CfE Chemistry

Summary Notes

S2/3

Pupil Name: _______________________

You must keep this booklet safe as it will also be used in

National 4/5 Chemistry.

2

Unit 1 – Solutions, Compounds and Mixtures

Elements and the Periodic Table

• Everything is made from elements in the Periodic Table.

• The Periodic Table contains around 109 different natural

elements.

• Man-made elements are shown with a * next to them in the

Periodic Table.

Each element has its own symbol.

E.g. Silver - Ag, Iron - Fe and Sodium - Na.

Each element has its own atomic number.

E.g. Aluminium is number 13.

Each element is made from only ONE type of atom.

E.g. Silver is only made from silver atoms.

Compounds

• Compounds are made by chemically joining two or

more different elements together.

• When the element sodium reacts together with

the element chlorine we get the compound sodium

chloride.

Naming Compounds

If a compound has the ending…

-IDE… this means there are only 2 elements present.

-ITE… this means the compound has 2 elements PLUS oxygen.

-ATE… this means the compound also has 2 elements PLUS oxygen.

Name of Compound Elements Present

Lithium bromide Lithium Bromine

Sodium nitrite Sodium Nitrogen Oxygen

Calcium carbonate Calcium Carbon Oxygen

3

Word Equations

Sodium metal can react with chlorine gas to produce the compound sodium

chloride. This can be written in a shorter format called a word equation:

REACTANTS PRODUCTS

Sodium + Chlorine Sodium Chloride

The ‘ ‘ means produces, makes or gives.

Chemical Reactions

All chemical reactions produce NEW substances.

Signs of a chemical reaction include:

1. A GAS forming.

This can be seen as fizzing or bubbling.

This is often known as effervescence.

2. A COLOUR CHANGE.

3. A PRECIPITATE formed.

This is when solutions react to form a solid.

4. During the reaction there is a change in ENERGY.

This can be exothermic – where heat is released – or

Endothermic – reaction gets cold.

Chemical reactions do NOT involve:

Melting boiling evaporating

Freezing condensing dissolving

The above are known as physical changes.

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Elements, Compounds and Mixtures

Mixture Compound

Chemically Joined

Easily separated

5

Solute, Solvent and Solutions

A solute is a substance which

can be dissolved in a liquid.

A solvent is a liquid in which a

solute can be dissolved.

A solution is what is produced

when a solute is dissolved in a

solvent.

A saturated solution is one in which no more solute can be dissolved.

A concentrate solution is one which contains a large amount of dissolved solute.

A dilute solution is one which contains a small amount of dissolved solute.

Solubility

If a substance dissolves in water it can be called soluble in water. Eg salt

If a substance does not dissolve in water it is called insoluble in water. Eg sand

• The higher the temperature of the solvent the faster the

substance will dissolve.

E.g. Sugar will dissolve faster in water at 80ºC than water at 25ºC

• The smaller the particle size of the substance the quicker

the substance will dissolve.

E.g. Sugar granules will dissolve quicker than sugar lumps.

Some substances that do not dissolve in water can be dissolved in

other solvents.

• Example: Nail varnish does not dissolve in water but does

dissolve in the solvent acetone.

6

Separating Mixtures

There are various methods used in the science lab to separate mixtures;

1. Filtration is used to separate a solid

from a liquid. The solid collected in

the filter paper is called the residue

and the liquid collected is called the

filtrate.

Eg. Used to separate sand and water.

2. Distillation can be used to separate

two different liquids. As different

liquids can have different boiling

points, one liquid can be boiled to

produce a gas which is then condensed

back into a liquid again.

Eg. Used to separate alcohol and water.

3. Chromatography can be used to find out which liquids are

contained within a mixture of liquids. Chromatography can

separate the liquids.

Eg. Used to separate the colours in ink.

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Unit 2 – Chemical Reactions

Rate of Reaction

How fast a new substance is made is called the ‘speed/rate of reaction’.

Reactions take place when particles in the chemicals collide with each other.

The rate of a reaction depends on:

1. The frequency or how often the collisions between the reactant particles

happen.

2. The energy with which reactant particles collide.

As the reaction progresses, the concentration of reactants decrease, and the concentration

of the products increases.

Concentration

• The units of concentration can be shown as M (eg. 1M) or mol/l (eg. 0.5mol/l).

• The higher the number, the greater the concentration of the solution.

• At a higher concentration, there are more particles in the same volume of space.

This means that the particles are more likely to collide and therefore more likely to

react.

High concentration concentra

tion

Low Concentration concentra

tion

8

Particle Size

Any reaction involving a solid can only take place on the surface of the solid.

If the solid is split into smaller pieces, the surface area increases.

This means that there is an increased area for the reactant particles to collide with.

The smaller the pieces, the larger the surface area. This means more collisions

and a greater chance of reaction.

Low Surface Area High Surface Area

• The following graph shows the

difference in reaction rate.

• Powder reacts faster than lumps and

therefore the powder graph has a

steeper gradient than the ‘lump’ graph.

This explains why potatoes chopped into smaller pieces cook faster than larger lumps.

Temperature

• If we increase the temperature of a reaction we give the particles

more energy.

• This means they will move faster and therefore are more likely to

collide with other particles.

• When the particles collide, they do so with more energy, and so the

number of successful collisions increases.

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Catalysts

• Another way we can speed up the rate of a chemical reaction is to add a catalyst.

• Unlike changing particle size, concentration or temperature, a catalyst is a substance

we add to a reaction to speed the reaction up.

• An advantage of using a catalyst is that it is not used up during the reaction and can

be used again.

• Enzymes are ‘biological catalysts’ and can be found in animals and plants.

E.g. Amylase, which is found in saliva.

Hazard Symbols

Corrosive Toxic Flammable Irritant

10

Unit 3 – Elements, The Periodic Table and Bonding

The Periodic Table

• The Periodic Table is

split into groups which

run vertically up and

down the table.

• The Table is also split

into periods which go

horizontally across.

Metals & Non-Metals

The thick zig zag line shown in the

diagram separates the metals and

non-metals.

Metals are found to the left of

the line and non-metals to the

right.

This information is also found on

page 3 of your databook.

Group Name Properties Examples

1 Alkali Metals Very reactive Sodium, Lithium

7 Halogens Reactive, used for

killing bacteria Chlorine, Fluorine

8/0 Noble Gases Unreactive Neon, Xenon

Middle section Transition Metals Can vary Copper, Silver

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Metals v. Non-Metals

Metals have many uses; copper for wires, aluminium for cans and planes and iron for bridges

and fences.

The Atom

The atom is made from 3 different

subatomic particles.

For every element:

The Atomic Number = The number of protons.

The Mass Number = The number of protons + neutrons.

To find the number of neutrons in an element we use:

Neutrons = Mass Number - Atomic Number

Metals Non-Metals

Solid.

(except MERCURY which is LIQUID) Can be solid, liquid or gas.

Conduct electricity. Do NOT conduct.

(except CARBON in the form of GRAPHITE)

High melting and boiling point. Lower melting and boiling point.

Higher density. Lower density.

Particle Name Mass Charge Location

PROTON 1 Positive Inside nucleus

NEUTRON 1 No charge Inside nucleus

ELECTRON 0 Negative Outside nucleus

NEUTRON

PROTON

ELECTRON

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**In a NEUTRAL atom the number of protons is EQUAL to the number of electrons.**

Nuclide Notation

No. of Protons = 17

No of Electrons = 17

No of Neutrons = 35 – 17= 18

Electron Arrangements

As the number of electrons increases they arrange themselves in a particular order in

energy levels/shells.

Energy level Maximum number of electrons

1 2

2 8

3 8 or 18

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The electron arrangement (Page 3 databook) shows how electrons are arranged in atoms.

E.g. Sodium

2 electrons in its 1st shell

8 electrons in its 2nd shell

1 electron in its 3rd shell

The Periodic Table group number tells us the number of outer electrons the element has

Eg. Group 3 elements have 3 electrons in the outer level.

Ions

• Noble Gases have a stable electron arrangement. • Noble Gases are like the rockstars of the Periodic Table, atoms want to be like them. • In order to do that they LOSE or GAIN electrons to have the same electron

arrangement. • When this happens atoms form IONS, which are charged particles.

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E.g. Na Na+ + e-

Electron arrangement: 2,8,1 2,8 1

Sodium now has the electron arrangement of Neon.

Rule: Metals lose electrons to form POSITIVE IONS. Non-metals gain electrons to form NEGATIVE IONS.

The Covalent Bond

Remember atoms want to have the electron arrangement of a noble gas so one way they can

do this is by sharing electrons.

A COVALENT bond is formed when 2 non-metal atoms share outer electrons.

The covalent bond is very strong and is difficult to break.

Covalent bonds can be single bond (Cl-Cl), double bond (O=O) or triple bond (N≡N).

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Diatomic Molecules

• A diatomic molecule contains 2 atoms. There are SEVEN elements that exist as

diatomic molecules on the periodic table.

Hydrogen - H2

Oxygen - O2

Nitrogen - N2

Fluorine - F2

Chlorine - Cl2

Bromine - Br2

Iodine - I2

Shapes of Molecules

When atoms bond together the molecules produced can form different shapes.

The four main shapes are…

Linear V-Shaped/Bent Pyramidal Tetrahedral

Chemical Formula

Chemical formula is a shorthand way of showing elements and compounds using

symbols from the Periodic Table.

Some covalent formula (non-metals) can be found by using the symbols given in the

name of the compound.

Eg. Silicon Tetrachloride - SiCl4

Formula using prefixes

If there is a prefix present in the name of the compound it tells you the number of atoms

present.

E.g. Carbon Dioxide – the carbon does not have a prefix so there must be only one C, the

oxygen has a DI prefix which means there are two O’s. The finished formula is CO2.

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Prefix Number

Mon or mono 1

Di 2

Tri 3

Tetra 4

Pent 5

Formula using the Crossover Rule

If the name of the compound does not contain any prefixes we can use the Crossover Rule.

The valency number of an element is used; this is the number of bonds an atom can form.

Group 1 2 3 4 5 6 7 8/0

Valency 1 2 3 4 3 2 1 0

Chemical formula is written using the following steps

1. Symbols

2. Valency number

3. Swap - crossover

4. Divide - if there is a common number*

5. Write the final formula – remember not to show

the 1 e.g C1H4 is shown as CH4

*The divide step is not always required.

As the Transition Metals do not have a group number their valency is given using Roman

Numerals

1 = (l) , 2 = (ll), 3 = (lll), 4 = (lV), 5 = (V), 6 =(Vl)

Eg. Nickel (II) bromide - Nickel has a valency of 2

Iron (III) oxide - Iron has a valency of 3

Symbols C H

Valency 4 1

Swap 1 4

Divide*

Formula CH4

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Examples

Iron (II) Oxide Hydrogen Oxide

Carbon Tetrachloride* Silicon Hydride

Aluminium Oxide Nitrogen Trihydride*

Carbon Sulphide

* do not use valency rules when the compound has a prefix

Symbols Fe O

Valency 2 2

Swap 2 2

Divide 1 1

Formula FeO

Symbols H O

Valency 1 2

Swap 2 1

Divide

Formula H2O

Symbols C Cl

Valency

Swap

Divide

Formula CCl4

Symbols Si H

Valency 4 1

Swap 1 4

Divide

Formula SiH4

Symbols Al O

Valency 3 2

Swap 2 3

Divide

Formula Al2O3

Symbols N H

Valency

Swap

Divide

Formula NH3

Symbols C S

Valency 4 2

Swap 2 4

Divide 1 2

Formula CS2

18

Unit 4 - Fuels

Fossil Fuels

• Coal, oil and natural gas were created millions of years ago and

are known as the Fossil Fuels.

• Coal was formed from dead plants that sank to the bottom of

swampy water and over millions of years were buried with layers

of mud.

• These layers of mud were then compressed by the pressure as more layers were

added and formed coal.

• Oil and gas were formed in a similar way from tiny sea creatures that sank to the

bottom of the sea millions of years ago.

• The pressure of more layers of sand caused the production of oil and natural gas.

Fuels

Fuels are substances which burn in oxygen to release energy.

E.g Coal, oil, gas, wood, peat and sugar

However, not all substances which give out energy are fuels.

A battery is NOT a fuel as it does not burn.

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Fractional Distillation

Crude Oil is a mixture of different chemical compounds.

• By using fractional distillation we can separate the different compounds as each has a

different boiling point.

• The fractional distillation tower is used to separate the mixture into groups with

similar boiling points called fractions.

• The tower is hot at the bottom to collect fractions with a high boiling point, and cool

at the top to collect fractions with a low boiling point.

Fractional Distillation

Top of the Tower Bottom of the Tower

Small molecules Large molecules

Low boiling point High boiling point

Low viscosity High viscosity

High flammability Low flammability

Easily evaporated Difficult to evaporate

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Composition of Air

Approximately only 20% of the air

around us is Oxygen. The majority is

made up of Nitrogen.

The ratio of Oxygen to Nitrogen in

air is 1 : 4.

Combustion

When a fuel is burned in oxygen it is called combustion. This is an example of an

exothermic reaction. There are 2 types of combustion.

1. Complete Combustion – A plentiful supply of oxygen.

FUEL + O2 CO2 (Carbon dioxide) + H2O (Water)

2. Incomplete Combustion – A limited supply of oxygen.

FUEL + O2 C (Carbon/Soot) + CO (Carbon monoxide) + H2O

**The test for oxygen is that it relights a glowing splint**

Alkanes

A homologous series is a set of compounds with similar chemical properties which can be

represented by a general formula.

Alkanes:

are a subset of the set of hydrocarbons (contain only Hydrogen and Carbon)

all end in the letters -ane

are a homologous series with general formula CnH2n+2

contain single C-C bonds (saturated)

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each alkane had a different name depending on how many carbon atoms are present.

E.g. The alkane with 3 carbons is called propane.

Note:

If you forget the prefixes then look at page 6

of the data booklet. The alkanes are listed in a

table, in order, so you can work out the number

of carbons from that.

Structure of the Alkanes

Name Formula Full Structural Formula Shortened Structural

Formula

Methane CH4

CH4

Ethane C2H6

CH3CH3

Propane C3H8

CH3CH2CH3

Butane C4H10

CH3CH2CH2CH3

Prefix Number of Carbons

meth- 1

eth- 2

prop- 3

but- 4

pent- 5

hex- 6

hept 7

oct- 8

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Pollution

The burning of fossil fuels causes POLLUTION!

1. Acid Rain

• Fossil fuels can often contain sulphur, which when burned

produces sulphur dioxide SO2. This acidic gas dissolves in clouds

and falls as acid rain.

• Nitrogen oxides NOx are made by lighting or by spark plugs in car

engines. This is another acidic gas which dissolves in clouds and

falls as acid rain.

2. Global Warming

• When a fuel is burned in enough oxygen it produces CO2

– a “Greenhouse gas”.

• This gas is building up in the atmosphere and causing the

world to heat up, also known as the “Greenhouse Effect”

• This could mean the melting of the polar ice caps with a

rising of sea levels and more severe weather.

3. Transportation of Crude Oil

• Oil tankers can crash spilling oil and causing damage to the sea.

Reducing Pollution

• To reduce pollution we have to use RENEWABLE ENERGY to make our electricity and

run our cars and buses.

HYDROELECTRIC --- uses falling water to make electricity.

SOLAR --- uses energy from the sun

WIND FARMS --- uses the wind to make electricity.

WAVE POWER --- uses the waves in the sea to make electricity.

BIO-FUELS --- from plants to help us run our cars and buses

Renewable Fuels

Ethanol, obtained from sugar cane, is a RENEWABLE fuel.

23

Unit 5 – Acids & Alkalis Ionic Introduction

Ionic bonding occurs between a metal and a non-metal.

The metal gives electrons away and the non-metal accepts electrons.

Metals form positive ions

Non-metals form negative ions

Ionic compounds from an ionic lattice. The ions are held together by electrostatic

attraction.

Ionic Formula

The charge for each ion can be obtained from the Periodic Table as follows:

** Groups 4 and 0 do NOT form ions**

Ionic formulae are worked out by writing the symbols or formulae for the positive and

negative ions. Then, the positive and negative charges must be "balanced" (if they are not

already the same) as below.

Eg. sodium chloride Na+ Cl-

copper (II) sulphide Cu2+ S2-

If the positive and negative ions don't have the same number of charges, we have to work

out how many of the positive ions and how many of the negative ions would be needed to

make the whole compound neutral.

When we need to show more than one ion in a formula, we put brackets round the ion as

below.

Eg. sodium oxide (Na+)2 O2-

calcium bromide Ca2+ (Br-)2

iron (III) chloride Fe3+ (Cl-)3

aluminium oxide (Al3+)2 (O2-)3

Group 1 2 3 4 5 6 7 0

Charge on ion 1+ 2+ 3+ ** 3- 2- 1- **

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Examples

Calcium oxide Sodium fluoride

Magnesium chloride Potassium sulphide

Formula involving complex ions

• Complex ions contain more than one kind of atom.

• These are found on pg 6 of the databook.

• E.g. CO32- , NO3

-, NH4+ and OH-

• The valency of these ions is the same as their charge. The formula is worked out in

the same way, using the crossover method.

Examples

Potassium nitrate Ammonium chloride

Symbols Ca O

Valency 2 2

Swap 2 2

Divide 1 1

Chem Formula CaO

Ionic Formula Ca2+O2-

Symbols Na F

Valency 1 1

Swap 1 1

Divide

Chem Formula NaF

Ionic Formula Na+F-

Symbols Mg Cl

Valency 2 1

Swap 1 2

Divide

Chem Formula MgCl2

Ionic Formula Mg2+(Cl-)2

Symbols K S

Valency 1 2

Swap 2 1

Divide

Chem Formula K2S

Ionic Formula (K+)2S2-

Symbols K NO3

Valency 1 1

Swap 1 1

Divide

Chem Formula KNO3

Ionic Formula K+NO3-

Symbols NH4 Cl

Valency 1 1

Swap 1 1

Divide

Chem Formula NH4Cl

Ionic Formula NH4+Cl-

25

Calcium hydroxide Sodium carbonate

The pH Scale

pH Scale measures how acidic or alkaline a solution is.

The pH can be found using universal indicator, litmus paper or using a pH meter.

Acid Alkali

Common lab

Hydrochloric acid HCl

Sulphuric acid H2SO4

Nitric acid HNO3

Sodium hydroxide NaOH

Potassium hydroxide KOH

Common household Coke

lemonade

Bleach

Oven cleaner

All acids contain the hydrogen ion H+

All alkalis contain the hydroxide ion OH-

Acids – pH less than 7 Neutral – pH 7 Alkali - pH more than 7

Symbols Ca OH

Valency 2 1

Swap 1 2

Divide

Chem Formula Ca(OH)2

Ionic Formula Ca2+(OH-)2

Symbols Na CO3

Valency 1 2

Swap 2 1

Divide

Chem Formula Na2CO3

Ionic Formula (Na+)2CO32-

11 22 33 44 55 66 77 88 99 1100 1111 1122 1133 1144

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HH++

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HHiigghh

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26

Making Acids

When we burn certain non-metals in the presence of O2,

a non-metal oxide is produced.

E.g. C(s) + O2(g) CO2(g)

The non-metal oxide, if soluble, can then be dissolved in water to produce an acid.

CO2(g) + H2O(l) (H+)2 CO32-(aq) (Carbonic acid)

Acid Rain

• Coal contains sulphur which when burned produces sulphur dioxide.

S(s) + O2(g) SO2(g)

• When the sulphur dioxide rises and is absorbed by the clouds sulphuric acid is

formed.

SO2(g) + H2O(l) (H+)2 SO42- (aq) (Sulphuric acid)

Other causes of acid rain

• The energy from lightening in thunderstorms and the energy from the spark plug in

an engine allows the nitrogen and oxygen in air to react together.

• The nitrogen dioxide is then absorbed into the clouds to form nitric acid.

NO2(g) + H2O(l) H+NO3- (aq) (Nitric Acid)

Making Alkalis

If a metal oxide or metal hydroxide dissolves in water, an alkali is formed.

i) Sodium oxide dissolves to form sodium hydroxide solution.

ii) Calcium hydroxide dissolves to form calcium hydroxide solution.

State Symbols

Solid s

Liquid l

Gas g

Aqueous aq

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Diluting Acids & Alkalis

When water is added to an acid the pH moves

towards 7 as the concentration of H+ ions

decreases.

When water is added to an alkali the pH moves

towards 7 as the concentration of OH-

decreases.

Neutralisation – Acid & Alkali

• Acids can be neutralised by certain substances to make neutral compounds (pH 7).

The reaction is called neutralisation and the substance which reacts with the acid is

called a neutraliser or a base.

• During neutralisation a SALT is always produced.

Everyday neutralisation reactions include:

a) adding lime to rivers and lochs to reduce the effects of acid rain.

b) treating a bee sting (acid) with baking soda (alkaline).

c) treating a wasp sting (alkaline) with vinegar (acid).

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Neutralisation Reactions

When a salt is produced in a neutralisation reaction part of the name comes from the acid

used in the reaction and the other part from the neutraliser/base.

Acid Salt Produced

Hydrochloric Chloride

Sulphuric Sulphate

Nitric Nitrate

1. Acid & Alkali

Acid + Alkali (neutraliser) Salt (neutral) + Water (neutral)

The hydrogen (H+) from the acid reacts with the hydroxide (OH-) from the alkali to form

water.

Eg.

Full equation H+Cl-(aq) + NaOH-(aq) Na+Cl- + H2O(l)

H+(aq) + OH-(aq) H2O(l) (pH 7)

2. Acid & Reactive metal

Reactive Metal + Acid Salt + Hydrogen

Eg. Magnesium + nitric acid Magnesium nitrate + Hydrogen

**The test for hydrogen gas is that is burns with a “pop”.**

3. Acid & Carbonate

Carbonate + Acid Salt + Carbon dioxide + Water

Eg. Calcium carbonate + Hydrochloric acid Calcium chloride + Carbon dioxide + Water

**The test for carbon dioxide is that it turns limewater cloudy/milky**

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The Mole

The mole is the formula mass of an element, compound or molecule, expressed in grams.

1. For Elements

The gram formula mass of Aluminium is 27 grams, this information is found in page 6 of your

databook. Therefore 1 Mole of Aluminium is 27 grams.

Note: For diatomic elements the databook value is doubled.

E.g. for chlorine Cl2 the mass of one mole is 71g not 35.5g.

2. For Molecules or Compounds

The chemical formula of a molecule or compound tells us which elements are present, and in

what quantity.

Calculating the mass of all the elements present will give the mass of one mole of the

molecule or compound.

E.g. Lithium Oxide (Li2O).

The formula tells us there are 2 atoms of lithium and 1 atom of oxygen present.

Therefore:

1 Mole of Li2O = 30g

E.g. Sodium Hydroxide (NaOH)

The formula tells us there is 1 atom of sodium, 1 atom of oxygen and 1 atom of

hydrogen present.

1 Mole of NaOH = 40g

30

Unit 6 – Metals, Reactivity Series, Electricity & Corrosion

Properties of metals

Iron is a metal that is strong and so it can be used to build bridges and

railway tracks.

Copper metal has many uses. It has good electrical conduction. This

means that it is very good at conducting electricity so it is used to make

electrical wires. Copper has good thermal conductivity. This means that

it is good at letting heat move through it and so it is sometimes used to

make cooking pots. Copper does not corrode so it is also good for making

water pipes.

Aluminium is a metal that has a very low density.

This means that it is light so it is used to make aeroplanes.

Metals that are used to make jewellery, have to be malleable.

This means that they are can be hammered easily into different shapes.

Gold, silver and platinum are very malleable metals. These metals are

also used because they do not corrode and stay shiny for a long time.

Tin metal does not corrode so it is a good metal to use to make food cans.

Alloys

Pure metals do not always have the properties that we need for a particular job.

An alloy is a substance made by melting and mixing metals with other elements.

Alloys are often more useful than pure metals because they have different

properties from the pure metals.

This can make them more suitable for certain uses.

E.g. Solder is an alloy of TIN and LEAD

Stainless steel is an alloy of IRON, CARBON and CHROMIUM

(a transition metal)

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Reactions of Metals

Metal + Water Metal hydroxide + Hydrogen

Metal + Acid Salt + Water

Metal + Oxygen Metal oxide

Reactivity Series

By observing the reactions of metals we are able to build the “Reactivity Series” which

shows how reactive metals are relative to each other.

Extraction of Metals

Some metals are uncombined. This means we find them pure in the ground; not joined

to other elements. Examples of uncombined metals are silver and gold.

Most metals are not like this. We find them combined with other elements in a

compound known as an ORE. Most ores are metals joined to oxygen (metal oxide).

To obtain a metal from a metal oxide we need to separate the metal atoms from the

oxygen atoms. The more reactive the metal the tighter it is joined to the oxygen so

the harder it is to produce!

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Extraction of Metals

Batteries

In a battery electricity comes from a chemical reaction.

Batteries need to be replaced when the chemicals inside are used up.

Some batteries are rechargeable; e.g. the lead-acid battery.

All batteries contain electrolytes. (Usually ammonium chloride paste)

The purpose of the electrolyte is to complete the circuit.

Mains Power v. Batteries

Advantages Disadvantages

Mains Power Cheaper Danger of being electrocuted

Uses finite resources

Batteries Safe to use

Portable

Expensive

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The Electrochemical Series (ECS)

Shows metals in order of their willingness to give up electrons.

Shown on page 7 of the data booklet (Reactivity Series).

Cells

Electricity can be produced by connecting two different metals together (with an

electrolyte eg ammonium chloride) to form a cell.

Electricity can also be produced in a cell by connecting two different metals in

solutions of their metal ions.

Electrons flow from

the metal higher in

the ECS to the

metal lower i.e.

from magnesium to

lead.

The purpose of the “ion bridge” is to complete the circuit.

The greater the distance between the metals in the electrochemical series, the

higher the voltage produced.

Displacement Reactions

Reactions which occur when a metal higher up in the electrochemical series is added

to a solution containing ions of a metal lower down in the series.

E.g. Magnesium metal added to copper ions;

Mg(s) + Cu2+SO42-(aq) Mg2+SO4

2-(aq) + Cu(s)

The higher magnesium metal will form ions in solution

The lower copper ions will form copper metal and come out of solution.

I.e. they will be displaced.

34

The ECS can be used to predict whether or not a displacement reaction will occur.

The metal being added must be higher than the ions in solution for displacement to

occur.

Corrosion

Corrosion is a chemical reaction which involves the surface of a metal changing from

an element to a compound (metal oxide).

This is an example of oxidation.

The corrosion of iron is known as rusting.

Water and oxygen are required for corrosion to occur.

Preventing Corrosion

A surface barrier to air (oxygen) and water can provide physical protection against

corrosion.

E.g.

Painting

Greasing

Coating with plastic

Galvanising - metal objects are dipped into molten zinc

Electroplating - silver, chromium and other metals can be deposited on the

surface of a metal

Tin-plating

Gas Tests - Summary

Gas Test

Oxygen Relights a glowing splint

Hydrogen Burns with a ‘pop’

Carbon dioxide Turns limewater milky/cloudy

35

Unit 7 – Bonding, Structure & Properties

Electrical Conductivity

Electric current is a flow of charged particles

Testing conductivity

Metal elements and carbon (graphite) are electrical conductors.

Non-metal elements are non-conductors of electricity.

Covalent compounds do not conduct in any state as they only

contain non-metals.

Ionic compounds conduct when molten or in solution.

Ionic compounds do not conduct when solid.

1. electrons flow through

metals.

2. ions flow through solutions or

melts.

State at Room

Temperature

Covalent

compounds can be

solids, liquids or

gases.

Ionic compounds

are all solids.

Testing a solution

Testing a solid

Testing a solution

36

Structure of Compounds

Ionic solids exist as lattices of Some covalent solids are network structures

oppositely charged ions. E.g. SiO2, diamond

E.g. Na+Cl-

Ionic lattices and covalent networks have strong bonds and

therefore high melting points.

Other covalent compounds exist as discrete molecules.

(see opposite)

Discrete covalent molecules have weak forces of attraction and

so these compounds have low melting points.

Solubility

Some covalent substances do not dissolve in water but will dissolve in other covalent

solvents.

Eg. Nail varnish being dissolved in propanone.

Paint dissolving using turpentine.

Ionic compounds generally dissolve in water. Eg. Sodium chloride & copper sulphate.

Dissolving ionic compounds break up the ionic lattice.

Solid ionic lattice Dissolved ionic compound

(ions are free to move)

37

Electrolysis

Electrolysis is used to break up a compound into its elements using electricity.

Positive metal ions are attracted to the negative electrode.

Negative non-metal ions are attracted to the positive electrode.

Electrodes are made from carbon in the form of GRAPHITE as it conducts electricity.

A d.c. (direct current) supply must be used if the products are to be identified.

Covalent compounds cannot be electrolysed as they do not conduct electricity and they

have no ions.

Coloured Ions

Some ions are coloured. Examples are shown in the table below.

Ion Colour

copper blue

nickel green

potassium colourless

chromate yellow

permanganate purple

sulphate colourless

38

Migration of Ions

In the experiment below;

the positively charged blue copper ions move towards the negative electrode

the negatively charged orange dichromate ions move towards the positive electrode.

39

Glossary

Definition

Acid Substance with a pH less than 7.

Acid rain Source of pollution caused by sulphur dioxide SO2 and nitrogen dioxide gas

NO2

Alkali Substance with a pH greater than 7.

Alkali metals Very reactive group 1 elements.

Atomic number Number of protons in the nucleus of an atom.

Base Substance which will neutralise an acid.

Blast furnace A tall oven used to extract iron from iron ore by burning it with carbon at

high temperatures.

Catalyst A substance that increases the rate of a chemical reaction without being

used up.

Chemical

reaction

A change in which new substances are made and cannot easily be reversed.

Combustion Reaction which involves burning a substance in oxygen (O2).

Concentration The number of molecules of a substance in a given volume. Units are mol/l,

mol l-1 or M

Concentrated

solution

A solution that contains a large amount of solid dissolved.

Corrosion Surface of a metal changing from a metal to a compound (metal oxide).

Covalent bond A shared pair of electrons between two non-metal atoms.

Crude oil A fossil fuel made up of a mixture of hydrocarbons.

Diatomic Containing only 2 atoms. E.g. HCl, H2, NO, CO, O2

Dilute solution A solution containing a small mass of dissolved solute.

Electrolysis Breaking a compound into its elements using electricity.

Electron Negatively charged subatomic particle found in electron shells.

Electron

arrangement

Shows the number of electron shells and how many electrons they contain.

Element A substance made up of only one type of atom.

Endothermic A reaction which releases heat to the surroundings ie gets cold.

Exothermic A reaction which releases heat.

Fossil fuel Fuels formed from fossils over millions of years at a high temperature and

pressure; coal, oil and gas.

40

Fractional

distillation

Method of separating crude oil depending on boiling point.

Group Vertical column of elements in the Periodic Table, with similar properties.

Graphite Form of carbon with is an electrical conductor. Used as electrodes.

Halogen Group 7 elements in the Periodic Table.

Hydrocarbon A molecule containing only hydrogen and carbon.

Insoluble A substance than cannot be dissolved in water.

Ion A charged particle formed when an atom loses or gains electrons.

Ionic bond A bond formed between a metal and a non-metal.

Loam A mixture of small pieces of rock or sand surrounded by decayed animal or

plant remains.

Metal

hydroxide

A compound containing metal, hydrogen and oxygen atoms which will

neutralise an acid and has a pH greater than 7.

Metal oxide A compound containing metal and oxygen atoms only.

Metamorphic

rock

Rock formed when heat and pressure cause changes in existing igneous or

sedimentary rocks over a long period of time.

Mixture Two or more substances brought together but are not chemically joined.

Eg. Air and crude oil.

Mohs scale Scale used to measure the ‘hardness’ of a mineral.

Molecule A small group of atoms that are held together by covalent bonds.

Noble gases Group 8/0 elements in the Periodic Table. Unreactive elements with a full

outer shell of electrons.

Nucleus The small, dense, postively charged centre of an atom, made up of protons

and neutrons.

Neutral Substance with a pH 7 which turns universal indicator green.

Period Horizontal line going across the Periodic Table.

Proton Positive subatomic particle found within the nucleus.

Precipitation When two solutions react and a solid is formed.

Reactive A substance that reacts quickly or easily. Eg. Alkali metals.

Reactivity How quickly or easily a substance will react.

Rusting The specific name for the corrosion of iron.

Saturated

solution

A solution in which no more solute can be dissolved.

Soluble A solute which can dissolve in a liquid.

Solute A solid which is dissolved in the solvent.

Solution A liquid containing a dissolved solid.

Solvent A liquid which can dissolved a solute.

41

Transition

metals

Metals that are found in the middle of the periodic table. They do not

have a group number given to them. Their valency is given in Roman

numerals. Eg. Titanium.

Universal

Indicator

Solution used to test the pH of substances.

Acids – red, alkali – blue and neutral – green.

Unreactive A substance that reacts very slowly or does not react at all. Eg. Noble

gases.

Valency How many bonds an atom can form. Eg. Group 4 elements have a valency

(combining power) of 4 and therefore form 4 bonds.