Core & Extension Metals I Reactivity Series & Redox.pptx

© 2004-09 Dorje Gurung
Core & Extension Metals I:
Reactivity Series & Redox

Core & Ext. Metals: Reactivity Series & Redox
Slide 2 of 53
Learning Objectives
Concepts:
– Reactivity, displacement, reduction, metal, order of reactivity
– redox reaction, reduction, oxidation, oxidizing agent (oxidant), reducing
agent (reductant), half-equation, ionic equation
Skills
Core:
– Place the following metals in order of reactivity: calcium, copper,
(hydrogen), iron, magnesium, potassium, sodium and zinc by reference
to the reactions, if any, of the metals with water or steam, dilute
hydrochloric acid and the reaction of their oxides with carbon
– Deduce an order of reactivity from a given set of experimental results
– Place carbon and hydrogen in the reactivity series.

Slide 3 of 53
Learning Objectives
Skills
– Define oxidation and reduction in terms of oxygen loss/gain
– Identify reduction, oxidation, reducing agent, oxidizing agent
– Describe methods of rust prevention: paint and other coatings to exclude
oxygen
Extension:
– Account for the apparent unreactivity of aluminum in terms of the oxide
layer which adheres to the metal

Slide 4 of 53
Metals
80% of all elements are metals.
Fig 1. Metals & non-metals in the periodic table
Other common metals include
iron, calcium and magnesium.
Fig 2. The amount of different metals in the earths surface
The most abundant metal in the earths
surface is aluminum.

Slide 5 of 53
The periodic Table

Slide 6 of 53
Some General Properties
Physical Properties
They are usually shiny. However they may look
dull if they they are covered in a layer of
metal oxide.
All metals will conduct electricity.
All metals are good conductors of heat
Most metals are malleable: they can be beaten
into different shapes
Many metals like copper are ductile: they can
be stretched into thin wires
Most metals have high melting points and
boiling points
Chemical Properties
They lose electrons to from
positively charged ions
They combine with non-metals to
form ionic compounds
Most metals react with acids to
produce a salt and hydrogen
They react with oxygen to produce
basic oxides

Slide 7 of 53
Reactivity Series
Remember, when metals react, they lose electrons to become positively
charged ions.
Some metals lose electrons easily, this makes them very reactive.
– For example, potassium reacts explosively with cold water and acids.
Other metals react more slowly.
– A piece of iron takes several days to rust when it is placed in water.
A few metals, such as gold and silver are very unreactive, they do not
combine with other elements, they never rust or oxidize.
The Group One metals are most reactive, followed by the Group Two
metals and then the transition metals.
– The Metals at the bottom of each group are more reactive than the
metals on the top of each group. (Do you remember why?)

Slide 8 of 53
Reactive
1. What does the word reactive actually mean?
It means how easily a chemical combines with another to form a
compound.
In the case of a metal, it’s how easily it will lose electrons to form a
cation (or the difficulty with which its ion will part with the anion to
which it is bonded).
2. Why is potassium very reactive?
It is very reactive because the outer electron is very easily removed.

Slide 9 of 53
Reactivity Series of metals
From your Form 3 Chemistry studies of patterns of reactivity, you should
know that the reactivity series is as follows:
Most reactive  Least reactive
Aluminium appears in the middle of the reactivity series, between magnesium
and zinc.
This suggests that it should corrode faster than iron, and react more quickly
with acids than zinc.
But aluminium does not corrode and does not react with acids as you have
already noted from past experiments!

Slide 10 of 53
Reactivity of metals: Aluminum
Why does aluminum appear to be very unreactive?
The surface of a piece of aluminum oxidizes very quickly and becomes covered
with a thin layer of aluminum oxide.
The layer of aluminum oxide has the same effect as a coat of paint , protecting
the aluminum beneath it.
Other metals such as iron also oxide, but iron oxide is soft easily removed from
the surface so it cannot protect the metal .

Slide 11 of 53
Reactivity Series
The order was arrived at by observing the manner and extent of
reaction of the metal
– with water (or steam),
– with dilute acid,
– oxide with carbon,
– oxide with other metals,
– and with salt solution of other metals
This year, we will also look at the stability of hydroxides and nitrates of
the metals to heat and the metal’s reaction with oxides of other
metals.
Most reactive  Least reactive

Slide 12 of 53
Reaction of Metals with Water
Most metals react with water; just that some react vigorously
with even cold water, while others require steam.
But the reaction always produces hydrogen gas.
Very reactive metals like potassium will react quickly with
cold water, producing a solution of potassium
hydroxide and hydrogen gas.
K (s) + H2O (l)  KOH (aq) + H2 (g)
– The heat of the reaction is great enough to ignite the
hydrogen explosively.
– Sodium may also ignite as shown to the right.
• The reaction is more and more violent down the group as you
have already noted from your studies of the periodic table.
• So the three placed in order of decreasing reactivity would be K
> Na > Li.

Slide 13 of 53
Less reactive metals like magnesium and calcium react very slowly with
cold water (fig.1), but quickly with steam (fig.2). Solid white
magnesium oxide powder and hydrogen gas and are produced.
Fig. 2
Fig. 1
hydrogen

Slide 14 of 53
When calcium react with water, the products are the same as when
group 1 metals react wit it.
Ca (s) + H2O (l)  Ca(OH)2 (aq) + H2 (g)
– The resulting solution of course is limewater, which when reacted with carbon
dioxide produces the white precipitate of calcium carbonate.
Magnesium however does not react readily with water.
It reacts instead with steam to produce magnesium oxide and hydrogen
gas.
Magnesium + steam  Magnesium oxide + hydrogen
Mg (s) + H2O (g)  MgO (s) + H2 (g)
Unreactive metals such as gold do not react with water at all.
That is why gold coins recovered from old ship wrecks appear shiny and
new.

Slide 15 of 53
Since calcium reacts with water while magnesium requires steam, order
of reactivity is therefore Ca > Mg.
In other words, reactivity increases going down group 2 as well just as
with group 1 metals.
Placing K, Na, Ca and Mg in order we get: K > Na > Ca > Mg.

Slide 16 of 53
Reactions of metals with acid
Many metals react with dilute acids to produce their respective salt and
hydrogen gas.
The vigor of the reaction varies with the reactivity of the metal.
The more reactive the metal the more vigorous the reactivity—bubbling
and heat is produced when a reaction does take place.
In general:
Metal + acid  salt solution + hydrogen gas
Mg (s) + HCl (aq)  MgCl2 (aq) + H2 (g)
Ionic equation:
Metal + 2H+
(aq)  Metal ion (aq) + H2 (g)
Mg(s) + 2H+
(aq)  Mg2+
(aq) + H2 (g)

Slide 17 of 53
Reactions of metals with acid
The most reactive metals such as potassium, sodium and lithium react
very rapidly, producing enough heat to ignite the hydrogen gas and
cause an explosion.
– Group 1 metals therefore are not tested with acids in a secondary lab.
Less reactive metals such as iron react quite slowly with acids, but the
reaction can be made faster by warming the acid gently with a
bunsen burner.
Copper reacts with only concentrated acids and, besides, the reaction is a
little different from those of the reactive metals.
Gold reacts only with a special mixture of concentrated nitric acid and
sulfuric acid.

Slide 18 of 53
Reactivity series
Therefore, group 1 metals are more reactive than Magnesium, Zinc, Iron
and Copper.
Of the metals tested, Magnesium was the most vigorous, zinc next, and
then iron while copper did not even react.
So therefore, of the four metals the most reactive is Magnesium,
followed by Zinc, and then Iron, and lastly Copper.
The reactivity series from the most to the least reactive looks like this:
Potassium, Sodium, Lithium, Calcium, Magnesium, Zinc, Iron, Copper.

Slide 19 of 53
Placement of carbon in the order
The experiments on the reactions of charcoal with metal oxides you
performed allowed you to determine the approximate location of
carbon (charcoal) in the order of reactivity.

Slide 20 of 53
Carbon was able to displace lead, zinc and copper from their oxides but
not aluminum oxide in the experiments you performed in the
laboratory.
zinc oxide + Carbon  Carbon dioxide + zinc
ZnO(s) + C(s)  CO2(g) + Pb(s)
Copper oxide + Carbon  Carbon dioxide + copper
CuO(s) + C(s)  CO2(g) + Cu(s)
Aluminum oxide + Carbon  No reaction

Slide 21 of 53
Based on these results, it was concluded
that carbon could be placed between
aluminum and zinc.
Displacement reactions where oxygen
exchange takes place between the
reactants is classified under a whole set
of reaction type called redox reaction.
And that is what we are going to look at
next.
Potassium
Sodium
Calcium
Magnesium
Aluminum
Zinc
Iron
Copper
Here

Slide 22 of 53
Redox reactions and oxygen
A redox reaction is a type of chemical reaction in which oxidation
and reduction occur.
When iron rusts, it reacts with oxygen in air to form iron oxide,
commonly called rust.
Fe (s) + O2 (g)  Fe2O3 (s)
Rusting of iron is an example of a redox reaction.
When material made from iron (for
example car body, nails, axes, railings
etc) is exposed to the elements, iron
rusts.

Slide 23 of 53
Redox in terms of oxygen
When other metals notably group 1 elements are exposed to air they
react with oxygen in air and undergoes a redox reaction.
Their shiny recently-exposed surface goes grey over time (which you
oberved in the lab).
Li (s) + O2 (g)  Li2O (s)
Na (s) + O2 (g)  Na2O (s)
K (s) + O2 (g)  K2O (s)
Aluminum also reacts with air forming an thin layer of oxide which
prevents the rest of the aluminum from corroding.
Al (s) + O2 (g)  Al2O3 (s)

Slide 24 of 53
Redox in terms of oxygen
Similarly, when coal is burnt to provide fuel,
then it reacts with oxygen in air and
undergoes a redox reaction.
C (s) + O2 (g)  CO2 (g) Carbon has undergone
oxidation; it has gained
oxygen.
The oxygen added to an element can be obtained from another
compound.

Slide 25 of 53
The Reactivity Series
This happens when a metal that
is higher up in the reactivity
series reacts with the oxide of
a less reactive element in the
series, such as another metal
or carbon or hydrogen
(water).
metals becoming
more reactive
Potassium
Sodium
Calcium
Magnesium
Aluminium
(Carbon)
Zinc
Iron
(Hydrogen)
Copper
Silver
Gold

Slide 26 of 53
Oxygen from an oxide
From your studies of metal reactivity, you know that a more reactive
metal (or substance) will displace a less reactive metal from its
oxide.
Metal 1 (or substance 1) + Metal 2 oxide  Metal 1 oxide + Metal2
Al has undergone oxidation, or
has been oxidized (to aluminum
oxide).
Oxidation is gain of oxygen.
Iron(III) oxide is said to have
undergone reduction, or it’s been
reduced (to iron).
Reduction is the loss of oxygen.
For example, the thermite reaction
Al(s) + Fe2O3(s)  Fe(s) + Al2O3(s)
(s) 2 3(s) (s) 2 3(s)
Al + Fe O Fe + Al O

Oxidation
Reduction

Slide 27 of 53
Again, since Fe2O3 is the provider of oxygen (causing Al to get oxidized),
it is referred to as the Oxidizing agent.
Similarly, since Al removes oxygen from Fe2O3 (causing it to be reduced
to Fe), Al is referred to as a the Reducing agent.
Similarly,
(s) 2 3(s) (s) 2 3(s)
Al + Fe O Fe + Al O

(s) (s) (s) (s)
Mg + CuO MgO + Cu

Oxidation
Reduction
Reducing agent
Oxidizing agent

Slide 28 of 53
Similarly,
2CuO(s) + C(s)  2Cu(s) + CO2(g)
2Fe2O3(s) + 3C(s)  4Fe(s) + 3CO2(g)
2PbO(s) + C(s)  2Pb(s) + CO2(g)
H2(g) + CuO (s)  Cu(s) + H2O (s)
Here CuO has undergone reduction, and H2 oxidation. CuO loses
oxygen, while H2 gains oxygen.
Identify the oxidizing and reducing agents in the reactions that appear on
this slide.
In the last reaction, PbO has
undergone reduction, and C
oxidation.
PbO loses oxygen, while C gains
oxygen.

Slide 29 of 53
Oxygen from Water
If we take water for example, any it can provide oxygen to any metal
higher up in the reactivity series than hydrogen.
Mg has undergone
oxidation, and water is
said to have undergone
reduction. Reduction is
the loss of oxygen.
When steam is passed over magnesium,
magnesium oxide and hydrogen gas is
produced for example.
Mg (s) + H2O (g)  MgO (s) + H2 (g)
Aluminum would react in a similar manner.
What is oxidizing agent and what the reducing agent in the above
reaction?
Calcium and the group 1 metals, as you already know, react with water
in a completely different manner.

Slide 30 of 53
Redox reaction definition
In terms of oxygen:
– Gain of oxygen is oxidation.
• That which gains (or accepts) oxygen is the reducing agent.
– Loss of oxygen is reduction.
• That which loses (or provides) oxygen is the oxidizing agent.

Slide 31 of 53
Redox reaction
Photography, respiration, reactions in batteries, wine going off, bleaching,
rusting etc. are all examples of what in chemistry we call redox
reactions.
Photography involves conversion of silver bromide to silver metal.
Respiration involves the breakdown of glucose in the presence of oxygen
to get energy.
Reactions in batteries involves using the different reactivities of metals to
make electrons flow through a circuit which again provides energy.
Wine going off, bleaching, rusting also involve reaction with oxygen
which are classified as redox reaction.

Slide 32 of 53
Another example of Redox Reaction:
Rusting
After:
A new material
(Fe2O3) is made
which is brown
and flaky.
Before:
The iron metal
(Fe) is shiny
and hard.
We say that iron and has undergone oxidation.
The definition of oxidation according to this
example given here is the addition of oxygen.

Slide 33 of 53
Prevention of Rusting
Rusting of iron and steel can be prevented by any method that stops
iron from coming in contact with oxygen.
One of the ways to do so is to paint
it.
Paint essentially forms a barrier
between oxygen in the air and
the iron.
Other methods which also work under the same principle involves coati
it with oil and grease which work particular well with moving parts,
and plastic (PVC) coating, used on items such as refrigerators and
lawn chairs.
Still others involve electroplating (with layer of chromium or tin) and
galvanizing (usually with zinc) (see Electricity & Chemistry:
Electroplating).

Slide 34 of 53
Summary
The reactivity series lists metals (and two nonmetals) in the order of their
relative reactivity.
Displacement reactions provide evidence for position of a metal and
those of hydrogen and carbon.
The displacement reaction may be
– Between a metal and salt solution of another
– Between a metal and the oxide of another metal or that of hydrogen or
carbon,
– Reaction with acid.

Slide 35 of 53
Summary
The more reactive the metal that higher up in the reactivity series it will
appear.
The more reactive metal/non-metal will displace a less reactive
metal/nonmetal from its salt solution or its oxide.
When oxygen exchange takes place between substances, the reaction
can be classified as a redox reaction.
The one that accepts the oxygen is said to have undergone oxidation and
therefore referred to as the reducing agent.
The one that loses (provides) the oxygen is said to have undergone
reduction and therefore referred to as the oxidizing agent.
Aluminum appears to be unreactive due to the protective layer of oxide
that coats its surface, which also prevents it from rusting further.

Slide 36 of 53
Paper 2 Practice Questions
1. J05/2/6d. The information below shows the reactivity of chromium,
copper and iron with warm hydrochloric acid.
chromium – few bubbles of gas produced every second
copper – no bubbles of gas produced
iron – many bubbles of gas produced every second
Put these three metals in order of their reactivity with hydrochloric
acid.
Most reactive  ………………………….
………………………….
Least Reactive  …………………………. [1]

Slide 37 of 53
2. N04/2/1e. State three properties shown by all metals. [3]
3. J03/2/6 (c) Uranium is between magnesium and zinc in the reactivity
series. Equal sized strips of magnesium, uranium and zinc were placed
in hydrochloric acid. The hydrochloric acid was the same
concentration. The results are shown in the table.

Slide 38 of 53
(i) Complete the result for uranium and hydrochloric acid. [1]
metal observations on adding to hydrochloric acid
magnesium
many bubbles of gas produced very rapidly and magnesium
dissolves quickly
uranium
zinc
a few bubbles produced at a steady rate and zinc dissolves
slowly

Slide 39 of 53
4. N01/2/6. In the ‘thermit’ reaction, aluminium powder reacts violently
with iron(III) oxide. A magnesium ribbon is lit to start the reaction.
The reaction gives out a great deal of heat.
The equation for the reaction is:
2Al(s) + Fe2O3(s)  Al2O3(s) + 2Fe(s)

Slide 40 of 53
(a) Complete the following sentence about the ‘thermit’ reaction using
words from the list.
added
electrolysed
neutralised
oxidised
reduced
In the ‘thermit’ reaction, the aluminium is ................................... to
aluminium oxide and the iron(III) oxide is ................................... to
iron. [2]
(b) Aluminium oxide which has been heated to a high temperature is called
fused aluminium oxide. Fused aluminium oxide does not react with
hydrochloric acid. Iron reacts with hydrochloric acid.

Slide 41 of 53
(i) What would you observe when hydrochloric acid is added to a mixture of
fused aluminium oxide and iron? [1]
(ii) After reaction with hydrochloric acid, the mixture contains fused aluminium
oxide solid and a solution of iron(II) chloride. Describe with the help of a
labelled diagram, how you would separate the aluminium oxide from the
iron(II) chloride solution. [4]

Slide 42 of 53
(c) The magnesium ribbon used to start the ‘thermit’ reaction burns in
oxygen and gives out heat. What term describes a reaction that gives out
heat? [1]
(d) In an oxyacetylene torch, oxygen is used with acetylene to produce a
flame with a temperature of about 3000°C. State one use of this flame.
[1]

Slide 43 of 53
1. J05/3/6. The position of aluminium in the reactivity series of metals is
shown below.
Magnesium
Aluminium
Zinc
Copper
(b) Aluminium reacts very slowly with aqueous copper(II) sulphate.
2Al(s) + 3CuSO4(aq)  Al2(SO4)3(aq) + 3Cu(s)
(i) Which of the two metals has the greater tendency to form ions? [1]
(ii) Describe what you would see when this reaction occurs. [1]

Slide 44 of 53
(iii) Explain why aluminium reacts so slowly. [1]
(c) Complete the following table by writing “reaction” or “no reaction” in the
spaces provided. [2]

Slide 45 of 53
2. N02/3/5d. Steel may be coated with another metal, eg zinc or
chromium, or with a polymer, eg poly(chloroethene), to prevent
rusting.
(i) Suggest a property of poly(chloroethene) that makes it suitable for this
purpose. [1]
(ii) Explain why the steel will rust when the protective coating of chromium
or polymer is broken. [1]
(iii) When the protective layer of zinc is broken, the steel still does not rust.
Suggest an explanation. [2]

Slide 46 of 53
3. N01/3/4. (d) A use of zinc is galvanising. When the zinc layer is
broken, the steel is exposed.
Explain why the exposed steel does not rust. [3]

Slide 47 of 53
Paper 6 questions
1. J04/6/2. The four tubes show an investigation of rusting. Each one of
these four tubes contains a nail and the reagents indicated.

Slide 48 of 53
Paper 6 questions
(a) Predict the order in which rust would appear.
First
Second [1]
(b) Explain your prediction. [2]

Slide 49 of 53
Paper 6 questions
2. N02/6/4. An investigation was carried out on the reactions of four
different metals. Equal masses of copper, magnesium, iron and zinc
were used.
Experiment 1
A 15 cm3 sample of dilute sulphuric acid was added to each of four
boiling tubes. The initial temperature of the acid was measured. Zinc
was added to the first tube, iron to the second tube, magnesium to
the third tube and copper to the fourth tube.
The maximum temperature reached in each tube was measured and
any observations were recorded in the table.
(a) Use the thermometer diagrams to complete the results table.

Slide 50 of 53

Slide 51 of 53
Paper 6 questions
Use your results and observations to answer the following questions.
(i) Which metal is most reactive with sulphuric acid? [1]
(ii) Give two reasons why you chose this metal.
1. . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . .
2. . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . [2]
(iii) Name the gas given off. [1]
The reaction between magnesium and aqueous copper(II) sulphate
was then investigated.

Slide 52 of 53
Paper 6 questions
Experiment 2
A 5 cm3 sample of aqueous copper(II) sulphate was measured into a test-
tube. The initial temperature of the solution was measured.
Magnesium powder was added to the test-tube and the maximum
temperature reached was measured. Use the thermometer diagrams to
complete the results table.

Slide 53 of 53
Paper 6 questions
(b) How do your observations show that the reaction of magnesium with
aqueous copper(II) sulphate is exothermic? [1]
(c) What type of exothermic reaction occurs when magnesium is added to
aqueous copper(II) sulphate? [1]
(d) Use your results from Experiments 1 and 2 to put the four metals in
order of reactivity.
least reactive ................................
.................................
.................................
most reactive ................................. [1]

Core & Extension Metals I Reactivity Series & Redox.pptx

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