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GCSE Chemistry Notes: Describing metal reactivity experiments

Reactivity series of metals   Reactivity series of metals

reactivityDoc Brown's Chemistry KS4 Science GCSE/IGCSE Chemistry Revision Notes

3. Metal Reactivity Series How to set up reactivity series experiments - observations, deducing the reactivity order and theoretically explaining the reactions

What do you see? The observations of the reactivity series of metals reacting with water, metals reacting with hydrochloric acid and metal displacement reactions are described and tabulated below. These revision notes on reactivity series of metals experiment observations should prove useful for the new AQA, Edexcel and OCR GCSE (91) chemistry science courses.

Sub-index of reactions described in this page

3a metal + water * 3b metal + acid

3c metal + salt solution  *  3d metal/carbon + solid oxide

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.

(c) doc b More details like word/symbol equations are given in Part 1.

SEE ALSO 1. Reactivity Series of Metals

2. RUSTING & Introducing REDOX reactions and ...

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

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

reactivity


chemistry revision notes: basic school chemistry science GCSE chemistry, IGCSE  chemistry, O level & ~US grades 8, 9 and 10 school science courses or equivalent for ~14-16 year old science students for national examinations in chemistry for topics including the reactivity series of metals

3a. Observations of the reaction of metals with cold water
  • All you need is a selection of metals, water, universal indicator, rack of test tubes and a splint.

  • Ideally all the metals would be in the same physical state, other than all being solids,

    • i.e. same mass of the same sized granules or same sized sheet of same thickness - to give the same surface area in contact with the acid.

    • However this is isn't possible within a school laboratory (or its budget!).

  • The experiment is extremely easy to do, simply add small similar sized portions of the metal to a few cm3 of cold water containing a few drops of universal indicator.

  • Before adding the metal the indicator should give a green colour indicating neutral ~pH 7.

  • Keep an eye on all of them at the same time and note any changes and the relative rate of evolution of gas bubbles - the latter is a crude measure of the speed of the reaction of the metal with water.

3a. Results: Table of possible observations in alphabetical order

Metal Possible observations
aluminium turnings no change - no reaction
calcium granules Slow reaction, effervescence, colourless gas evolved. The universal indicator turns blue-purple - alkaline solution (pH ~13)
copper strip no change - no reaction
iron filings no change - no reaction
lead strip no change - no reaction
lithium - pea sized lump! reactivityModerately fast reaction - effervescence, colourless gas formed  TEACHER DEMONSTRATION in large trough of water!  The universal indicator turns blue-purple - alkaline solution (pH ~13)
magnesium ribbon Very slow reaction, bubbles slowly form on the surface of the magnesium ribbon. Around the magnesium ribbon the universal indicator turns pale blue - weakly alkaline solution (pH ~11)
potassium - pea sized lump! reactivityVery fast reaction - effervescence, colourless gas formed  TEACHER DEMONSTRATION in large trough of water!  The universal indicator turns blue-purple - alkaline solution (pH ~13)
sodium - pea sized lump! reactivityFast reaction - effervescence, colourless gas formed, may ignite with a larger lump  TEACHER DEMONSTRATION in large trough of water!  The universal indicator turns blue-purple - alkaline solution (pH ~13)
tin strip no change - no reaction
zinc granules no change - no reaction
  • reactivityNo reaction = no observed change, in some cases where theoretically there should be a reaction, you might not see any change.

  • If there is any reaction, bubbles of colourless gas will form.

  • In this experiment you are judging the relative reactivity by observing how fast each metal appears to react with water.

  • This is simply judged by the evolution of colourless gas bubbles of hydrogen, readily tested for using the calcium reaction - the colourless gas gives a 'pop'! with a lit splint - hydrogen.

  • Any change in the indicator colour also indicates a chemical change is taking place.

  • -

3a. Conclusions and comments.

  • From the observations you can reasonably deduce the following reactivity order

    • potassium > sodium > lithium > calcium > magnesium > rest of metals that don't seem to react with cold water

    • but the rest do not react, so all you can say is that they appear less reactive than magnesium.

  • All the equations you need are on the page 1. The Reactivity Series of Metals, but some examples are given below

    • sodium + water ==> sodium hydroxide + hydrogen

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

    • You can write similar equation for lithium and potassium (Li/K instead of Na)

    • and

    • calcium + water ==> calcium hydroxide + hydrogen

    • Ca(s) + 2H2O(l) ==> Ca(OH)2(aq/s) + H2(g)

    • You can write a similar equation for magnesium (Mg instead of Ca)

    • The half equations for the reactions are written as:

    • Na  ==> Na+  + e-  and  Ca  ==>  Ca2+  +  2e-

    • The metal atoms are oxidised by electron loss.

    • Reminder: The more easily the electrons are lost, the more reactive the metal.

  • NOTE: 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

  • Comments on some of uses of the metal relevant to these observations

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

    • Zinc is used in galvanising, i.e. coating iron or steel to prevent rusting. The zinc corrodes preferentially forming a protective oxide layer.

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

    • 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!

    • 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).

  • -

 

3b. The Reaction of Metals with Hydrochloric Acid

  • All you need is a selection of metals, dilute hydrochloric acids, rack of test tubes and a splint.

  • Fair test points:

    • Ideally all the metals would be in the same physical state, other than all being solids.

    • i.e. use the same mass of the sized/shape granules or same sized sheet of same thickness - to give the same surface area in contact with the acid.

    • Powdered metals will automatically react faster than big lumps of another metal, especially if they were of similar, but not identical, reactivity.

    • However this is isn't possible within a school laboratory (or its budget!).

    • You should do the experiments with same concentration of acid and all at the same laboratory temperature.

  • The experiment is extremely easy to do, simply add similar sized small portions of the metal to a few cm3 of cold dilute hydrochloric acid.

  • Keep an eye on all of them at the same time and note any changes and the relative rate of evolution of gas bubbles - which is a crude measure of the relative speed of reaction of the metal with the acid.

  • The observations are somewhat crude unless you use a gas syringe to measure the volume of hydrogen evolved in a certain time.

3b. Results: Table of possible observations

Metal Possible observations
aluminium turnings You may see a few bubbles after a long time!
calcium granules Very fast reaction, effervescence, colourless gas evolved, obviously exothermic - test tube heats up!
copper strip no change - no reaction
iron filings Slow reaction, colourless gas
lead strip no change - no reaction
lithium  MUST NOT BE DONE - TOO FAST and EXPLOSIVE
magnesium ribbon Fast reaction, bubbles rapidly form on the surface of the magnesium ribbon
potassium - pea sized lump!  MUST NOT BE DONE - TOO FAST and EXPLOSIVE
sodium - pea sized lump!  MUST NOT BE DONE - TOO FAST and EXPLOSIVE
tin strip You may see a few bubbles after a long time!
zinc granules Moderately fast reaction, bubbles of colourless gas
  • reactivityNo reaction = no observed change, in some cases, where theoretically there should be a reaction, you might not see any change - this can happen with aluminium.

  • If there is any reaction, bubbles of colourless gas will form.

  • In this experiment you are judging the relative reactivity by observing how fast each metal appears to react with the dilute hydrochloric acid.

  • This is simply judged by the evolution of colourless gas bubbles of hydrogen, readily for tested using the magnesium reaction - the colourless gas gives a 'pop'! with a lit splint.

  • Aluminium, again gives problems with the observations because of the oxide layer inhibiting the reaction with the salt solution of a less reactive metal theoretically it is quite reactive.

  • -

3b. Conclusions and comments.

  • From the observations you can reasonably deduce the following reactivity order

    • calcium > magnesium > zinc > iron

    • but the rest hardly react, so all you can say is that they appear less reactive than iron.

  • All the equations you need are on the page 1. The Reactivity Series of Metals, but some examples are given below

    • magnesium + hydrochloric acid ==> magnesium chloride + hydrogen
    • Mg(s) + 2HCl(aq) ==> MgCl2(aq) + H2(g)
    • zinc + hydrochloric acid ==> zinc chloride + hydrogen
    • Zn(s) + 2HCl(aq) ==> ZnCl2(aq) + H2(g)
    • aluminium + hydrochloric acid ==> aluminium chloride + hydrogen
    • 2Al(s) + 6HCl(aq) ==> 2AlCl3(aq) + 3H2(g)   (a more tricky equation to balance!)
  • These reactions can be seen as redox reactions i.e. an oxidation-reduction reaction in terms of electron loss and gain.

    • Oxidation is electron loss and the neutral metal atom loses electrons to form a positive ion (cation).

      • eg.  Mg  ==>  Mg2+  +  2e-,   Zn  ==>  Zn2+  +  2e-   and  Al  ==>  Al3+  +  3e-

      • Reminder: The more easily the electrons are lost, the more reactive the metal.

    • Reduction is electron gain and the hydrogen ions gain electrons to form a neutral hydrogen molecule.

      • 2H+  +  2e-  ==>  H2

    • These half-equations show the movement and transfer of electrons between the species involved in the reaction.

    • You can then combine these half reaction to give the full ionic equation that excludes spectator ions like the chloride ion, which simply remain in solution.

      • This involves balancing the electron losses and gains, so that no electron shows up in the full ionic equation.

      • These are the full redox equations (ionic equations) for acid-metal reactions and I've included state symbols too.

        • Note the electrons are not shown in the full ionic equations because electrons lost = electrons gained, so they cancel out in a sort of 'algebraic' way.

      • Mg(s)  +  2H+(aq)  ===>  Mg2+(aq)  +  H2(g)

      • -

      • Zn(s)  +  2H+(aq)  ===>  Zn2+(aq)  +  H2(g)

      • -

      • 2Al(s)  +  6H+(aq)  ===>  2Al3+(aq)  + 3H2(aq)    (a bit tricky to balance!)

      • -

  • and for more on redox theory behind displacement reactions

  • You can do this experiment more accurately using a gas syringe system to collect the hydrogen.

  • The speed-rate of the reaction would be measured in cm3/min.

  • Factors affecting the rates of Reaction - theory and methods of measuring the speed of a reaction (c) Doc Brown

  • For more details see How can we measure the speed or rate of a chemical reaction?

 

3c. Metal Displacement Reactions (metal + salt solution)

  • A displacement reaction is where a more reactive element displaces a less reactive element.

  • Displacement reactions involve a simultaneous oxidation and reduction, which is why they are called redox reactions.

    • In this experiment it is important to understand the 'electron' definitions of oxidation and reduction.

    • Oxidation is electron loss - the atom/ion/molecule losing one or more electrons is said to be oxidised.

    • Reduction is electron gain - the atom/ion/molecule gaining one or more electrons is said to be reduced.

    • You must learn these definitions!

    • There are lots more examples on my Introducing REDOX reactions page.

  • All you need is a selection of metals, salt solutions of the same metals (eg. chloride, nitrate or sulphate) and rack of test tubes.

  • The physical state of the metal in terms of granule size or area of sheet doesn't matter in this experiment.

  • You just pop the bits of metals into the solutions and carefully observe what happens on the surface.

  • See also metal extraction experiments in section 3(d)

3c. Results: Table of possible observations

salt\metal aluminium film copper strip iron filings lead strip magnesium ribbon zinc granules
aluminium chloride no reaction - same metal no reaction no reaction no reaction You may see a slight change on the Mg surface no reaction
copper sulphate faint pink coating of copper no reaction - same metal pink-orange-brown-dark? layer of copper on the iron filings, blue colour fades You may see a pink-orange-brown-dark? layer of copper on the lead strip, blue colour fades pink-orange-brown-dark? layer of copper on the magnesium strip, blue colour fades pink-orange-brown-dark? layer of copper on the zinc granules, blue colour fades
iron(II) sulphate theoretically reacts - but doubt if you see anything no reaction no reaction - same metal no reaction 'dark' crystals of iron on the magnesium ribbon 'dark' crystals of iron on the zinc granules
lead(II) nitrate theoretically reacts - but doubt if you see anything no reaction theoretically reacts - but doubt if you see anything no reaction - same metal 'dark' crystals of lead on the magnesium ribbon 'dark' crystals of lead on the zinc granules
magnesium sulphate no reaction no reaction no reaction no reaction no reaction - same metal no reaction
zinc sulphate theoretically reacts - but doubt if you see anything no reaction no reaction no reaction 'dark' crystals of zinc on the magnesium ribbon no reaction - same metal
  • observations when adding magnesium to copper sulfate solution displacement reactionNo reaction = no observed change, in some cases where theoretically there should be a reaction, you might not see any change (see Al note below).

  • Some of the best observations are with the more reactive metals than copper.

  • The picture on the right shows what happens if you put a strip of magnesium in copper sulfate solution.

    • The solution starts of a deep blue of the copper sulfate solution.

    • Then the blue colour fades as the brown copper deposit forms on the magnesium.

    • Eventually most of the magnesium dissolves and the solution becomes colourless as magnesium sulfate is formed (which is colourless!).

  • Aluminium, again gives problems with the observations because of the oxide layer inhibiting the reaction with the salt solution of a less reactive metal.

  • The displaced copper can display a variety of colours depending on how the precipitate-coating forms, and how much of it - I'm afraid that's the way it is!

  • In some cases the metal crystals formed by the displacement reaction are very small and scatter the light so that they can look quite dark - almost black, rather than a silver-shiny precipitate.

3c. Conclusions and comments

  • reactivityAny colour change indicates a displacement reaction has occurred,

  • AND the rule is that

  • a more reactive metal will displace a less reactive metal from its salt solution (irrespective of which salt)

  • The way that you deduce the order is quite simple and logical.

    • e.g. zinc cannot displace magnesium, but it can displace copper, and magnesium displaces both zinc and copper,

    • therefore zinc is less reactive than magnesium and zinc is more reactive than copper,

    • therefore you can deduce the reactivity order for these three metals is Mg > Zn > Cu.

  • Of the series of metals tested magnesium appears to be the most reactive because it displaces all the other metals being investigated here.

  • Copper appears to be the least reactive in this limited series because it does not displace any of the other metals.

  • Therefore by considering what will displace what from the results table, you can quite legitimately deduce that in terms of reactivity order

    • magnesium > aluminium > zinc > iron > lead > copper

    • BUT the observations for aluminium and lead may not show up clearly.

    • What you can definitely deduce from the observations in a more restricted experiment is the reactivity series order

    • magnesium > zinc > iron > copper

    • There can be problems in observations of the reactivity of aluminium and lead due to a protective oxide layer.

  • Examples of displacement equations

  • observations when adding magnesium to copper sulfate solution displacement reactionThe first metal on the left is the most reactive, and the last metal on the right, is the displaced less reactive metal.

    • magnesium + copper sulphate ==> magnesium sulphate + copper
      • Mg(s) + CuSO4(aq) ==> MgSO4(aq) + Cu(s)
      • The blue of the copper sulphate solution fades as colourless magnesium sulfate is formed and the brown deposit of copper increases on the surface of the magnesium.
      • -
    • zinc + copper sulphate ==> zinc sulphate + copper
      • Zn(s) + CuSO4(aq) ==> ZnSO4(aq) + Cu(s)
      • -
    • aluminium + copper(II) sulphate ==> aluminium sulphate + copper
  • Again, these reactions can be seen as redox reactions i.e. an oxidation-reduction reaction in terms of electron loss and gain.

    • Oxidation is electron loss and the neutral metal atom loses electrons to form a positive ion (cation).

      • eg.  Mg  ==>  Mg2+  +  2e-,   Zn  ==>  Zn2+  +  2e-   and  Al  ==>  Al3+  +  3e-

      • Reminder: The more easily the electrons are lost, the more reactive the metal.

    • Reduction is electron gain and the hydrogen ions gain electrons to form a neutral hydrogen molecule.

      • Cu2+  +  2e-  ==>  Cu

    • You can then combine these half reaction to give the full ionic equation that excludes spectator ions like the chloride ion, which simply remain in solution.

      • These are the full ionic equations (redox equations) for metal displacement reactions, and I've included state symbols too.

        • Mg(s)  +  Cu2+(aq)  ==>  Mg2+(aq)  +  Cu(s)

        • concept diagram magnesium oxidised displacing copper ion reduced oxidation = electron loss reduction = electron gain
        • Zn(s)  +  Cu2+(aq)  ==>  Zn2+(aq)  +  Cu(s)

        • concept diagram zinc oxidised displacing copper ion from copper sulfate reduced oxidation = electron loss reduction = electron gain

        • 2Al(s)  +  2Cu2+(aq)  ==>  2Al3+(aq)  + 3Cu(s)  (a bit tricky!)

        • Note the electrons are not shown because electrons lost = electrons gained, so cancel out.

  • and for more on redox theory behind displacement reactions


3d. Simple Metal Extraction Experiments  (with reducing agents)

  • Metals of lower reactivity (below aluminium) can be displaced-extracted using carbon as the reducing agent.
  • Heating oxides or carbonates with powdered charcoal (mainly carbon)
    • If you strongly heat copper(II) carbonate with finely powdered charcoal (mainly carbon) you can reddish-brown specks of copper in resulting mixture. The dark green copper(II) carbonate turns black initially as copper(II) oxide is formed, but this is then reduced to copper by the charcoal.
      • initially a thermal decomposition: CuCO3 ==> CuO + CO2
      • then the reduction reaction (O loss): 2CuO + C ==> 2Cu + CO2
      • You can of course start the experiment with copper(II) oxide, but copper carbonate is closer to the sort of naturally occurring copper ore that is mined.
    • You can do a similar experiment by heating lead(II) oxide with powdered carbon and you can get silvery lead formed BUT in a fume cupboard please, since lead fumes are very poisonous!
      • 2PbO + C  ==> 2Pb + CO2
      • Lead is a sufficiently unreactive metal for carbon to displace it.
    • If you heat iron oxides with powdered carbon nothing happens however strongly you heat the mixture in a test tube. Although carbon is reactive enough to displace iron, the temperature in the test tube isn't high enough - you need a blast furnace!
    • If you heat the white powder of aluminium oxide with carbon powder, nothing happens because carbon isn't 'reactive' enough to displace aluminium ie aluminium is too reactive forming very stable compounds.
  • Metal displacement reactions with solid oxides plus reactive metal
    • Here you can use a more reactive metal to displace a less reactive one from its oxide.
    • Your teacher can demonstrate two metal extraction displacement reactions.
    • Neither are used commercially, but they do illustrate how chromium and titanium are extracted.
    • Both these reactions are very exothermic - lots of heat released
    • The Thermit Reaction
      • You ignite a mixture of brown iron(III) oxide and silvery grey aluminium powder using a magnesium fuse.
      • The mixture goes off like a firework in a shower of sparks - a VERY exothermic reaction!
      • On examining the cold residue you find a lump of iron and specks of white aluminium oxide.
      • 2Al + Fe2O3 ==> Al2O3 + 2Fe
      • Aluminium is more reactive than iron and so will displace iron from iron compounds.
    • Heating a mixture of magnesium powder and copper(II) oxide
      • Silvery grey magnesium powder is mixed with black copper oxide powder and heated strongly in a test tube.
      • The mixture glows red hot and on examining the cold mixture you see specks of white magnesium oxide and reddish-brown specks of copper.
      • Mg + CuO ==> MgO + Cu
      • Magnesium is more reactive than copper and so will displace copper from copper compounds.
  • See also Metal Displacement Reactions (metal + salt solution) section 3c
  • and 'The Reactivity Series of Metals'


4. Action of heat on carbonates

Often, but not always, the more reactive a metal, the more thermally stable is the metal carbonate i.e.

reactivity trend with water and acids: sodium > calcium > copper,

thermal stability of carbonate: Na2CO3 > CaCO3 > CuCO3

BUT many carbonates do not fit into such a sequence, so you can't regards this as a general rule

See notes on the Limestone: Detailed section on the thermal decompositions of carbonates page


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reactivity

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

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

  Easy KS3 science multiple choice quiz start on metal reactivity and (c) doc b KS3 word-fills

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)


Revision notes on reactivity series of metals observations KS4 Science GCSE/IGCSE/O level Chemistry Information on reactivity series of metals observations for revising for AQA GCSE Science, Edexcel Science chemistry IGCSE Chemistry notes on reactivity series of metals observations OCR 21st Century Science, OCR Gateway Science notes on reactivity series of metals observations WJEC gcse science chemistry notes on reactivity series of metals observations CIE O Level chemistry CIE IGCSE chemistry notes on reactivity series of metals observations CCEA/CEA gcse science chemistry (help for courses equal to US grade 8, grade 9 grade 10) science chemistry courses revision guides explanation chemical equations for reactivity series of metals observations educational videos on reactivity series of metals observations guidebooks for revising reactivity series of metals observations textbooks on reactivity series of metals observations what do you see when aluminium reacts with water, what do you see when calcium reacts with water, what do you see when copper reacts with water, what do you see when iron reacts with water, what do you see when lead reacts with water, what do you see when lithium reacts with water, what do you see when magnesium reacts with water, what do you see when potassium reacts with water, what do you see when sodium reacts with water, what do you see when tin reacts with water, what do you see when zinc reacts with water what do you observe when aluminium reacts with dilute hydrochloric acid, what do you observe when calcium reacts with dilute hydrochloric acid, what do you observe when copper reacts with dilute hydrochloric acid, what do you observe when iron reacts with dilute hydrochloric acid, what do you observe when lead reacts with dilute hydrochloric acid, what do you observe when magnesium reacts with dilute hydrochloric acid, tin reacts with dilute hydrochloric acid, what do you observe when zinc reacts with dilute hydrochloric acid what do you observe when aluminium reacts with salt solutions of other metals in displacement reactions?, what do you observe when copper reacts with salt solutions of other metals in displacement reactions?, what do you observe when iron reacts with salt solutions of other metals in displacement reactions?, what do you observe when lead reacts with salt solutions of other metals in displacement reactions?, what do you observe when magnesium reacts with salt solutions of other metals in displacement reactions?, what do you observe when zinc reacts with salt solutions of other metals in displacement reactions? gcse chemistry revision free detailed notes on doing metal reactivity series experiments observations to help revise igcse chemistry igcse chemistry revision notes on doing metal reactivity series experiments observations O level chemistry revision free detailed notes on doing metal reactivity series experiments observations to help revise gcse chemistry free detailed notes on doing metal reactivity series experiments observations to help revise O level chemistry free online website to help revise doing metal reactivity series experiments observations for gcse chemistry  free online website to help revise doing metal reactivity series experiments observations for igcse chemistry free online website to help revise O level doing metal reactivity series experiments observations chemistry how to succeed in questions on doing metal reactivity series experiments observations for gcse chemistry how to succeed at igcse chemistry how to succeed at O level chemistry a good website for free questions on doing metal reactivity series experiments observations to help to pass gcse chemistry questions on doing metal reactivity series experiments observations a good website for free help to pass igcse chemistry with revision notes on doing metal reactivity series experiments observations a good website for free help to pass O level chemistry

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