction to the REACTIVITY SERIES OF METALS:

Reactivity series of metals Reactivity series of metals

potassium > sodium > calcium > magnesium > aluminium > zinc > iron > tin > lead > copper > silver > gold > platinum

A short introduction: When metals react with other substances like water, oxygen or acids, the metal atoms form positive ions (characteristic chemical behaviour of metals.

The chemical reactivity of a metal (e.g. how fast it reacts) is related to its tendency to form positive ions (cations). From quite simple experiments, metals can be arranged in order of their reactivity in what we call a reactivity series.

In order of decreasing reactivity, the metals potassium, sodium, lithium, calcium, magnesium, zinc, iron and copper can be put in order of their reactivity from their reactions with water and dilute acids.

The non-metals hydrogen and carbon are often included in the reactivity series, and this quite important when considering the method by which a metal can be extracted from its ore.

You should also note that a more reactive metal will displace a less reactive metal from one of its compounds.

reactivityFor 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 on this page are in order of decreasing position in the reactivity series of metals)

  • The reactivity series of metals is an important concept in chemistry and has implications for corrosion chemistry and its prevention, metal extraction, storage of reactive metals.
  • The higher the metal in the series, the more reactive it is and you usually observe a more vigorous - faster and more exothermic (heat releasing) reaction with oxygen, water or an acid.
  • At a more theoretical level, the more reactive a metal, the greater tendency it has to form a positive ion (cation) by losing electrons in the context of a chemical reaction
    • e.g. for sodium Na ==> Na+ or iron Fe ==> Fe2+

    • The more reactive a metal, the more easily it loses electrons to form a positive ion, an electron loss is a process called oxidation.

      • e.g. sodium loses its outer electron much more easily than iron loses two of its outer electrons, consequently, sodium is a much more reactive metal than iron.

      • Positive ions like Na+,  Mg2+ or Al3+ are known as cations (they migrate to the negative cathode in electrolysis).

    • A largish atom like potassium is very reactive because the single outer electron is readily lost in reacting with water of acids - the potassium ion formation and change in electronic structure is shown below.

    • (c) doc b ===> Reactivity series of metals  +  e–

    • K   or   [2.8.8.1]  ==> K+   or   [2.8.8]+   +   e–

    • A smallish atom like magnesium is not quite as reactive because the two outer electrons are not as easily lost in reacting with water or acids - the magnesium ion formation and change in electronic structure is shown below.

    • Reactivity series of metals ===> Reactivity series of metals  +  2e–

    • Mg   or   [2.8.2]  ==> Mg2+   or   [2.8]2+   +   2e–

    • This argument applies whatever the metal reacts with e.g. oxygen, water, acids or salt solutions of other metals.

    • On the alkali metals page is more detailed argument focussing on the group reactivity trend.

    • The more reactive a metal, the more easily it is oxidised - the more easily the metal atoms lose electrons to form a positive ion.

      • This means the more reactive a metal, the more likely it is to be corroded - means the metal is more susceptible to be attacked by a combination of oxygen (from air) and water.

      • e.g. two extremes - sodium exposed to air at room temperature is rapidly coated in sodium oxide whereas gold or platinum do not corrode at all.

      • For more details see section 2. RUSTING & Introducing REDOX reactions

  • The reactivity trend implies that the reverse reaction becomes more difficult the more reactive a metal.
    • This has implications for a metal is extracted from its ore.
    • i.e. the more reactive a metal, the more difficult it is to extract the metal from its ore, in other words its more difficult to remove the other element the metal is combined with e.g. oxygen or sulfur.
    • It also means that the more reactive a metal the more susceptible it is to corrosion from 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). Details of the reactions are given on this page and details of some experiments and possible observations are given in section 3. (c) doc b Metal Reactivity Series Experiments-Observations (separate page)
  • Most metals react with oxygen to form an oxide
    • This is an oxidation reaction because the metal gains oxygen (see redox theory, redox is shorthand for ;oxidation and reduction'). Technically any oxidation is accompanied by a reduction - but not concerned with that here.
    • Very reactive metals like sodium rapidly tarnish in air with the formation of an oxide layer.
    • Magnesium readily burns in air once magnesium ribbon is strongly heated in a bunsen flame.
      • 2Mg  +  O2  ===>  2MgO
    • Less reactive metals don't burn as brightly and some like iron or tin won't burn but will be oxidised to form the oxide on heating in air.
      • In reactions to extract metal from oxide ores, the ore is some agent that removes the oxygen, so the oxide is reduced, oxygen loss - the opposite of the oxidation reactions described above.
  • Displacement reactions - one of the best ways of establishing the reactivity series of metals
    • A metal in the series, can displace any metal below it in the series, from the less reactive metal's solid oxide, or chloride, sulfate or other soluble compound from solution.
    • 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:
      • eg    magnesium + copper oxide ==> magnesium oxide + copper
        • Mg + CuO ==> MgO + Cu
        • Mg(s) + CuO(s) ==> MgO + Cu(s)
        • Mg oxidised by oxygen gain, copper oxide reduced, by oxygen loss.
      • The more reactive magnesium displaces the less reactive copper as it does in the 2nd example below.
      • See also the Thermit reaction.
    • or adding a metal to a salt solution of another metal e.g. adding magnesium to blue copper(II) sulfate solution, the blue colour fades as colourless magnesium sulfate is formed and brown bits of copper metal form a precipitate:
      • eg      magnesium + copper sulfate ==> magnesium sulfate + copper
        • Mg + CuSO4 ==> MgSO4 + Cu
          • Mg(s) + CuSO4(s) ==> MgSO4(s)  + Cu(s)   ('normal' equation with state symbols)
          • Mg(s) + Cu2+(aq) ==> Mg2+(aq) + Cu(s)    (the ionic equation, redox equation)
          • Magnesium atoms (Mg) are oxidised to magnesium ions (Mg2+) by electron loss, and the copper ions (Cu2+) are reduced to copper atoms by electron gain.
      • concept diagram magnesium oxidised displacing copper ion reduced oxidation = electron loss reduction = electron gain
      • The electron transfer redox theory behind displacement reactions is explained later on another page.
    • If no reaction happens, then it means the added metal is less reactive than the metal in the oxide or sulfate etc.
  • You can also determine the reactivity series of metal activity by reacting the metals with water or acids.
  • Some general word equations where the metal does react with water or acid:
    • (a) metal + cold water ==> metal hydroxide + hydrogen (metals above aluminium)
      • eg sodium + water ==> sodium hydroxide + hydrogen
        • 2Na + 2H2O ==> 2NaOH + H2
          • 2Na(s) + 2H2O(l) ==> 2NaOH(aq) + H2(g)
    • (b) heated metal + steam ==> metal oxide + hydrogen (for metals above tin?)
      • eg magnesium + water ==> magnesium oxide + hydrogen
        • Mg + H2O ==> MgO + H2
          • Mg(s) + H2O(g) ==> MgO(s) + H2(g)
    • (c) metal + acid ==> metal salt + hydrogen
      • (for metals above hydrogen)
      • eg magnesium + hydrochloric acid ==> magnesium chloride + hydrogen
        • Mg + 2HCl ==> MgCl2 + H2
          • Mg(s) + 2HCl(aq) ==> MgCl2(aq) + H2(g)
        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 sulfuric acid makes a metal sulfate 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 is formed rather than hydrogen (not usually studied at GCSE level these days, but often in IGCSE courses).
        • 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 on another page.
  • Within the general Reactivity or Activity Series of Metals there are some Periodic Table Trends too …
    • 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 > Be
    • 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 > (Zn) > Fe > Cu etc. (on Period 4)
  • The reactivity of a metal has an important bearing on the (c) doc b 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.
      • In sense this means metals above carbon in the reactivity series cannot be 'displaced' from their compounds by carbon.
    • Metals below carbon in the series can be extracted by heating the oxide with carbon or carbon monoxide.
      • In other words, metals above carbon in the reactivity series cannot be displaced by carbon OR you can say carbon is not reactive enough to displace metals above it in the reactivity series.
    • Metals below hydrogen will not usually displace hydrogen from acids and can be extracted by heating the oxide in hydrogen, though is rarely done e.g.
      • The method costs more. Carbon from coke (made from coal) is cheaper and safety reasons - hydrogen is a very explosive gas in air.
      • Again, you can think of metals above hydrogen in the reactivity series as being reactive enough to displace hydrogen from acids in aqueous solution OR you can say hydrogen is not reactive enough to displace the metals above in the reactivity series.
  • (c) doc b Notes on the corrosion of metals and the prevention of iron rusting are dealt with on another page and the theory of (c) doc b OXIDATION and REDUCTION and their application to REDOX reactions is also dealt with on the same 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.
    • 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 page)

 

(b) A summary diagram of important ideas to do with the reactivity series of metals!

Reactivity series of metals - method of metal extraction - relative ease of oxidation reaction with acids

METAL in DECREASING REACTIVITY ORDER

(and where in the Periodic Table)

(c) Reactivity of the metals and their reactions

The compounds formed in the reactions are white insoluble solids, (s), or soluble colourless solutions, (aq), unless otherwise stated eg in the case of blue copper compounds or pale green iron compounds. 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

reactivity

The Reactivity Series of Metals

(reactivity of francium and compared with the non-metals carbon and hydrogen)

 See RADIOACTIVITY NOTES - symbol for a readioactive substance 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! Its chemistry is identical to cesium described below [just change Cs (below) for a Fr in any of the formulae or equations, because their chemistry is identical].

Francium is the most reactive metal known because it most easily loses its outer electron to form a positive ion (Fr+). It behaves like any other alkali metal ...

francium + water ==> francium hydroxide + hydrogen

2Fr(s) + 2H2O(l) ==> 2FrOH(aq) + H2(g)

Francium is much too dangerous a metal to add to acids because of its high reactivity apart from spreading the danger of radioactivity around the laboratory!

caesium Cs

Reactivity series of metals

Group 1 Alkali Metal

reactivity

The Reactivity Series of Metals

(reactivity of cesium and compared with the non-metals carbon and hydrogen)

  • 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.

  • Caesium burns vigorously with a blue flame when heated in air/oxygen to form the white powder caesium oxide.

    • caesium + oxygen ==> caesium oxide

    • 4Cs(s) + O2(g) ==> 2Cs2O(s)

      • Cesium is oxidised, oxygen gain, oxidation reaction.

      • also forms caesium peroxide, Cs2O2 and caesium superoxide, CsO2

      • When caesium oxide is dissolved in water it forms caesium hydroxide and the solution turns universal indicator solution or litmus paper blue-purple. Using pH indicator paper or a pH meter you find the alkaline solution has a pH of 13-14.

  • 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) + H2(g)

    • Cesium is much too dangerous a metal to add to acids because of its high reactivity.

  • Caesium was first  extracted in 1860  by electrolysis of the molten chloride CsCl.

rubidium Rb

Reactivity series of metals

Group 1 Alkali Metal

reactivity

The Reactivity Series of Metals

(reactivity of rubidium and compared with the non-metals carbon and hydrogen)

  • 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.

  • Rubidium burns vigorously with a red flame when heated in air/oxygen to form the white powder rubidium oxide.

    • rubidium + oxygen ==> rubidium oxide
    • 4Rb(s) + O2(g) ==> 2Rb2O(s)

      • Rubidium is oxidised, oxygen gain, oxidation reaction.

      • also forms rubidium peroxide, Rb2O2 and rubidium superoxide, RbO2

      • When rubidium oxide is dissolved in water it forms rubidium hydroxide and the solution turns universal indicator solution or litmus paper blue-purple. Using pH indicator paper or a pH meter you find the alkaline solution has a pH of 13-14.

  • Rubidium is 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) + H2(g)

    • Cesium is much too dangerous a metal to add to acids because of its high reactivity.

  • Rubidium was first extracted in 1861 by electrolysis of the molten chloride RbCl

  • (c) doc b GCSE/IGCSE/O level revision study notes on Group 1 The Alkali Metals

    • where the reactivity trend for group 1 alkali metals is discussed in detail.

  • Advanced Level Inorganic Chemistry Part 7 GCE revision notes on the s-block Group 1 Alkali Metals and Group 2 Alkaline Earth Metals

    TOP OF PAGE and sub-index

potassium K

Reactivity series of metals

Group 1 Alkali Metal

reactivity

The Reactivity Series of Metals

(reactivity of potassium and compared with the non-metals carbon and hydrogen)

 

  • 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 purple-lilac flame when heated in air/oxygen to form the white powder potassium oxide.
    • potassium + oxygen ==> potassium oxide
    • 4K(s) + O2(g) ==> 2K2O(s)

      • Potassium is oxidised, oxygen gain, oxidation reaction.

      • also forms potassium peroxide, K2O2 and potassium superoxide, KO2

      • When potassium oxide is dissolved in water it forms potassium hydroxide and the solution turns universal indicator solution or litmus paper blue-purple. Using pH indicator paper or a pH meter you find the alkaline solution has a pH of 13-14.

  • 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.
    • This is always a great demonstration to do as a chemistry teacher - but don't cut to big a piece 'guys' - even with a safety screen and safety glasses/goggles, it can be a bit unpredictable and dangerous!
    • potassium + water ==> potassium hydroxide + hydrogen
    • 2K(s) + 2H2O(l) ==> 2KOH(aq) + H2(g)

    • Potassium is much too dangerous a metal to add to acids because of its high reactivity.

  • Potassium was first extracted in 1807 by electrolysis of the molten chloride KCl

  • (c) doc b GCSE/IGCSE/O level revision study notes on Group 1 The Alkali Metals

  • (c) doc b Advanced Level Inorganic Chemistry Part 7 GCE revision notes on the s-block Group 1 Alkali Metals and Group 2 Alkaline Earth Metals

    TOP OF PAGE and sub-index

sodium Na

Reactivity series of metals

Group 1 Alkali Metal

reactivity

The Reactivity Series of Metals

(reactivity of sodium and compared with the non-metals carbon and hydrogen)

 

  • 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) + O2(g) ==> 2Na2O(s)
      • Sodium is oxidised, oxygen gain, oxidation reaction.
      • also forms some sodium peroxide, Na2O2

      • When sodium oxide is dissolved in water it forms sodium hydroxide and the solution turns universal indicator solution or litmus paper blue-purple. Using pH indicator paper or a pH meter you find the alkaline solution has a pH of 13-14.

  • 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 + 2H2O ==> 2NaOH + H2
      • 2Na(s) + 2H2O(l) ==> 2NaOH(aq) + H2(g)

    • Sodium is much too dangerous a metal to add to acids because of its high reactivity.

  • Sodium was first  extracted in 1807 by electrolysis of the molten chloride NaCl. Extraction of sodium

lithium Li

Reactivity series of metals

Group 1 Alkali Metal

reactivity

The Reactivity Series of Metals

(reactivity of lithium and compared with the non-metals carbon and hydrogen)

  • 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 reddish-crimson flame when heated in air/oxygen to form the white powder lithium oxide.

    • lithium + oxygen ==> lithium oxide
    • 4Li(s) + O2(g) ==> 2Li2O(s)

    • Lithium is oxidised, oxygen gain, oxidation reaction.

    • When lithium oxide is dissolved in water it forms lithium hydroxide and the solution turns universal indicator solution or litmus paper blue-purple. Using pH indicator paper or a pH meter you find the alkaline solution has a pH of ~13.

  • Lithium has 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) + H2(g)
    • By the time you get down to lithium, the electron is not quite as easily lost to form the positive ion (Li+)

  • Lithium is perhaps too dangerous a metal to add to acids because of its high reactivity?

  • Lithium was first extracted in 1821 by electrolysis of the molten chloride LiCl

  • (c) doc b GCSE/IGCSE/O level revision study notes on Group 1 The Alkali Metals

  • (c) doc b Advanced Level Inorganic Chemistry Part 7 GCE revision notes on the s-block Group 1 Alkali Metals and Group 2 Alkaline Earth Metals

    TOP OF PAGE and sub-index

calcium Ca

Reactivity series of metals

Group 2 Alkaline Earth Metal

reactivity

The Reactivity Series of Metals

(reactivity of calcium and compared with the non-metals carbon and hydrogen)

 
  • 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) + O2(g) ==> 2CaO(s)
    • Calcium is oxidised, oxygen gain, oxidation reaction.

    • When calcium oxide is slightly soluble in water and forms calcium hydroxide and the solution turns universal indicator solution or litmus paper blue-purple. Using pH indicator paper or a pH meter you find the alkaline solution has a pH of ~13.

  • Calcium is 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) + H2(g)
  • Calcium is 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) ==> CaCl2(aq) + H2(g)
    • Ionic equation: Ca(s) + 2H+(aq) ==> Ca2+(aq) + H2(g)
    • In this reaction the calcium atom is oxidised and loses electrons to form the calcium ion (Ca2+). The hydrogen ion (H+) from the acid is reduced by electron gain to give the hydrogen molecule (H2).
    • See 'explaining oxidation and reduction' with lots of examples!
  • Calcium is not very reactive with dilute sulfuric acid because the colourless calcium sulfate formed is not very soluble and coats the metal inhibiting the reaction, so not many bubbles of hydrogen.
    • calcium + sulfuric acid ==> calcium sulfate + hydrogen
    • Ca(s) + H2SO4(aq) ==> CaSO4(aq/s) + H2(g)
  • Calcium was first extracted in 1808 by electrolysis of the molten chloride CaCl2

  • See also setting up metal reactivity series experiments–observations-deductions

  • Advanced Level Inorganic Chemistry Part 7 GCE revision notes on the s-block Group 1 Alkali Metals and Group 2 Alkaline Earth Metals

    TOP OF PAGE and sub-index

magnesium Mg

Reactivity series of metals

Group 2 Alkaline Earth Metal

reactivity

The Reactivity Series of Metals

(reactivity of magnesium and compared with the non-metals carbon and hydrogen)

 
  • 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) + O2(g) ==> 2MgO(s)
    • Magnesium is oxidised, oxygen gain, oxidation reaction.

    • When magnesium oxide is very slightly soluble in water and forms magnesium hydroxide and the solution turns universal indicator solution or litmus paper blue-purple. Using pH indicator paper or a pH meter you find the alkaline solution has a pH of ~12.

  • Magnesium reacts slowly with cold 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) + H2(g)

  • If magnesium is heated in steam, the magnesium will burn with a bright white flame and the white powder magnesium oxide is formed and hydrogen gas.

    • The reaction between cold water and magnesium is slow. but heating both reactants speeds up the reaction, but note that magnesium oxide is formed and not magnesium hydroxide.

    • You can ignite a strip of magnesium in a bunsen flame and plunge it carefully into steam above a flask of boiling water ie heating magnesium to a high temperature in steam.
    • magnesium + water ==> magnesium oxide + hydrogen
    • Mg(s) + H2O(g) ==> MgO(s) + H2(g)
    • Mg oxidised, O gain and H2O reduced, O loss.
    • OR you can do the experiment in a boiling tube as illustrated below.
    • In this 2nd method you can burn the hydrogen formed!
    • reacting magnesium in steam & burning the product hydrogen

  • In fact magnesium is so reactive, it will even burn in carbon dioxide, the products being white magnesium oxide powder and black specks of elemental carbon!

    • You can ignite a strip of magnesium held on the end of a deflagrating spoon and lid, plunge into a gas jar of carbon dioxide, replace the lid-spoon and it will continue to burn.
    • magnesium + carbon dioxide ==> magnesium oxide + carbon
    • 2Mg(s) + CO2(g) ==> 2MgO(s) + C(s)
    • Mg oxidised, O gain, CO2 reduced, O loss.
    • Magnesium is oxidised (oxygen gain) and carbon dioxide is reduced (oxygen loss)
    • See 'explaining oxidation and reduction' with lots of examples!
  • Magnesium is very reactive with dilute hydrochloric acid forming the colourless soluble salt magnesium chloride and hydrogen gas.
    • magnesium + hydrochloric acid ==> magnesium chloride + hydrogen
    • Mg + 2HCl ==> MgCl2 + H2
      • Mg(s) + 2HCl(aq) ==> MgCl2(aq) + H2(g)
      • Ionic equation: Mg(s) + 2H+(aq) ==> Mg2+(aq) + H2(g)
    • In this reaction the magnesium atom is oxidised and loses electrons to form the magnesium ion (Mg2+). The hydrogen ion (H+) from the acid is reduced by electron gain to give the hydrogen molecule (H2).
    • See 'explaining oxidation and reduction' with lots of examples!
  • Magnesium is very reactive with dilute sulfuric acid forming colourless soluble magnesium sulfate and hydrogen.
    • magnesium + sulfuric acid ==> magnesium sulfate + hydrogen
    • Mg(s) + H2SO4(aq) ==> MgSO4(aq) + H2(g)
    • Ionic equation: Mg(s) + 2H+(aq) ==> Mg2+(aq) + H2(g)
    • Again, in this reaction the magnesium atom is oxidised and loses electrons to form the magnesium ion (Mg2+). The hydrogen ion (H+) from the acid is reduced by electron gain to give the hydrogen molecule (H2).
  • 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) + 2HNO3(aq) ==> Mg(NO3)2(aq) + H2(g)
    • which competes with the reaction ...
    • (ii) magnesium + nitric acid ==> magnesium nitrate + water + nitrogen(II) oxide [nitric oxide]
    • 3Mg(s) + 8HNO3(aq) ==> 3Mg(NO3)2(aq) + 4H2O(l) + 2NO(g)
    • and followed rapidly by ...
    • (iii) nitrogen(II) oxide + oxygen ==> nitrogen(IV) oxide
    • 2NO(g) + O2(g) ==> 2NO2(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) + 4HNO3(aq) ==> Mg(NO3)2(aq) + 2H2O(l) + 2NO2(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)
    • Copper oxide is reduced (oxygen loss) and magnesium is oxidised (oxygen gain).
  • (ii) Adding magnesium powder to copper(II) sulfate solution, remove the blue colour of the copper(II) salt, leaving a colourless solution of magnesium sulfate and a pinky-brown deposit of copper.
  • magnesium + copper sulfate ==> magnesium sulfate + copper
    • Mg(s) + CuSO4(aq) ==> MgSO4(aq) + Cu(s)
    • ionic equation: Mg(s) + Cu2+(aq) ==> Mg2+(aq) + Cu(s)
    • Magnesium atoms (Mg) are oxidised to magnesium ions (Mg2+) by electron loss, and the copper ions (Cu2+) are reduced to copper atoms by electron gain.
  • Magnesium was first extracted in 1808 by electrolysis of the molten chloride MgCl2
  • See 'explaining oxidation and reduction' with lots of examples!

  • and setting up metal reactivity series experiments–observations-deductions

  • Advanced Level Inorganic Chemistry Part 7 GCE revision notes on the s-block Group 1 Alkali Metals and Group 2 Alkaline Earth Metals

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aluminium Al

Reactivity series of metals

Group 3 Metal

reactivity

The Reactivity Series of Metals

(reactivity of aluminum and compared with the non-metals carbon and hydrogen)

 

  • 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 aluminium is 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.
  • This property of aluminium makes it a useful metal for out-door purposes e.g. aluminium window frames, greenhouse frames.

    • aluminium + oxygen ==> aluminium oxide

    • 4Al(s) + 3O2(g) ==> 2Al2O3(s)

    • Aluminium is oxidised, oxygen gain, oxidation reaction.

    • Aluminium oxide is insoluble with water.

  • 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) sulfate 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) sulfate solution is added to a zinc-dilute sulfuric 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 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 Thermit 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)

  • Equation and redox theory applied to the Thermite reaction
    • aluminium + iron(III) oxide ==>  aluminium oxide + iron
    • 2Al(s) + Fe2O3(s) ==> Al2O3(s) + 2Fe(s)
    • The iron oxide is reduced to iron
      • reduction is oxygen loss (Fe2O3 ==> Fe),
        • aluminium is the reducing agent, gains the lost oxygen
      • Fe3+ ions in Fe2O3 gain three electrons to form Fe atoms
        • (electron gain is a more advanced definition of reduction),
    • The aluminium is oxidised to aluminium oxide.
      • oxidation is oxygen gain (Al ==> Al2O3),
        • technically, iron oxide is the oxidising agent
      • Al atoms lose three electrons to form Al3+ ions (in Al2O3)
        • (electron loss is a more advanced definition of oxidation)
    • See 'explaining oxidation and reduction' with lots of examples!

  • This is a typical displacement reaction by a more reactive metal displacing a less reactive metal from one of its compounds.

    • In the blast furnace iron is displaced from iron oxide by using cheap carbon as the reducing agent.
    • Aluminium is an expensive metal made by the costly process of electrolysis (Extraction of Aluminium), so the Thermit reaction would be a ridiculously expensive way of producing iron!
  • 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) ==> 2AlCl3(aq) + 3H2(g)
    • Ionic equation: 2Al(s) + 6H+(aq) ==> 2Al3+(aq) +  3H2(g)
    • In this reaction the aluminium atom is oxidised and loses electrons to form the aluminium ion (Al3+). The hydrogen ion (H+) from the acid is reduced by electron gain to give the hydrogen molecule (H2).
  • The reaction with dilute sulfuric acid is very slow to form colourless  aluminium sulfate and hydrogen. (see above)
    • aluminium + sulfuric acid ==> aluminium sulfate + hydrogen
    • 2Al(s) + 3H2SO4(aq) ==> Al2(SO4)3(aq) + 3H2(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) sulfate solution and a patchy pink colour of copper metal slowly appears over many hours/days?
    • aluminium + copper(II) sulfate ==> aluminium sulfate + copper
    • 2Al(s) + 3CuSO4(aq) ==> Al2(SO4)3(aq) + 3Cu(s)

  • See also setting up metal reactivity series experiments–observations-deductions

  • Aluminium was first extracted in 1825 by electrolysis of its molten oxide Al2O3 (bauxite ore).

(Carbon C, a non-metal)

reactivityElements 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

Reactivity series of metals

At the end of the 1st block-series of Transition Metals

reactivity

The Reactivity Series of Metals

(reactivity of zinc and compared with the non-metals carbon and hydrogen)

 

  • 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) + O2(g) ==> 2ZnO(s)
    • Zinc is oxidised, oxygen gain, oxidation reaction.
    • Zinc oxide is insoluble with water.

  • Zinc has no reaction with cold water.

  • When the zinc is heated strongly in steam zinc oxide and hydrogen are formed.
    • zinc + water ==> zinc oxide + hydrogen
    • Zn(s) + H2O(g) ==> ZnO(s) + H2(g)
    • Zn oxidised, O gain, H2O reduced, O loss.
  • Zinc is 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) ==> ZnCl2(aq) + H2(g)
    • Ionic equation: Zn(s) + 2H+(aq) ==>  Zn2+(aq) + H2(g)
    • In this reaction the zinc atom is oxidised and loses electrons to form the zinc ion (Zn2+). The hydrogen ion (H+) from the acid is reduced by electron gain to give the hydrogen molecule (H2).
  • Zinc is quite reactive with dilute sulfuric acid forming the colourless soluble salt zinc sulfate and hydrogen gas.
    • zinc + sulfuric acid ==> zinc sulfate + hydrogen
    • Zn(s) + H2SO4(aq) ==> ZnSO4(aq) + H2(g)
    • (this reaction is catalysed by adding a trace of copper sulfate solution which form a deposit on the zinc surface)
    • Again, in this reaction the zinc atom is oxidised and loses electrons to form the zinc ion (Zn2+). The hydrogen ion (H+) from the acid is reduced by electron gain to give the hydrogen molecule (H2).
  • Zinc forms very little hydrogen with dilute nitric acid, though zinc nitrate is formed. 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) + 2HNO3(aq) ==> Zn(NO3)2(aq) + H2(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) + 8HNO3(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) + O2(g) ==> 2NO2(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
    • Zn(s) + 4HNO3(aq) ==> Zn(NO3)2(aq) + 2H2O(l) + 2NO2(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) sulfate solution, removes the blue colour of the copper(II) salt, leaving a colourless solution of zinc sulfate and a pinky-brown deposit of copper.
    • zinc + copper sulfate ==> zinc sulfate + copper
    • Zn(s) + CuSO4(aq) ==> ZnSO4(aq) + Cu(s)
    • ionic equation: Zn(s) + Cu2+(aq) ==> Zn2+(aq) + Cu(s)
    • The zinc atom is oxidised to the zinc ion by electron loss and the copper ion is reduced to a copper atom by electron gain.
    • This is a typical displacement reaction by a more reactive metal displacing a less reactive metal from one of its compounds.
    • See 'explaining oxidation and reduction' with lots of examples!

    • and setting up metal reactivity series experiments–observations-deductions

  • 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) + CO2(g)
  • zinc oxide is reduced (oxygen loss) and carbon is oxidised (oxygen gain).
  • A zinc coating (galvanising) is used to protect iron from rusting. The more reactive zinc oxidises 1st. 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.
  • (c) doc b Extraction of Zinc notes
  • Advanced Level Inorganic Chemistry Part 10 GCE revision notes 3d block TRANSITION METALS including zinc

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iron Fe

Reactivity series of metals

In the 1st block-series of Transition Metals

reactivity

The Reactivity Series of Metals

(reactivity of iron and compared with the non-metals carbon and hydrogen)

 

  • 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) + 2O2(g) ==> Fe3O4(s)
      • Iron oxide is insoluble with water.
      • Iron is oxidised, oxygen gain, oxidation reaction.

  • Iron has no reaction with cold water to form hydrogen (rusting is a joint reaction with oxygen).

  • When iron 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) ==> Fe3O4(s) + 4H2(g)
    • This is a reversible reaction - if you pass hydrogen over heated iron tetroxide it is reduced to iron and water is formed.
    • water is reduced (oxygen loss) and iron is oxidised (oxygen gain).
    • iron tetroxide + hydrogen ==> iron + water (condenses)
      • Fe3O4(s) + 4H2(g) ==> 3Fe(s) + 4H2O(g)
      • iron oxide is reduced (oxygen loss) and hydrogen is oxidised (oxygen gain).
  • Iron has a relative slow-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) ==> FeCl2(aq) + H2(g)
      • Ionic equation: Fe(s) + 2H+(aq) ==> Fe2+(aq) + H2(g)
      • In this reaction the iron atom is oxidised and loses electrons to form the iron(II) ion (Fe2+). The hydrogen ion (H+) from the acid is reduced by electron gain to give the hydrogen molecule (H2).
    • 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).
  • Iron has a slow reaction with dilute sulfuric acid forming the soluble pale green salt iron(II) sulfate and hydrogen gas.
    • iron + sulfuric acid ==> iron(II) sulfate + hydrogen
      • Fe(s) + H2SO4(aq) ==> FeSO4(aq) + H2(g)
      • In this reaction the iron atom is oxidised and loses electrons to form the iron(II) ion (Fe2+). The hydrogen ion (H+) from the acid is reduced by electron gain to give the hydrogen molecule (H2).
  • 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
      • Fe2O3(s) + 3CO(g) ==> 2Fe(l-s) + 3CO2(g)
    • iron tetroxide + carbon monoxide ==> iron + carbon dioxide
      • Fe3O4(s) + 4CO(g) ==> 3Fe(l-s) + 4CO2(g)
    • See 'explaining oxidation and reduction' with lots of examples!
  • Iron will displace less reactive metals from their salt solutions.

tin Sn

Reactivity series of metals

A Group 4 metal

reactivity

The Reactivity Series of Metals

(reactivity of tin and compared with the non-metals carbon and hydrogen)

 

  • Tin has slow reaction when heated in air to slowly form white tin(IV) oxide or tin dioxide

    • tin + oxygen ==> tin oxide [tin dioxide, tin(IV oxide]
    • Sn(s) + O2(g) ==> SnO2(s)
    • Tin is oxidised, oxygen gain, oxidation reaction.
    • Tin oxide is insoluble with water.

  • Tin has no reaction with cold water or when heated in steam.

  • Tin has a 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) ==> SnCl2(aq) + H2(g)
    • Ionic equation: Sn(s) + 2H+(aq) ==> Sn2+(aq) + H2(g)
    • In this reaction the tin atom is oxidised and loses electrons to form the tin(II) ion (Sn2+). The hydrogen ion (H+) from the acid is reduced by electron gain to give the hydrogen molecule (H2).
    • See 'explaining oxidation and reduction' with lots of examples!
  • Tin has a very slow reaction with dilute sulfuric acid forming the colourless slightly soluble tin(II) sulfate and hydrogen gas.
    • tin + sulfuric acid ==> tin(II) sulfate + hydrogen
    • Sn(s) + H2SO4(aq) ==> SnSO4(aq) + H2(g)

  • See also setting up metal reactivity series experiments–observations-deductions

  • 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!
  • Tin's lack of reactivity enables it to be used as a protective layer in steel cans of fruit - tinned cans!

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lead Pb

Reactivity series of metals

A Group 4 metal

reactivity

  • Reactivity of lead

  • Lead has a 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) + O2(g) ==> 2PbO(s)

    • and 3Pb(s) + 2O2(g) ==> Pb3O4(s)

    • Lead is oxidised, oxygen gain, oxidation reaction.
    • Lead oxides are insoluble with water.

  • Lead has no reaction with cold water or when heated in steam.

  • Lead reacts very slowly, and effectively, no real reaction with dilute hydrochloric acid or dilute sulfuric acid.

  • See also setting up metal reactivity series experiments–observations-deductions

  • 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!
  • Lead's lack of reactivity has enabled it in the past to be used for water pipes, though it is being replaced by plastic tubing or piping for two reasons - (i) lead is a toxic metal and plastic is cheaper!

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Hydrogen H

non-metal

reactivityNon 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

Reactivity series of metals

In the 1st block-series of Transition Metals

reactivity

The Reactivity Series of Metals

(reactivity of copper and compared with the non-metals carbon and hydrogen)

 

  • Surface blackens when a copper strip is strongly heated in air/oxygen to form copper(II) oxide (you see flashes of green and blue in the flame prior to the formation of the black layer of copper(II) oxide.

    • copper + oxygen ==> copper oxide [copper(II) oxide]

    • 2Cu(s) + O2(g) ==> 2CuO(s)

    • Copper is oxidised, oxygen gain, oxidation reaction.
    • Copper(II) oxide is insoluble with water.

  • Copper has no reaction with cold water or when heated in steam.

  • Copper has no reaction with dilute hydrochloric acid or dilute sulfuric acid.
  • Copper can be extracted by reducing the hot black metal oxide on heating with carbon
  • Although copper doesn't readily react with dilute hydrochloric acid and dilute sulfuric acid (low in reactivity series), if heated with nasty oily concentrated sulfuric acid you make nasty pungent irritating sulphur dioxide gas and white anhydrous copper(II) sulfate, but this is NOT a reaction on which to base its place in the metal reactivity series and hydrogen gas isn't produced.
  • 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) + 8HNO3(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) + O2(g) ==> 2NO2(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) + 4HNO3(aq) ==> Cu(NO3)2(aq) + 2H2O(l) + 2NO2(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 as 'native copper' 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!
  • Copper can be used for roofing, where it corrodes superficially, and very slowly, to give a green protective layer of a basic carbonate (its a mixture of insoluble hydroxide and carbonate).
  • Advanced Level Inorganic Chemistry Part 10 GCE revision notes 3d block TRANSITION METALS including copper

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silver Ag

 Reactivity series of metals

a transition metal (2nd series)

silver is very low in the reactivity series of metals

  • Reactivity of silver

  • Silver has no reaction when heated in air, it isn't oxidised.

  • Silver has no reaction with cold water or when heated in steam.

  • Silver has no reaction with dilute hydrochloric acid or dilute sulfuric acid.
  • Metals like silver are very unreactive because they do not readily lose electrons to form a positive ion.
  • Silver reacts with hot concentrated sulfuric acid to form silver sulfate and sulfur dioxide gas.
  • Silver reacts with hot concentrated nitric acid to form silver nitrate and gaseous nitrogen oxides.
  • 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.
  • Its very low reactivity makes it a valuable jewellery metal as it doesn't corrode easily and retains its attractive silvery appearance.

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gold Au

Reactivity series of metals

a transition metal (3rd series)

gold is very low in the reactivity series of metals

  • Reactivity of gold

  • Gold has no reaction when heated in air, it isn't oxidised.

  • Gold has no reaction with cold water or when heated in steam.

  • Gold has no reaction with dilute hydrochloric acid or dilute sulfuric acid.

  • Metals like gold are very unreactive because they do not readily lose electrons to form a positive ion.

  • Gold will react with, and dissolve in, a mixture of concentrated nitric acid and concentrated hydrochloric acid (known as 'aqua regia') to form gold(III) chloride.

  • 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! 24/24 to 9/24 fraction of gold!
  • Gold's extremely low reactivity makes it a valuable jewellery metal as it doesn't corrode easily and retains its shiny attractive yellow appearance.

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platinum Pt

Reactivity series of metals

a transition metal (3rd series)

 platinum is so low in the reactivity series of metals, its less reactive than gold!

  • Reactivity of platinum

  • Platinum has no reaction when heated in air, it isn't oxidised.

  • Platinum has no reaction with cold water or when heated in steam.

  • Platinum has no reaction with dilute hydrochloric acid or dilute sulfuric acid.
  • Metals like platinum are very unreactive because they do not readily lose electrons to form a positive ion.
  • 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.
  • Platinum is used in expensive jewellery, laboratory ware (e.g. inert crucible container) and catalytic converters in car exhausts.
  • Platinum's very low reactivity makes it a valuable jewellery metal as it doesn't corrode easily and retains its attractive silvery appearance.
  • Platinum crucibles are used for some high temperature chemical procedures because they are so stable and unreactive.

(d) OTHER ASSOCIATED PAGE LINKS

SEE ALSO (c) doc b 2. RUSTING & Introducing REDOX reactions

and 3. (c) doc b Metal Reactivity Series Experiments-Observations

and GCSE/IGCSE m/c QUIZZES on metal reactivity

Foundation-tier Level (easier) multiple choice quiz on the Reactivity Series of Metals

or Higher-tier Level (harder) multiple choice quiz on the Reactivity Series of Metals

and (c) doc b GCSE/IGCSE reactivity gap-fill worksheet or (c) doc b Rusting word-fill worksheet

KS4 Science GCSE/IGCSE/O level Chemistry revision notes pages:

(c) doc b The Periodic Table  *  (c) doc b Group 1 Alkali Metals  *  (c) doc b Methods of Metal extraction

(c) doc b Transition Metals  *  (c) doc b Alloys-uses of metals  *  (c) doc b Electrochemistry-Electrolysis

(c) doc b Rates of Reactions Experiments (e.g. metal-acid)

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