GCSE chemistry Metal reactivity series experiments results observations conclusions deducing reactivity order gcse igcse KS4 science chemistry revision notes revising O level notes

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!	Moderately 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. The solution around the magnesium ribbon turns the universal indicator pale blue - formation of a weakly alkaline solution (pH ~11)
potassium - pea sized lump!	Very 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!	Fast 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

No 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) + H_2(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) + H_2(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 i.e. heating magnesium to a high temperature in steam.
- magnesium + water ==> magnesium oxide + hydrogen
- Mg(s) + H₂O(g) ==> MgO(s) + H_2(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!
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).

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

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

No 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) ==> MgCl_2(aq) + H_2(g)
- zinc + hydrochloric acid ==> zinc chloride + hydrogen
- Zn(s) + 2HCl(aq) ==> ZnCl_2(aq) + H_2(g)
- aluminium + hydrochloric acid ==> aluminium chloride + hydrogen
- 2Al(s) + 6HCl(aq) ==> 2AlCl_3(aq) + 3H_2(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.
For more details see How can we measure the speed or rate of a chemical reaction?

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

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

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 sulfate) 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 sulfate	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) sulfate	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 sulfate	no reaction	no reaction	no reaction	no reaction	no reaction - same metal	no reaction
zinc sulfate	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

No 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.

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3c. Conclusions and comments

Any 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 because oxide layer inhibits the reaction with water, acids or with another less reactive metal salt solution.
- 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
The first metal on the left is the most reactive, and the last metal on the right, is the displaced less reactive metal.
- magnesium + copper sulfate ==> magnesium sulfate + copper
- Mg(s) + CuSO_4(aq) ==> MgSO_4(aq) + Cu_(s)
- The blue of the copper sulfate solution fades as colourless magnesium sulfate is formed and the brown deposit of copper increases on the surface of the magnesium.
  - -
- zinc + copper sulfate ==> zinc sulfate + copper
- Zn(s) + CuSO_4(aq) ==> ZnSO_4(aq) + Cu_(s)
  - _-
- aluminium + copper(II) sulfate ==> aluminium sulfate + copper
- 2Al(s) + 3CuSO₄(aq) ==> Al₂(SO₄)₃(aq) + 3Cu(s)
  - -
See also 'The Reactivity Series of Metals'
- -
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)
- - -
  - Zn(s) + Cu²⁺(aq) ==> Zn²⁺(aq) + Cu(s)
- - 2Al(s) + 2Cu²⁺(aq) ==> 2Al³⁺(aq) + 3Cu(s) (a bit tricky to balance!)
  - Note the electrons are not shown because electrons lost = electrons gained, so cancel out.
and for more on redox theory behind displacement reactions

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

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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 i.e. aluminium is too reactive forming very stable compounds.
  - For more see notes on Metal Extraction Fe, Cu, Al etc.

3e. Metal displacement reactions with solid oxides plus reactive metal as the reducing agent

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'

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

Higher-tier Level (harder) multiple choice quiz on 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

and advanced level chemistry notes on the thermal decomposition of carbonates

What next? Associated Pages

Index of notes on reaction of acids

GCSE Level (~US grades 8-10) School Chemistry Notes (students age ~14-16)

Find your GCSE science course for more help links to revision notes

Advanced Level (~US grades 11-12) pre-university Chemistry Notes (for aged ~17-18)

Email doc brown - comment? query?

OTHER ASSOCIATED PAGE LINKS to do with metal reactivity

2. RUSTING & Introducing REDOX reactions

3. Metal Reactivity Series Experiments-Observations

and GCSE level m/c QUIZZES on metal reactivity

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

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

GCSE level reactivity gap-fill worksheet

and Rusting word-fill worksheet

Other QUIZ links to do with metals (UK GCSE and ~US grade 9-10 level)

Higher level QUIZ on Group 1 Alkali Metals and Transition Metals

Lower level QUIZ on Group 1 Alkali Metals and Transition Metals

Higher level QUIZ on The relative reactivity and extraction of metals

Lower level QUIZ on The relative reactivity and extraction of metals

Higher level QUIZ on Group 1 Alkali Metals, Transition Metals and relative reactivity & extraction of metals

Lower level QUIZ on Group 1 Alkali Metals, Transition Metals and relative reactivity & extraction of metals

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