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A Level Chemistry > Elements of Group 1 and 2 - Inorganic chemistry > Flashcards

Elements of Group 1 and 2 - Inorganic chemistry Flashcards

(51 cards)

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1
Q

What is the bonding in the group 1 and 2 elements?

A

Metallic

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2
Q

What are the reasons for the trend in melting and boiling point and hardness of group 1 and 2 metals?

A

Metallic radius increases down the group due to the extra electron shells, so the metallic bonding weakens down the group, and so melting, boiling points and hardness all decrease down the group

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3
Q

Why are group 2 metals harder than group 1 metals and have higher melting and boiling points?

A

Number of valency electrons increases and metallic radius slightly decreases, so the strength of the metallic bonding increases

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4
Q

What are the reasons for the trend in ionisation energies down group 1 and 2?

A

The increase in nuclear charge is outweighed by the increased distance between the electron being lost and the nucleus, and the increased shielding effect of complete shells of electrons between the electron being lost and the nucleus

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5
Q

What is lithium’s flame test colour?

A

(Carmine) red

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6
Q

What is sodium’s flame test colour?

A

Intense yellow

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7
Q

What is potassium’s flame test colour?

A

Lilac

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8
Q

What is magnesium’s flame test colour?

A

No colour

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9
Q

Why does magnesium’s flame test not produce a colour?

A

The energy emitted has a wavelength outside the visible spectrum

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10
Q

What is calcium’s flame test colour?

A

Brick red

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11
Q

What is strontium’s flame test colour?

A

Red/Crimson

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12
Q

What is barium’s flame test colour?

A

(Apple) green

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13
Q

What is rubidium’s flame test colour?

A

Red

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14
Q

What is caesium’s flame test colour?

A

(Sky) blue

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15
Q

What is beryllium’s flame test colour?

A

No colour

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16
Q

Why do flame tests work?

A

Electrons are excited to a higher energy level by the Bunsen burner flame, and as they fall back down to their ground state (they are unstable in the higher energy state), they emit energy in the form of visible light

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17
Q

How should a flame test be carried out?

A

1) Use a nichrome wire (very unreactive and so cannot give out a flame colour)
2) Dip the wire in concentrated hydrochloric acid (out it in the flame to check if any colour is being produced, keep cleaning it until there isn’t one)
3) Grind up any solid that isn’t already a powder
4) Dip the wire in the solid’
5) Put the wire into the flame so a colour is produced

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18
Q

What colour does magnesium burn in oxygen?

A

Intense white light

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19
Q

What colour does calcium burn in oxygen?

A

Brick red flame

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20
Q

What colour does strontium burn in oxygen?

A

Dark red flame

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21
Q

What colour does barium burn in oxygen?

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A

Pale green flame

22
Q

What compound do all group 2 metals form when they react with chlorine?

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A

MCl2

Beryllium forms a covalent, anhydrous chloride, all others produce ionic chlorides

23
Q

Beryllium’s reaction with water

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A

There is no reaction

24
Q

Magnesium’s reaction with water

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A

Magnesium reacts very slowly with cold water, but burns when heated in steam
Mg (s) + H2O (g) -> MgO (s) + H2 (g)

25
Calcium, strontium and barium's reactions with water
React rapidly with cold water | Ba (s) + 2H2O (g) -> Ba(OH)2 (aq) + H2 (g)
26
How does the reactivity of group 2 elements change down the group?
It increases
27
Barium reacting with oxygen
2Ba (s) + O2 (g) -> 2BaO (s)
28
Beryllium oxide with water
No reaction
29
Magnesium oxide with water
MgO (s) + H2O (l) -> Mg(OH)2 (s) | Reacts slowly and incompletely to form a slightly alkaline suspension of magnesium hydroxide
30
Calcium oxide with water
CaO (s) + H2O (l) -> Ca(OH)2 (s) | Reacts very exothermically to form an alkaline suspension of calcium hydroxide (limewater)
31
Strontium and barium with water
Sr^2+ + 2OH- -> Sr(OH)2 (aq) (same for barium)
32
Magnesium oxide and sulfuric acid (all group 2 metal oxides + acids)
MgO (s) + H2SO4 (aq) -> MgSO4 (aq) + H2O (l) (Base + acid -> salt + water) MgO (s) + 2H+ (aq) -> Mg^2+ (aq) + H2O (l) (sulfate ions do not change so are not included in the ionic equation)
33
Calcium hydroxide and nitric acid
Ca(OH)2 (aq) + 2HNO3 (aq) -> Ca(NO3)2 (aq) + 2H2O (l) | 2OH- (aq) + 2H+ (aq) -> 2H2O (l)
34
Beryllium oxide and hydrochloric acid
BeO (s) + 2HCl (aq) -> BeCl2 (aq) + H2O (l) | BeO (s) + 2H+ -> Be^2+ (aq) + H2O (l)
35
Why does beryllium oxide react differently?
It is an amphoteric oxide and so it reacts as both a base and an acid
36
Reactions of group 1 oxides and hydroxides
All group 1 oxides react with water to give hydroxides. Group 1 hydroxides all react with acids to form salts which are soluble in water, producing alkaline solutions with a high pH
37
What is the trend in solubility of the group 2 sulphates?
As you go down the group, the solubility of the group 2 sulfates decreases
38
What is the trend in solubility of group 2 hydroxides?
As you go down the group, the solubility of the group 2 hydroxides increases
39
What determines solubility?
Energy is required to break up the ionic lattice (the reverse lattice energy) and energy is given out as water molecules are attracted to the ions (hydration energy), and the solubility depends on the balance between the endothermic reverse lattice energy and the exothermic hydration energy
40
Testing for sulfate ions
Add barium chloride or barium nitrate solution to the solution being tested, and a white ppt will form Barium carbonate is also insoluble, so dilute hydrochloric acid is added to convert any barium carbonate to barium chloride Ba^2+ (aq) + SO4^2- -> BaSO4 (s)
41
How to distinguish between magnesium ions and barium ions
Adding sulfate ions: magnesium sulfate is soluble, but barium sulfate is not, so a white precipitate forms with barium sulfate, but not with magnesium sulfate Adding hydroxide ions: Opposite to sulfate - Mg(OH)2 is insoluble and would form a white ppt
42
What is the trend of thermal stability in group 1 and 2 carbonates?
Increases down the group, become more thermally stable and require more heat to decompose
43
Thermal decomposition of group 1 carbonate
Li2CO3 (s) -> Li2O (s) + CO2 (g) | Only lithium carbonate decomposes in group 1, to produce lithium oxide and carbon dioxide
44
Thermal decomposition of group 2 carbonates
MgCO3 (s) -> MgO (s) + CO2 (g) All group 2 carbonates decompose on heating - beryllium carbonate is so unstable it does not exist at room temperature - barium carbonate requires extremely strong heating before it is decomposed Group 2 carbonates decompose to from metal oxide and carbon dioxide
45
What is the trend of thermal stability in group 1 and 2 nitrates?
Down the group, they become more thermally stable
46
Thermal decomposition of lithium nitrate
4LiNO3 -> 2Li2O + 4NO2 + O2 (more like a group 2 nitrate) | Nitrogen dioxide is visible as a brown gas. Its decomposition is different to all other group 2 nitrates
47
Thermal decomposition of other group 1 nitrates
2NaNO3 -> 2NaNO2 + O2 | Very strong heating causes them to melt, giving off oxyden and leaving a molten nitrite
48
Thermal decomposition of group 2 nitrates
2Ca(NO3)2 -> 2CaO + 4NO2 + O2 | They all decompose to form the metal oxide, nitrogen dioxide (visible as a brown gas) and oxygen
49
Explanation for the trend in thermal stability
From group 1 to group 2, charge density increases, and down the group charge density decreases as the ions get bigger. Ions with lower charge density polarise the anions less
50
Method for investigating the thermal decomposition of carbonates
Heat a known mass of carbonate in a sidearm boiling tube and pass the gas produced through limewater. Time how long it takes for the first permanent cloudiness to appear. Repeat for different carbonates using the same moles of carbonate, same volume of limewater, same bunsen flame and height of the tube above the flame
51
Method for investigating the thermal decomposition of nitrates
Set up test tube with the nitrate being tested in a clamp stand, with a bung. Start a timer when heating begins. Test for oxygen at regular intervals and look for the first signs of a brown gas, and record the time it takes for it to appear

A Level Chemistry (25 decks)

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