Download N4 Metals Topic 10

Topic 10 National 5 Chemistry Summary Notes
Notesnotes
Metals
Reactivity of Metals
LI 1
The reactivity series, which is shown in the table below is a list of metals
in order of their reactivity. The most reactive metals are at the top and
the least reactive are at the bottom.
Metal
Potassium
Reactivity with
Oxygen
react with
oxygen
Reactivity with
Water
react with water
Reactivity with
Acid
too reactive to
try in acid
Sodium
Calcium
Magnesium
react with acid
Aluminium
do not react
with water
Zinc
Iron
Nickel
Tin
Lead
Copper
Mercury
only forms metal
oxide on the
surface of the
metal
do not react
with acid
Silver
Gold
do not react
with oxygen
An easy way to remember the reactivity series is the following sentence:
Police Sergeant Charlie MAZINTL Caught Me Stealing Gold
The general word equations for metals reacting with oxygen, water or
acid are given below.
1.Metal + Oxygen
metal
+
e.g. magnesium
+
oxygen
metal oxide
oxygen
magnesium oxide
2.Metal + Water
metal
+
e.g. potassium
water
metal hydroxide
+ water
+
hydrogen
potassium hydroxide +
hydrogen
3.MAZINTL Metal + Acid
MAZINTL
metal
e.g. zinc
+
acid
salt
+ hydrochloric
acid
zinc chloride
2
+
hydrogen
+
hydrogen
Ionic Equations
LI 2
An ionic equation is an equation which shows any ions that may be present
among the reactants and products.
If you are asked to write an ionic equation for a reaction then you must
remember that not all of the substances in the reaction will be ionic.
When writing ionic equations remember the following points:

If a substance is not ionic then its formula will be no different
than usual.

If an ionic substance is present in the solid form then its’ ionic
formula is written in the usual way but with the state symbol (s)
placed after it.

If an ionic substance is dissolved in water then the ions are
separated in a special way and the state symbol (aq) is placed after
each ion.

Acids are ionic substances and should be shown with their ions
separated.
The following balanced ionic equations are for the reactions mentioned on
the previous page.
1. Mg(s) + O2(g)
 MgO(s)
equation
2Mg + O2
 2MgO
balanced equation
2Mg(s) + O2(g)

2Mg2+ O2-(s)
2. K(s)
+ H2O(l)

KOH(aq) + H2(g)
2K
+ 2H2O

2KOH
2K(s) + 2H2O(l)
3. Zn(s) + HCl(aq)
Zn
+
2HCl
balanced ionic equation
+ H2
 2K+(aq) + 2OH-(aq) + H2(g)

ZnCl2(aq) + H2(g)

ZnCl2
+ H2
Zn(s) + 2H+(aq) + 2Cl-(aq)  Zn2+(aq) + 2Cl-(aq) + H2(g)
3
LI 3
Extracting Metals
Less reactive metals can be found uncombined (not joined up with other
elements) in the earth’s crust and consequently were the first to be
discovered.
More reactive metals are always found combined and have to be
extracted (obtained) from ores. (see * below)
Metals have to be extracted from their ores by different methods. The
method used is shown in the table below and depends on the reactivity of
the metal.
Metal
Potassium
Extraction Method
electrical energy
required
Sodium
Calcium
Magnesium
i.e. electrolysis is the
splitting up of an ionic
compound into its original
elements using electricity
Extracting a metal from
its ore is an example of a
REDUCTION REACTION.
Aluminium
Zinc
Iron
*An ORE is a compound of
a metal that occurs
naturally. For example, iron
oxide is iron ore.
heat with carbon or
carbon monoxide
Nickel
Tin
Example
To extract silver metal from
silver(I)oxide it only has to be
heated. Write the balanced
ionic equation for this reaction.
Lead
Ag2O(s) 
Ag(s) + O2(g)
Copper
2Ag2O(s) 
4Ag(s) + O2(g)
Mercury
Silver
The balanced ionic equation is:
heat alone
2(Ag+)2O2-(s)  4Ag(s) + O2(g)
Gold
4
The more reactive metals hold on more strongly to oxygen than the less
reactive metals. Therefore, it is much easier to remove oxygen from
compounds where it is joined to less reactive metals.
The most reactive metals hold on to oxygen more strongly than carbon
does. Heating with carbon or carbon monoxide therefore does not work.
Wars and the invention of electricity led to the large scale extraction of
more reactive metals.
Very reactive metals are extracted from their ores using huge amounts
of electrical energy.
We can carry out the electrolysis of copper(II) chloride in the lab as
shown below:




DC current is used to ensure one electrode remains positively
charged and the other negatively charged.
Negative non-metal ions are attracted to the positive electrode
and positive metal ions are attracted to the negative electrode.
Graphite electrodes are used since they conduct electricity and will
not react with the solution being electrolysed.
During electrolysis, chemical reactions take place at each
electrode.
5
Percentage Composition
LI 4
The percentage composition is the percentage by mass of each element in
a compound. To work out the percentage composition, follow the steps
given in the example below.
Example
What is the percentage composition of iron (III) oxide?
1. Formula
Fe2O3
2. Formula Mass
160
3. % of elements
i.e.
mass of element present
formula mass of compound
% of Fe
=
(2x56) x
160
% of O
=
(3x16)
160
x 100%
100% = 70%
x 100% = 30%
Note: Calculate the percentage by mass of iron and oxygen in
iron (III) oxide - this is asking the same question as the example above.
6
LI 5
Redox Reactions
Redox reactions are reactions where REDuction and OXidation take place.
The word OILRIG is useful when thinking about oxidation and reduction
reactions.
Oxidation
Is
Loss
Reduction
Is
Gain
(of electrons)
(of electrons)
p.10 of the data booklet gives reduction equations. Remember, just
reverse to get the oxidation equation. These equations are commonly
called the ION-ELECTRON EQUATIONS (also known as half reactions
or half equations or ion-electron half equations).
Metals higher up the table on p.10 of the data booklet undergo oxidation
reactions, whereas, metals lower down undergo reduction reactions.
Displacement Reactions – A Type of Redox Reaction
A displacement reaction is a reaction which occurs when a metal higher up
The Electrochemical Series is added to a solution containing ions of a
metal lower down in the Electrochemical Series. For example,
iron (grey)
+
copper sulfate
solution (blue)
copper (red/brown)
+
iron sulfate
solution(colourless)

The copper and iron have changed places!
7
In this reaction the following has happened:
Iron atoms give electrons to copper ions i.e.
Cu2+(aq)
+
2e-
Cu(s)
REDuction
The copper ions are reduced to copper atoms which appear as a
red/brown solid.
Fe2+(aq)
Fe(s)
2e-
+
OXidation
The iron atoms are oxidised to iron ions which dissolve into solution
forming iron sulphate.
The same happens in all displacement reactions i.e. the metal higher up in
the electrochemical series always loses electrons and forms ions, and the
metal lower down always gains these electrons and forms atoms.
Rule : A metal higher up in the electrochemical series always displaces a
metal lower down.
Note: all displacement reactions are redox reactions.
To get the redox equation for the previous displacement reaction,
combine the ion-electron equations i.e.
Cu2+(aq) +
Fe(s)

2e-

Fe2+(aq) +
Cu(s)
REDuction
2e-
OXidation
Cu2+(aq) + 2e- + Fe(s)  Cu(s) + Fe2+(aq) + 2e-
Cu2+(aq) + Fe(s)
 Cu(s)
+ Fe2+(aq)
8
add and cancel
overall redox equation
LI 6
Cells/Batteries
Note: whenever you see the word CELL in these notes it can be replaced
with the word BATTERY.
A CELL is an arrangement which converts chemical energy into electrical
energy (electricity).
Electricity can be produced by connecting different metals together and
dipping them in an electrolyte (see note below) to form a cell.
Example – The Zinc/Copper Cell
.
An ELECTROLYTE is a liquid or solution which conducts.
The purpose of the electrolyte is to COMPLETE THE CIRCUIT

Acids and ammonium chloride solution are examples of electrolytes.
9
LI 7
The Electrochemical Series
We can use the equipment shown below to compare the voltage produced
by different pairs of metals. The two metals are connected by an
electrolyte.
The results obtained are given in the table below.
Metal Pair
Voltage Reading
(millivolts)
copper and copper
0
copper and tin
10
copper and iron
40
copper and zinc
50
copper and magnesium
60
copper and silver
-10
These results show that different pairs of metals give different voltages
and this leads to THE ELECTROCHEMICAL SERIES which is shown on
p.10 of the data booklet.
The electrochemical series places metals in order of their ability to
supply electrons (it is very similar to the reactivity series but not exactly
the same) The metals at the top of the series supply electrons most
easily.
Electrons always flow from the metal higher up the electrochemical
series to the metal lower down.
The further apart the metals are in the electrochemical series, the
10
higher the voltage they produce.
LI 8

Oxidising and Reducing Agents
Oxidising agents cause other species to be oxidised and are
therefore themselves reduced.

Reducing agents cause other species to be reduced and are
therefore themselves oxidised.
Example 1
Using the equations below, circle the oxidising agent with a dotted circle
and the reducing agent with a full circle.
Ag+(aq)
Cu(s)
e-
+
Cu2+(aq)


Ag(s)
+
2e-
Example 2
Fe2O3
+
3CO
2Fe
+
3CO2
The above reaction shows the final reaction in the production of iron
from iron ore. This takes place in industry in a blast furnace.
It shows that the iron ore (Fe2O3) is reduced to iron when it reacts with
the carbon monoxide. Therefore the carbon monoxide is the reducing
agent.
2Fe3+
+ 3e-
2Fe
reduction
11
LI 9
More Complicated Cells - Half Cells
The Zinc/Copper Cell Again!
In the above set-up:

At the zinc rod the reaction taking place is:
Zn2+(aq)
Zn(s) 
2e-
+
(oxidation)
The zinc rod is getting LIGHTER as its atoms turn into ions which then enter the
solution.

At the copper rod the reaction taking place is:
Cu2+(aq)
+
2e-

Cu(s)
(reduction)
The copper ions are gaining electrons to become copper atoms which sink into the
copper rod, making it HEAVIER.
As before, combining these two equations gives the redox equation for the overall
cell reaction.
Zn +
Zn(s) +
Cu2+ + 2e-

Zn2+
Cu2+(aq)

Zn2+(aq)
+ 2e + Cu
+
Cu(s)
12
add and cancel
redox equation
Electrons flow from the zinc rod to the copper rod through the wires and
the meter.
ELECTRONS always flow through the wires and the meter.
The purpose of the ION BRIDGE is to complete the circuit – it is
the movement of ions in the ion bridge which completes the circuit.
Ions flow through the
ion bridge.
Electrons flow through wires.
Electrons always flow through the wires and meter from the metal higher up
the electrochemical series to the metal lower down.
When setting up a cell like the zinc/copper cell, for electricity to be
produced the metals have to be:
1. different
2. placed in a solution of their own metal ions. For example, zinc has
to be placed in a zinc solution e.g. zinc chloride, it cannot be placed
in a copper solution such as copper chloride.
13
Cells with Non-Metals
The half-cells in a cell need not involve metal atoms.
graphite
electrodes
solution containing
sulfite ions (SO32-)
ion
bridge
iodine
solution (I2)
In the above set-up the following show the ion-electron half equations
involved:
Oxidation
SO32-(aq)
+
H2O(l)

SO42-(aq)
+
2H+(aq)
+
2e-
Reduction
I2(aq) +
2e-

2I-(aq)
Combining these two equations gives the redox equation for the overall
cell reaction.
SO32-(aq) + H2O
(l)
+ I2(aq)
+ 2e  SO42-(aq) + 2H+(aq) + 2e + 2I-(aq)
add and cancel
SO32-(aq) + H2O(l) + I2(aq)  SO42-(aq) + 2H+ (aq) + 2I-(aq)
redox equation
14
LI 10
More on Redox Reactions
Fuel cells and rechargeable batteries are two examples of technologies
which make use of redox reactions.
Fuel Cells
A fuel cell is a device that converts the chemical energy from a fuel into
electricity through a chemical reaction with oxygen or other oxidising
agents.
Hydrogen is the most common fuel used and these fuel cells are called
hydrogen fuel cells.
The ion-electron equations and overall redox equation for a hydrogen fuel
cell are shown below.
H2(g)  2H+(aq) + 2e
+
O2(g) + 4H
(aq)
+ 4e  2H2O(l)
(x2)
oxidation
(leave)
reduction
2H2(g)  4H+(aq) + 4e
oxidation
+
reduction
O2(g) + 4H
(aq)
+ 4e  2H2O(l)
2H2(g) + O2(g) + 4H+(aq) + 4e 4H+(aq) + 4e + 2H2O(l)
add and cancel
2H2(g)
redox
+ O2(g)  2H2O(l)
As can be seen from this redox equation, using fuel cells helps reduce
carbon dioxide emissions.
Fuel cells are increasingly being used in place of internal combustion
engines for transport.
15
Rechargeable Batteries
Rechargeable batteries are batteries which can be made to work again
when they go flat by charging. Today, many items we use on a daily basis,
for example mobile phones, are powered by rechargeable batteries. The
lead-acid battery is the oldest type of rechargeable battery and it is still
used today to start car engines.
The ion-electron equations and overall redox equation for this type of
battery whilst it is recharging are shown below.
PbSO4(s) + 2H2O(l)  PbO2(s) + 4H+(aq) + SO42-(aq) + 2e
oxidation
PbSO4(s) + 2e  Pb(s) + SO42-(aq)
reduction
2PbSO4(s) + 2H2O(l)  Pb(s) + PbO2(s) + 4H+(aq) + 2SO42-(aq)
redox
When a battery is being recharged the energy change is:
electrical
chemical
16
Topic 10 Pupil Self Evaluation
Metals – National 5
If there is an E in any part of the notes or the success criteria is in italics, then this is National 5 level work.
Number
Learning Intention
Success Criteria
1
I am going to find out about:
the reactivity of metals
2
ionic equations
3
extracting metals from their ores
I can:
 state the order of metals in The Reactivity Series
 state if a metal reacts with oxygen, water or acid and write the
word equation for the reaction
 state the balanced ionic equation for a metal reacting with
oxygen, water or acid4
 state the definition of an ore
 state the method of extraction required to extract a particular
metal from its ore
 explain why this method of extraction is required
 state the balanced ionic equation for the extraction of a
particular metal from its’ ore


4
percentage composition



explain why unreactive metals were the first to be discovered
give examples of what led to the large scale extraction of the
more reactive metals
state which metal is produced in the blast furnace
state the reactions which take place in the blast furnace
work out the percentage of a particular element in a compound
5
Redox Reactions





6
cells
7
The Electrochemical Series






8
oxidising and reducing agents
9
more complicated cells



state the definition of a redox and a displacement reaction
state the definition of the terms oxidation and reduction
state the oxidation and reduction reactions for a given reaction
explain what happens in a displacement reaction stating the
oxidation and reduction reactions involved
work out the redox equation for a redox reaction
describe how electricity can be produced using metals
draw a set-up of how electricity can be produced using metals
state the purpose of an electrolyte
explain The Electrochemical Series
state the direction of electron flow if two different metals are
connected in a cell
the size of voltage produced to the position of metals in The
Electrochemical Series
state the definition of an oxidising and reducing agent
given a balanced equation state the reducing agent
draw a set-up of how electricity can be produced using metals
and solutions of their own ions
explain what happens in these set-ups stating the oxidation and
reduction reactions involved
 work out the redox equation for these set-ups
 state ion-electron equations and work out the redox equation for
cells with non-metals

1
10
two examples of technologies which
make use of redox reactions







state where electrons flow in this set-up
state where ions flow in this set-up
state the purpose of an ion bridge
state the definition of a fuel cell and a rechargeable battery
work out the redox equation given the ion-electron equations
involved in the reaction in a fuel cell
state the effect the use of fuel cells has on carbon dioxide
emissions
work out the redox equation given the ion-electron equations
involved in the reaction in a rechargeable battery
2
Points to Note






a rough draft, I still need to read over it myself!
Haven’t given a general blurb on metals as I thought the teacher
would set the scene.
Nat 5 - Metallic bonding covered in Topic 3 p.16 therefore not
added to these notes.
Nat 5 Support Notes LHS of Table ‘balanced ionic
equations…reduction reactions’ - not sure what they are looking for
here, may have covered it in what i’ve put together already?Let me
know what you think?
Nat 5 page numbers help….can’t get them sorted!
Cells stuff – really needs a check!
Topic 9 – Metals and Alloys Experiments – Nat 5
Note: The experiments listed below the dotted line are optional as they
are National 4 experiments
1. Redox Reactions – SGrade Topic 11 - Displacement Reactions
2. Cells with Non-metals – SGrade Topic 10 – Demo of SO42-/I2 set up
3. More on Redox Reactions – a car battery
--------------------------------------------------------------------------4. Materials - a selection of different materials - ??? – the ones
mentioned in the notes????
5. Reactivity of Metals


Alkali Metal demo
SGrade Topic 11 – Metals & Water/Acid/Oxygen - could test
for hydrogen if released.
6. Extraction of Metals - ??????????? do we have anything?
7. Corrosion – SGrade Topic 12 – nails expt in water etc…
8. Rusting – ferroxyl indicator and Fe2+ ion and OH- ion solutions
9. Preventing Corrosion – iron/magnesium cell set up in a u-tube with
salt water and ferroxyl indicator
10. Cells/Batteries






zinc rod
copper rod
dilute sulphuric acid
voltmeter
wires
lemon, wires, voltmeter, zinc & copper rods ?????????
11. The Electrochemical Series – SGrade Topic 10 Electrode Potential
12. The Zinc/Copper Cell Again!








zinc rod
copper rod
voltmeter
wires
filter paper
salt solution
zinc chloride solution
copper sulphate solution
13. Alloys


circuits boards
a selection of different alloys - ??? – the ones mentioned in
the notes???
1