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
The Reactivity Series of Metals
(reactivity of francium and compared with the non-metals
carbon and hydrogen) |
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

Group 1 Alkali Metal
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.
-
Because it is extremely reactive, it
reacts and explodes violently with cold water forming the alkali caesium hydroxide and
flammable-explosive
hydrogen gas.
==>
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

Group 1 Alkali Metal
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 was
first extracted in 1861 by electrolysis of the molten chloride RbCl
-
GCSE/IGCSE/O
level revision study notes on
Group 1
The Alkali Metals
-
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

Group 1 Alkali Metal
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).
-
Potassium was
first extracted in 1807 by electrolysis of the molten chloride KCl
-
GCSE/IGCSE/O level revision study notes on
Group 1 The Alkali Metals
-
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

Group 1 Alkali Metal
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 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 was
first extracted in 1807 by electrolysis of the molten chloride
NaCl. Extraction
of sodium
|
lithium
Li

Group 1 Alkali Metal
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
-
GCSE/IGCSE/O level revision study notes on
Group 1 The Alkali Metals
-
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

Group 2 Alkaline Earth Metal
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 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
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

Group 2 Alkaline Earth Metal
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 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.
-
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
TOP OF PAGE and
sub-index
|
aluminium
Al

Group 3 Metal
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)
|
Elements higher than carbon i.e. aluminium or more reactive, must be extracted by electrolysis
(or displacing it with an even more reactive metal).
Metals below it, i.e.
zinc
or a less reactive, can be extracted by reducing the hot metal oxide with carbon. |
zinc
Zn

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

In the 1st block-series of Transition Metals
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

A Group 4 metal
The Reactivity Series of Metals
(reactivity of tin and compared with the non-metals carbon
and hydrogen)
|
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

A Group 4 metal
|
-
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 |
Non of the metals below hydrogen can react with acids to form hydrogen
gas.
They are least easily corroded metals and partly accounts for their value and uses in jewellery,
electrical contacts etc.
|
copper
Cu

In the 1st block-series of Transition Metals
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 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

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

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

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.
|
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(d) OTHER ASSOCIATED PAGE
LINKS
SEE ALSO
2. RUSTING &
Introducing REDOX reactions
and 3.
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
GCSE/IGCSE
reactivity gap-fill worksheet or
Rusting
word-fill worksheet
KS4 Science GCSE/IGCSE/O level
Chemistry revision notes
pages:
The Periodic Table
*
Group 1 Alkali Metals
*
Methods of Metal
extraction
Transition Metals
*
Alloys-uses of metals
*
Electrochemistry-Electrolysis
Rates of Reactions Experiments (e.g.
metal-acid)
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