Metals

Reactivity of Metals

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 Reactivity with

Oxygen

Reactivity with

Water

Reactivity with

Acid

Potassium

Sodium

Calcium

Magnesium

Aluminium

Zinc

Iron

Nickel

Tin

Lead

Copper

Mercury

Silver

Gold

react with

oxygen

only forms metal

oxide on the

surface of the

metal

do not react

with oxygen

react with water

do not react

with water

too reactive to

try in acid

react with acid

do not react

with acid

An easy way to remember the reactivity series is the following sentence:

LI 1

Police Sergeant Charlie MAZINTL Caught Me Stealing Gold

Topic 10 National 5 Chemistry Summary Notes

Notesnotes

2

The general word equations for metals reacting with oxygen, water or

acid are given below.

1.Metal + Oxygen

e.g. magnesium + oxygen magnesium oxide

2.Metal + Water

e.g. potassium + water potassium hydroxide +

3.MAZINTL Metal + Acid

e.g. zinc + hydrochloric zinc chloride + hydrogen

acid

metal + oxygen metal oxide

metal + water metal hydroxide + hydrogen

hydrogen

MAZINTL + acid salt + hydrogen

metal

3

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) balanced ionic equation

2. K(s) + H2O(l) KOH(aq) + H2(g)

2K + 2H2O 2KOH + H2

2K(s) + 2H2O(l) 2K+(aq) +

2OH-(aq) + H2(g)

3. Zn(s) + HCl(aq) ZnCl2(aq) + H2(g)

Zn + 2HCl ZnCl2 + H2

Zn(s) + 2H+(aq) + 2Cl

-(aq) Zn2+

(aq) + 2Cl-(aq) + H2(g)

LI 2 Ionic Equations

4

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 Extraction Method

Potassium

Sodium

Calcium

Magnesium

Aluminium

Zinc

Iron

Nickel

Tin

Lead

Copper

Mercury

Silver

Gold

electrical energy

required

i.e. electrolysis is the

splitting up of an ionic

compound into its original

elements using electricity

heat with carbon or

carbon monoxide

heat alone

Extracting a metal from

its ore is an example of a

REDUCTION REACTION.

LI 3

*An ORE is a compound of

a metal that occurs

naturally. For example, iron

oxide is iron ore.

Example

To extract silver metal from

silver(I)oxide it only has to be

heated. Write the balanced

ionic equation for this reaction.

Ag2O(s) Ag(s) + O2(g)

2Ag2O(s) 4Ag(s) + O2(g)

The balanced ionic equation is:

2(Ag+)2O2-

(s) 4Ag(s) + O2(g)

5

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.

6

Percentage Composition

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 x 100%

formula mass of compound

% of Fe = (2x56) x 100% = 70%

160

% of O = (3x16) x 100% = 30%

160

Note: Calculate the percentage by mass of iron and oxygen in

iron (III) oxide - this is asking the same question as the example above.

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7

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 (of electrons)

Reduction

Is

Gain (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 (red/brown)

+ +

copper sulfate iron sulfate

solution (blue) solution(colourless)

The copper and iron have changed places!

LI 5

8

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.

Fe(s) Fe2+(aq) + 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) + 2e- Cu(s) REDuction

Fe(s) Fe2+(aq) + 2e- OXidation

Cu2+(aq) + 2e- + Fe(s) Cu(s) + Fe2+

(aq) + 2e- add and cancel

Cu2+(aq) + Fe(s) Cu(s) + Fe2+

(aq) overall redox equation

9

An ELECTROLYTE is a liquid or solution which conducts.

The purpose of the electrolyte is to COMPLETE THE CIRCUIT

Cells/Batteries

Note: whenever you see the word CELL in these notes it can be replaced

with the word BATTERY.

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

.

Acids and ammonium chloride solution are examples of electrolytes.

A CELL is an arrangement which converts chemical energy into electrical

energy (electricity).

LI 6

10

The further apart the metals are in the electrochemical series, the

higher the voltage they produce.

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.

LI 7

11

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) + e- Ag(s)

Cu(s) Cu2+(aq) + 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

LI 8

12

More Complicated Cells - Half Cells

The Zinc/Copper Cell Again!

LI 9

In the above set-up:

At the zinc rod the reaction taking place is:

Zn(s) Zn2+(aq) + 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 + Cu2+ + 2e- Zn2+ + 2e + Cu add and cancel

Zn(s) + Cu2+(aq) Zn2+

(aq) + Cu(s) redox equation

13

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.

Ions flow through the Electrons flow through wires.

ion bridge.

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.

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.

Electrons always flow through the wires and meter from the metal higher up

the electrochemical series to the metal lower down.

14

graphite

electrodes

Cells with Non-Metals

The half-cells in a cell need not involve metal atoms.

solution containing ion iodine

sulfite ions (SO32-) bridge 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.

SO3

2-(aq) + H2O (l) + I2(aq) + 2e SO4

2-(aq) + 2H+

(aq) + 2e + 2I-

(aq)

add and cancel

SO3

2-(aq) + H2O(l) + I2(aq) SO4

2-(aq) + 2H+

(aq) + 2I-(aq)

redox equation

15

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 (x2) oxidation

O2(g) + 4H+(aq) + 4e 2H2O(l) (leave) reduction

2H2(g) 4H+(aq) + 4e oxidation

O2(g) + 4H+(aq) + 4e 2H2O(l) reduction

2H2(g) + O2(g) + 4H+(aq) + 4e 4H+

(aq) + 4e + 2H2O(l) add and cancel

2H2(g) + O2(g) 2H2O(l) redox

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.

LI 10

16

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) + SO4

2-(aq) + 2e oxidation

PbSO4(s) + 2e Pb(s) + SO42-

(aq) reduction

2PbSO4(s) + 2H2O(l) Pb(s) + PbO2(s) + 4H+(aq) + 2SO4

2-(aq) redox

When a battery is being recharged the energy change is:

electrical chemical

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

I am going to find out about:

Success Criteria

I can:

1 the reactivity of metals

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

2 ionic equations state the balanced ionic equation for a metal reacting with

oxygen, water or acid4

3 extracting metals from their ores 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

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

4 percentage composition

work out the percentage of a particular element in a compound

Topic 10 Pupil Self Evaluation

Metals – National 5

1

5 Redox Reactions 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

6 cells 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

7 The Electrochemical Series 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

8 oxidising and reducing agents state the definition of an oxidising and reducing agent

given a balanced equation state the reducing agent

9 more complicated cells 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

2

state where electrons flow in this set-up

state where ions flow in this set-up

state the purpose of an ion bridge

10 two examples of technologies which

make use of redox reactions

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

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…

1

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???