C4 revision

Easter revision

Atomic structure

Proposing a theory

• Sometimes indirect evidence and a bit of imaginative thinking has to be used to account for data.

• Since 1800 there have been a number of different atomic models proposed. Each one as a result of trying to explain observations that could not simply be done by using data.

A helium atom has a diameter of just 0.18 nm.

1 nm is one billionth of a metre or 0.000 000 001 m.

Atoms are so small that even with a powerful microscope we cannot see what they are made up of.

However, results of investigations have enabled scientists to create a model atom.

The nucleus is surrounded by a cloud of electrons.

The nucleus is many thousands of times smaller than the atom.

The protons and neutrons are in the centre of the atom. This is the nucleus.

All the mass of the atom is in the nucleus.

It helps to be able to see all the particles in the nucleus and the electrons at the same time…

…so the model we use is not to scale.

An atom is made up of:

Protonscharge = +1mass = 1

Neutronscharge = 0mass = 1

Electronscharge = -1mass = negligible

Arrangement of electrons…

• Electrons are arranged in shells, with different energy levels, around the nucleus.

• Each electron shell can contain only a limited number of electrons.

• The innermost shell (lowest energy) fills first.

• When it is full, the electrons go to the next shell.

11 protons

therefore

11 electrons.

The electron arrangement can be written as 2,8,1

How to work out the number of neutrons…

The relative atomic mass = protons + neutrons

How could you work out how many neutrons are in beryllium? Discuss with others on your table.

Mass no. – proton no. E.g. for beryllium 9 – 4 = 5 neutrons

So beryllium has 4 protons, 4 electrons and 5 neutrons.

Why are atoms neutral

• Why is an atom neutral and has no charge?

Positive proton charges cancel negative electron charges

4 protons (4+) will cancel 4 electrons (-4)

So what happens if the number of electrons changes?

Ions

• Electrons can be given or taken from atoms, this results in ions being formed

• Ions have a different number of electrons to protons so the atom becomes charged

• If electrons are lost – the ion is POSITIVE

• If electrons are gained – the ion is NEGATIVE

Atomic number

Mass number

Charge Number of

protons

Number of

electrons

Number of

electronsLithium

atom3 7 0 3 3 4

Li+ 3 7 +1 3 2 4

Ions

• You will find out why atoms change into ions later in the topic but…

In general, elements are more stable if they have a full outer shell of electrons

Isotopes

• Isotopes of an element have different numbers of neutrons in their atoms

• Chemically isotopes have exactly the same properties – all that changes is the number of neutrons:

Isotope Electrons Protons Neutrons

H – 1 1 1 0

H – 2 1 1 1

H – 3 1 1 2

56

Feiron

26

protonsneutronselectrons

??

26

56

Feiron

26

protonsneutronselectrons

3026

26

?

Nasodium

11

protonsneutronselectrons

1211

?

23

Nasodium

11

protonsneutronselectrons

1211

11

Group 1 metals

Electron Arrangement: 2,1

Facts about Lithium • Lithium is soft enough to be cut with a knife.

• Is a shiny silvery-white colour metal.

• Turns black in air to form an outside layer of Lithium Oxide.

• Stored under oil to stop it reacting with the air.

• Lithium is not found in pure form naturally due to its reactivity.

• Found in large quantities in seawater (in compounds).

• Discovered in 1817.

Formula of Compounds

Lithium Hydroxide LiOH

Lithium Chloride LiCl

Lithium Oxide Li2O

Group 1 Period 2

Atomic Number: 3

LiLithiumLithium

Atomic Mass: 7

Properties

Melting Point

180oC

Boiling Point

1342oC

Density 0.534 g/cm3

Atomic structure:

Formula of Compounds

Sodium Hydroxide NaOH

SodiumChloride NaCl

Sodium Oxide Na2O

Group 1 Period 3

Atomic Number: 11

NaSodiumSodium

Atomic Mass: 23

Properties

Melting Point

78oC

Boiling Point

883oC

Density 0.968 g/cm3

Atomic structure:Facts about Sodium • Sodium is soft enough to be cut with a knife.

• Is a shiny silvery-white colour metal.

• Turns white in air to form an outside layer of Sodium Oxide.

• Stored under oil to stop it reacting with the air.

• Sodium and chemicals with sodium in burn with a yellow flame.

•Sodium chloride is a white crystalline solid.

• It was discovered in 1807, but has never been found on its own as it is very reactive.

Facts about Potassium• Potassium is soft and can be easily cut with a knife.

• Is a silvery grey colour metal that gets a coating of potassium oxide very quickly.

• Stored under oil to stop it reacting with the air.

• Potassium and chemicals with it in burn with a purple flame.

•Its chloride is a white crystalline solid.

•It is never found on its own as it is very reactive. The first time the actual metal was made was in 1807.

• Chocolate and Bananas are a good source of Potassium.Formula of Compounds

Potassium Hydroxide KOH

Potassium Chloride KCl

Potassium Oxide K2O

Group 1 Period 4

Atomic Number: 19

KPotassiuPotassiu

mmAtomic Mass: 39

Properties

Melting Point

63oC

Boiling Point

579oC

Density 0.89 g/cm3

Atomic structure:

Facts about Rubidium • Rubidium is soft, like plastercine.

• Is a silvery grey/white colour metal.

• Stored in a glass container with a noble gas atmosphere to stop it reacting.

• Rubidium and chemicals with it in burn with a violet flame.

•Rubidium chloride is a white crystalline solid.

•It is never found on its own as it is very reactive.

• The first time the actual metal was made was in 1861 by 2 scientists, one of which was Robert Bunsen.Formula of Compounds

Rubidium Hydroxide RbOH

Rubidium Chloride RbCl

Rubidium Oxide Rb2O

Group 1 Period 5

Atomic Number: 37

RbRubidiuRubidiu

mmAtomic Mass: 85

Properties

Melting Point

39oC

Boiling Point

688oC

Density 1.532 g/cm3

Atomic structure:

Facts about Caesium • Caesium is the softest of all solid elements.

• Is a silvery gold colour metal.

• Stored in a glass container with a noble gas atmosphere to stop it reacting.

• Caesium and chemicals with it in burn with a blue flame.

•Caesium chloride is a white crystalline solid.

•It is never found on its own as it is very reactive.

• The first time the actual metal was made was in 1860 by 2 scientists, one of which was Robert Bunsen.

Formula of Compounds

Caesium Hydroxide CsOH

Caesium Chloride CsCl

Caesium Oxide Cs2O

Group 1 Period 6

Atomic Number: 55

CsCaesiumCaesium

Atomic Mass: 133

Properties

Melting Point

28oC

Boiling Point

671oC

Density 1.93 g/cm3

Atomic structure:

Group 7: Halogens

The Group 7 Elements

• The halogens are the group 7 elements• The halogens have low melting points and boiling points.

Fluorine has the lowest melting point and boiling point• The melting points and boiling points then increase as

you go down the group• At room temperature (25oC) this temperature, fluorine

and chlorine are gases, bromine is a liquid, and iodine and astatine are solids

• The colour of the halogens elements gets darker as you go down the group. Iodine is purple, and, as we would expect, astatine is black

The Group 7 elements

Reactions with metals

•The halogens react with metals to make salts called metal halides

Metal + halogen → metal halide

•For example, sodium reacts with chlorine to make sodium chloride (common salt)

Sodium + chlorine → sodium chloride2Na(s) + Cl2(g) → 2NaCl(s)

Group 7 Structure

• The halogens are diatomic

• This means they exist as molecules, each with a pair of atoms. Chlorine molecules have the formula Cl2, bromine Br2 and iodine I2

Uses of halogens• Halogens are bleaching agents. They will remove the

colour of dyes

• Chlorine is used to bleach wood pulp to make white paper

• Halogens kill bacteria. Chlorine is added to drinking water at very low concentrations. This kills any harmful bacteria in the water, making it safe to drink. Chlorine is also added to the water in swimming pools

Handling halogens

• The halogens are very reactive and poisonous, care must be taken when using them

• Chlorine is used only in a fume cupboard

• Iodine should not be handled (it will damage the skin)

• Gloves may be used, and goggles should be worn.

Halogen reactivity

Iodine

Bromine

Chlorine

Fluorine

Reactivity increases as you move up the group – how is this different

to group 1?

Remember the colours of the

vapour...

Displacement of halogensIf a halogen is added to a solution of a compound containing a less reactive halogen, it will react with the compound and form a new one.

sodiumchloride

sodiumfluoride chlorinefluorine + +

F2 (aq) 2NaCl (aq) 2NaF (aq) Cl2 (aq)++

A more reactive halogen will always displace a less reactive halide from its compounds in solution.

This is called displacement.

Displacement reactions

Mr LithiumMiss Iodine Miss Chlorine+ Miss Chlorine Mr Lithium Miss Iodine+

The halogen with the highest reactivity will take the metal from the halogen less reactive

Displacement reactions: Answers

The x shows no reaction. You MUST revise the reactions that occur with halogens

Salt

Halogen

Potassium fluoride

Potassium chloride

Potassium bromide

Potassium iodide

Fluorine Potassium fluoride and

chlorine

Potassium fluoride and

bromide

Potassium fluoride and iodine

Chlorine X Potassium chloride and

bromine

Potassium chloride and

iodine

Bromine X X Potassium bromide and

iodine

Iodine X X X

Halogen Displaceme

nt

Chlorine water

Bromine water

Iodine water KBr adde

d

KI adde

dKCl

added

KI adde

dKCl

added

KBr adde

d

Colour change ?

Yes Yes YesNo No No

Ionic bonding

What is an Ionic Bond?• A very strong force of attraction between the

positively charge ion and negatively charged ion.

• One element is a metal, the other is a non-metal.

Here comes my friend, Sophie Sodium

Hey Johnny. I’m in Group 1 so I have one electron in my outer shell.

This electron is far away from the nucleus so I’m quite happy to get rid

of it. Do you want it?

Cl

Now we’ve both got full outer shells and we’ve both gained a charge.

We’ve formed an IONIC bond.

Na

Okay

ClNa

+ -

Hi. My name’s Johnny Chlorine. I’m in Group 7, so I have 7 electrons in my outer shell

Ionic bonding

Cl

How do atoms form ions?

Na

Cl

Na

2,8,1

2,8,7

+

-[2,8]+

[2,8,8]-

Very common exam question.

Atoms forming ions

• Atoms that have too many electrons want to LOSE them and become oxidised (metals)

• Atoms that have too few electrons want to GAIN them and become reduced (non-metals)

• Gaining electrons (-ves) makes the ion negative

• Losing electrons (-ves) makes the ion positive

Metals form POSITIVE ions (Li+, K+)

Non metals for NEGATIVE ions (Cl-, I-)

Oxidation is loss of electrons

Reduction is gain of electrons

OIL RIG

Covalent bonding

These sorts of compounds are usually gases or liquids at room temperature. If they are solids then they have low melting points.

They also do not conduct electricity.

These properties are related to the way the

atoms bond together in the compounds

Covalent bonding

Oxygen:

•6 electrons in its outer shell

•It needs 8 to be stable.

Hydrogen:

•1 electron in its outer shell

•Needs 2 to be stable

Covalent bonding

The atoms get so close to each other that their outer shells overlap.

Where the overlap occurs each atom contributes an electron to make a shared pair. In this diagram the electrons from the Hydrogen have been represented with a cross. Those from the oxygen are represented by a dot. In reality all electrons are identical.

Shared pair

A covalent bond

The oxygen atom can now ‘claim’ to have 8 electrons in its outer shell

Each hydrogen atom can now ‘claim’ to have 2 electrons in its outer shell

Covalent bonds can be represented using ‘dot and cross’ diagrams

Methane has 4 Hydrogen atoms sharing 4 pairs of electrons with a

carbon atom.

Some elements also bond together to form molecules.

Here chlorine atoms share a pair of

electrons.

Cl – Cl

Some molecules contain atoms that share two pairs of electrons. These are called ‘double bonds’. Carbon dioxide is a good example

of this but there are others.

Two pairs of electrons shared Double bond

• Covalent bonding produces individual molecules.

• There is little or no attraction between the molecules (weak intermolecular forces).

• This is why covalent compounds have low melting and boiling points.

As electrons are shared on covalent compounds, they do not have electric charges. This means they will not conduct electricity.

Group number and period number

• The group number relates to the number of electrons in the outer shell of electrons

• The period the element is in relates to the number of shells filled. Lithium has 2 shells so it is in period 2

Metallic bonding

Metallic bonding

“The electrostatic attraction between a lattice of positive ions surrounded by delocalised electrons”

Metal atoms achieve stability by “off-loading” electrons to attain the electronic structure of the nearest noble gas.

This results in a lattice of positive ions and a “sea” of delocalised electrons. These electrons float about and are not associated to a particular atom.

Metallic bonding: electrical conductivity

Because the electron cloud is mobile, electrons are free to move throughout its structure.

When the metal is part of a circuit, electrons leaving create

a positive end and electrons entering create a negative end. These new arrivals join the “sea” already present.

Metallic bonding: malleability

Metals are malleable: they can be hammered into shapes.

The delocalised electrons allow metal atoms to slide past one another without being subjected to strong repulsive forces that would cause other materials to shatter.

This allows some metals to be extremely workable. For example, gold is so malleable that it can make translucent sheets.

Increasing electron cloud density as moreelectrons are donated per atom.

This means the ions are held more strongly

Metallic bonding: melting points

Na (2,8,1) Mg (2,8,2) Al (2,8,3)Melting point 89°C 650°C 659°C

Boiling point 890°C 1110°C 2470°C

The melting point is a measure of how easy it is to separate the individual particles. In metals it is a measure of how strong the electron cloud holds the positive ions.

Na+ Al3+Mg2+

<<

Transition metals

Transition metals

• Transition metals are found in the centre box on the periodic table

Properties of the transition elements

All conduct heat (metals)They are all shinyAll conduct electricityAll sonorousAll malleable (beaten into sheets)All ductile (drawn into wires)Often coloured in solutionOften catalysts: Ni used in margarine making and Fe

used in Haber process

Look at their colours – take a

note of the colour of copper

and iron

Thermal decomposition of transition elements

• Occurs when a substance is broken down into at least two other substances by heat

• Transition metals form carbonates that are often coloured (FeCO3 – iron carbonate)

• When the coloured carbonate is heated (thermally decomposed), it will be broken down into at least two products – a gas and a metal oxide

• When carbonates are heated they produce a gas…what gas do you think would be formed?

Carbon dioxide – tested for with LIMEWATER

• Metal carbonates thermally decompose to form a metal oxide and carbon dioxide:

Iron carbonate iron oxide + carbon dioxide

FeCO3 FeO + CO2

What would the word equations be for the thermal decomposition of copper carbonate, zinc carbonate and manganese carbonate? (Extension: symbol eq)

Thermal decomposition of transition elements