1. C4 revision
Easter revision

2. Atomic structure

3. 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.

4. A helium atom has a diameter of just 0.18 nm.
1 nm is one billionth of a metre or 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.

5. The protons and neutrons are in the centre of 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.
The nucleus is surrounded by a cloud of electrons.
The nucleus is many thousands of times smaller than the atom.

6. 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.

7. An atom is made up of: Protons charge = +1 mass = 1 Neutrons
Electrons
charge = -1
mass = negligible

8. 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.

9. The second shells can hold 8.
The first shell that is closest to the nucleus can hold up to 2 electrons.
The second shells can hold 8.
Once the second shell is full the third shell starts to fill. It also holds 8.
If there are more electrons they occupy further shells.
11 protons
therefore
11 electrons.
The electron arrangement can be written as 2,8,1

10. 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.

11. So what happens if the number of electrons changes?
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?

12. Ions
Atomic number
Mass number
Charge
Number of protons
Number of electrons
Lithium atom 3 7 4
Li+ +1 2
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

13. 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

14. 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
H – 2
H – 3 2

15. 56
Fe iron 26
protons
neutrons
electrons ?

16. 56
Fe iron 26
protons
neutrons
electrons 30

17. ?
Na sodium 11
protons
neutrons
electrons 12

18. 23
Na sodium 11
protons
neutrons
electrons 12

19. Group 1 metals

20. Li Lithium Group 1 Period 2 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.
Atomic structure:
Atomic Number: 3 Li
Lithium
Atomic Mass: 7
Electron Arrangement: 2,1
Properties
Melting Point
180oC
Boiling Point
1342oC
Density
0.534 g/cm3
Formula of Compounds
Lithium Hydroxide LiOH
Lithium Chloride LiCl
Lithium Oxide Li2O

21. Na Sodium Facts about Sodium Group 1 Period 3
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.
Group 1 Period 3
Atomic structure:
Atomic Number: 11 Na
Sodium
Atomic Mass: 23
Properties
Melting Point
78oC
Boiling Point
883oC
Density
0.968 g/cm3
Formula of Compounds
Sodium Hydroxide NaOH
SodiumChloride NaCl
Sodium Oxide Na2O

22. K Potassium Facts about Potassium Group 1 Period 4
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.
Group 1 Period 4
Atomic structure:
Atomic Number: 19 K
Potassium
Atomic Mass: 39
Properties
Melting Point
63oC
Boiling Point
579oC
Density
0.89 g/cm3
Formula of Compounds
Potassium Hydroxide KOH
Potassium Chloride KCl
Potassium Oxide K2O

23. Rb Rubidium Facts about Rubidium Group 1 Period 5
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.
Group 1 Period 5
Atomic structure:
Atomic Number: 37 Rb
Rubidium
Atomic Mass: 85
Properties
Melting Point
39oC
Boiling Point
688oC
Density
1.532 g/cm3
Formula of Compounds
Rubidium Hydroxide RbOH
Rubidium Chloride RbCl
Rubidium Oxide Rb2O

24. Cs Caesium Facts about Caesium Group 1 Period 6
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.
Group 1 Period 6
Atomic structure:
Atomic Number: 55 Cs
Caesium
Atomic Mass: 133
Properties
Melting Point
28oC
Boiling Point
671oC
Density
1.93 g/cm3
Formula of Compounds
Caesium Hydroxide CsOH
Caesium Chloride CsCl
Caesium Oxide Cs2O

25. Group 7: Halogens

26. 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

27. Metal + halogen → metal halide
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 chloride
2Na(s) + Cl2(g) → 2NaCl(s)

28. 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

29. 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

30. 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.

31. Remember the colours of the vapour...
Halogen reactivity
Fluorine
Remember the colours of the vapour...
Chlorine
Bromine
Iodine
Reactivity increases as you move up the group – how is this different to group 1?

32. Displacement of halogens
If 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.
This is called displacement.
sodium chloride
sodium fluoride
chlorine
fluorine + 
F2 (aq)
2NaCl (aq)
2NaF (aq)
Cl2 (aq)
A more reactive halogen will always displace a less reactive halide from its compounds in solution.

33. Displacement reactions
Miss Iodine
Mr Lithium +
Miss Chlorine
Miss Chlorine
Mr Lithium +
Miss Iodine
The halogen with the highest reactivity will take the metal from the halogen less reactive

34. 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
Potassium bromide and iodine
Iodine

35. Halogen Displacement Colour change ? KBr added KI added KI added
Bromine water
KCl added
Chlorine water
Iodine water
KBr added
KCl added
Colour change ?
Yes
Yes No
Yes No No

36. Ionic bonding

37. 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.

38. Here comes my friend, Sophie Sodium
Ionic bonding Cl
Hi. My name’s Johnny Chlorine. I’m in Group 7, so I have 7 electrons in my outer shell
Here comes my friend, Sophie Sodium Na
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?
Okay - + Na Cl
Now we’ve both got full outer shells and we’ve both gained a charge. We’ve formed an IONIC bond.

39. - + How do atoms form ions? Na Na Cl Cl [2,8]+ [2,8,8]-
Very common exam question. + Na Na
2,8,1
[2,8]+ - Cl Cl
[2,8,8]-
2,8,7

40. Atoms forming ions Metals form POSITIVE ions (Li+, K+)
Oxidation is loss of electrons
Reduction is gain of electrons
OIL RIG
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-)

41. Covalent bonding

42. They also do not conduct electricity.
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

43. 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

44. The atoms get so close to each other that their outer shells overlap.
Shared pair
A covalent bond
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.

45. 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

46. 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

47. Two pairs of electrons shared
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

48. 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.

49. 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

50. Metallic bonding

51. 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.

52. 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.

53. 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.

54. Metallic bonding: melting points
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 (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
<
<
Na+
Mg2+
Al3+
Increasing electron cloud density as more electrons are donated per atom. This means the ions are held more strongly

55. Transition metals

56. Transition metals
Transition metals are found in the centre box on the periodic table

57. Properties of the transition elements
All conduct heat (metals)
They are all shiny
All conduct electricity
All sonorous
All malleable (beaten into sheets)
All ductile (drawn into wires)
Often coloured in solution
Often 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

58. 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

59. Thermal decomposition of transition elements
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)