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Unknown chloride: MCl (M= metal)
Method

1. Weigh out ~0.3g of MCl - RECORD THE MASS

2. Dissolve into ~100ml of water

3. Add 3ml of 6M nitric acid. Stir

4. Add 20ml of silver nitrate

5. Add 5ml of silver nitrate

6. Place in 60 degree water bath for 10 mins

7. Leave to cool

8. Measure mass of filter paper and fold

9. Pour ppt solution into funnel

10. Wash ppt using wash solution and get solid out of beaker

11. test liquid with ~3 drops of HCl (2M)

12. Wash again- keep repeating until all solids are in the filter

13. Record the mass of the ppt after it dried and subtract the mass of the filter paper to find the overall

Our data
Mass of MCl - 0.330g
Mass of filter paper - 1.460g
Mass of ppt - 5.835g
5.240g
3.300g

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Chlorine

Added to water to make sure it’s safe to drink

Kills bacteria and viruses

Can prevent outbreaks of disease

Fluoride

Added to water to reduce tooth decay

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Halogens in metal halides always have an oxidation number of -1

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Add dilute aqueous ammonia to precipitates done in test 1 if they are hard to tell appart as this will distinguish them.
Chloride is soluble therefore when ammonia is added, a colourless solution is formed. Bromide and Iodide will remain unchanged.

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A precipitation reaction with silver nitrate solution and dilue nitric acid.
Silver Nitrate= AgNO3
Ag⁺(aq) + X⁻(aq) —> Ag⁺X⁻(s)
Example
Ag+  + Br-  —> Ag⁺Br⁻

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The ability to cause oxidation (to make other elements lose electrons)

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X = any halogen

All halogens exist as X₂ = diatomic molecules

As we go down group 7, volatility increases.

As we go down group 7, the bigger molecules contain more electrons therefore van der Waals' forces increase.

As we go down group 7, the oxidising power decreases and we see the trend in reactivity shown.

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Oxidising agents accept electrons, becoming reduced in the process.

Reducing agents give electrons, becoming oxidised in the process. - RAGE - Reducing Agents Give Electrons.

Example
Write half equations to show fluorine acting as an oxidising agent and lithium acting as a reducing agent and add to show overall reactions.

F2 + 2e⁻  —> 2F-
2Li   —> Li+ + 2e-
————————————————
F2 + 2Li + 2e⁻  —> 2F- + Li+ + 2e⁻
F₂ + 2Li —> 2Li⁺F-

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Example
Ionic Equation
Cl₂ + 2Br⁻ —> 2Cl- + Br2
Chlorine + Bromide —> Chloride + Bromine
Chemical Equation
Cl2(aq) + 2KBr(aq) —> 2KCl(aq) + Br2(aq)
K is the spectator ion

Chlorine has gained electrons from the bromide ions because chlorine is a stronger oxidising agent

Chlorine is more reactive than bromine therefore it can displace bromine from solution

Chlorine’s oxidation number has decreased therefore it is a reduction

Bromine’s oxidation number has increased therefore it is an oxidation

Observations
Bromine Water
KCl- No reaction
KBr- No reaction
KI- Turned yellow

Br2 + 2KI —> 2KBr + I₂

Br 0–> -1 reduction

I -1–>0 oxidation

K +1–>+1 unchanged

Chlorine Water
KCl - No reaction
KBr- turned yellow

Cl2 + 2KBr —> 2KCl + Br₂

Cl 0–>-1 reduction

Br -1–>0 oxidation

K +1–>+1 unchanged

KI - turned orange

Cl2 + 2KI —> 2KCl + I₂

Cl 0–>-1 reduction

I -1–>0 oxidation

K +1–>+1 unchanged

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Group 1 and 2 hydroxides are basic. A base dissolves into water to form an alkali.

The oxides of group 1 and 2 metals also have basic properties. This means that they react with an acid to form a salt and water.
Base + Acid --> Salt + Water

Example
CaO + HCl --> CaCl2 + H₂0

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Elements in groups 1 and 2 show signs of being more reactive as we go down the group.

Reaction between group 1/2 elements and water will always produce a metal hydroxide and hydrogen gas.
Example
2Na + 2H₂O --> 2NaOH + H2

Trend in thermal stability
The thermal stability of group 2 hydroxides and carbonates increases down the group.

Trend in solubility
In group 2 the hydroxides become more soluble as you go down the group and the sulphates become less soluble.

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Lithium - Red
Sodium - Golden Yellow
Potassium - Lilac
Calcium - Brick Red
Strontium - Crimson
Barium - Apple Green
Magnesium - No colour

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Group 1 - Alkali Metals
Group 2 - Alkalie Earth Metals
Group 7 - The Halogens
Group 8 - The Noble Gases

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Melting temperatures increase across a period until group 4 and then decrease.

For metals (eg. Group 1), melting temperatures decrease down the group

For the elements of group 7, melting temperatures increase down the group due to increasing van der Waals forces.

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Electronegativity decreases down a group because more shells are added which increases shielding and the distance of outer electrons from their nucleus.

Electronegativity increases across a period because nuclear charge increases. THere is a larger ‘pull’ on bonding electrons.

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First IE decreases down a group due to increased shielding and distance from the nucleus

First IE increases across a period because protons are being added which increases nuclear charge, making the electrons harder to lose.

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Oxidation - Loss of electrons - If an atom is oxidised, it’s oxidation number will go up.
Reduction - Gain of electrons - If an atom is reduced, its oxidation number will go down.

Remember OILRIG
Oxidation Is Lost Reduction Is Gain.

Example
Fe3+ + I-  —> Fe2+ + 1/2 I2.
Fe is a reduction, as its oxidation number has reduced from 3+ to 2+
I is an oxidation as its oxidation number has increased from -1 to 0

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Elements are grouped together in order of their atomic number

The group number indicates the number of valency (outer) electrons.

The period number indicates the number of electron shells

The electron structures of the elements relate to their position in the s, p and d blocks