5.1 Explain how the methods of extraction if the metals in this section are related to their positions in the reactivity series

The reactive series

Anything LESS reactive than carbon can be displaced from its ore by carbon (e.g.: iron)
Anything MORE reactive than carbon can't and so is extracted by electrolysis (e.g.: aluminium)

2.36 Understand the sacrificial protection of iron in terms of the reactivity series.

Sacrificial protection of iron means coating it with a layer of metal that is more reactive. This means that the reaction will take place with the metal that is more reactive rather than the iron itself. Sometimes, ships will use blocks of magnesium because electrons involved in the reaction will flow like a current to the more reactive metal.

2.35 Describe how the rusting of iron may be prevented by grease, oil, paint, plastic and galvanising

Anything that prevents air or water from coming in contact with the iron will stop or delay rusting. This is what grease, oil, paint and plastic does. Galvanising is putting a layer of zinc over the iron. This is also effective because, unlike iron, zinc doesn't rust.

2.34 Describe the conditions under which iron rusts

Water and oxygen cause iron to rust, forming hydrated iron (III) oxide. Okay well water isn't needed, but there's oxygen in water, which causes rusting.
Salty water can make iron rust faster because it is an electrolyte (a liquid containing ions which can conduct electricity). If it can conduct electricity it will allow the electrons to flow more easily, and hence speeding up electron loss (because we know that oxidisation can also be described as the loss of electrons)

Fig 1: Rusting over time

2.33 Understand the terms redox, oxidising agent, reducing agent

Redox: When a more reactive metal gains oxygen from the less reactive metal - the less reactive one is reduced and the more reactive one is oxidized. (aka a reaction in which both REDuction and OXidation occur)

Oxidizing agent:

Reducing agent:

 Thank you, Google :)

2.32 Understand oxidation and reduction as the addition and removal of oxygen respectively

Oxidisation is when something gains oxygen
Reduction is when something loses oxygen

(This is in terms of oxygen gain/loss)

OILRIG
(Oxidisation Is Loss - Reduction Is Gain)
This is in terms of electrons

Oxidation is when something gains ELECTRONS (use the dash in the acronym to remember, as electrons are negative - )

Reduction is when something loses ELECTRONS

2.31 Deduce the position of a metal within the reactivity series using displacement reactions between metals and their oxides, and between metals and their salts in aqueous solutions

If a more reactive metal (A) is placed into a test tube (or other container) with a metal oxide (B) dissolved in water, it will displace that metal (B).

If a less reactive metal (C) is placed into the test tube with the metal oxide (B), no reaction will occur.

This applies to the solution of their salts as well. From this, you can see which is more reactive. I.e. if chlorine displaces bromine, you can tell that chlorine is more reactive. :)

2.30 Describe how reactions with water and dilute acids can be used to deduce the following order of reactivity: potassium, sodium, lithium, calcium, magnesium, zinc, iron and copper

Potassium, sodium, calcium and lithium all react with water, and very vigorously with acid. Calcium has a much less vigorous reaction.

Magnesium, zinc and iron react with acid and very slowly with water. 

Copper reacts with neither (boring!).

By watching their chemical reactions, we can put them into order of reactivity:

• Potassium
• Sodium
• Lithium
• Calcium
• Magnesium
• Zinc
• Iron
• Copper

Potassium + water ;)

2.29 Understand that metals can be arranged in a reactivity series based on the reactions of the metals and their compounds: potassium, sodium, lithium, calcium, magnesium, aluminium, zinc, iron, copper, silver and gold

These metals, in order of reactivity are:

• Potassium
• Sodium
• Lithium 
• Calcium
• Magnesium
• Aluminium
• Zinc
• Iron
• Copper
• Silver
• Gold

2.15 Understand these displacement reactions as redox reactions.

When a more reactive halogen displaces a less reactive halogen, it's called a redox reaction. This is because an element has gained something and the other has lost something:

OIL RIG:
Oxidisation
Is
Loss (of electrons)
Reduction
Is
Gain (of electrons)

2.14 Describe experiments to demonstrate that a more reactive halogen will displace a less reactive halogen from a solution of one of its salts

A more reactive halogen will displace a less reactive one in a solution of its salt but only if the salt is dissolved in water or a gas.

Eg. if chlorine (Cl) is added to a solution of sodium bromide, it will form a solution of sodium chloride instead, and bromine will become a 'waste' product.

The balanced equation would be:

Cl₂ + 2NaBr → 2NaCl + Br₂

2.8 Explain the relative reactivities of the elements in Group 1 in terms of distance between the outer electrons and the nucleus.

Group 1 elements only have 1 electron in their outer shell. Electrons are held to the atom by the forces of attraction between the nucleus (which has protons, so a positive charge) and the electrons themselves (which have a negative charge). If there are more shells, the last electron (which will be lost if the metal reacts with anything)  is further away from the nucleus, so the forces of attraction between the nucleus and the electron are weaker. The more shells, the further away, and it's more reactive.

As you go down the group, they become more reactive because there is an increase in the number of shells :)

(Does this make sense or am I talking gibberish?)