Elements of Group 1 and 2 - Inorganic chemistry Flashcards
What is the bonding in the group 1 and 2 elements?
Metallic
What are the reasons for the trend in melting and boiling point and hardness of group 1 and 2 metals?
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
Why are group 2 metals harder than group 1 metals and have higher melting and boiling points?
Number of valency electrons increases and metallic radius slightly decreases, so the strength of the metallic bonding increases
What are the reasons for the trend in ionisation energies down group 1 and 2?
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
What is lithium’s flame test colour?
(Carmine) red
What is sodium’s flame test colour?
Intense yellow
What is potassium’s flame test colour?
Lilac
What is magnesium’s flame test colour?
No colour
Why does magnesium’s flame test not produce a colour?
The energy emitted has a wavelength outside the visible spectrum
What is calcium’s flame test colour?
Brick red
What is strontium’s flame test colour?
Red/Crimson
What is barium’s flame test colour?
(Apple) green
What is rubidium’s flame test colour?
Red
What is caesium’s flame test colour?
(Sky) blue
What is beryllium’s flame test colour?
No colour
Why do flame tests work?
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
How should a flame test be carried out?
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
What colour does magnesium burn in oxygen?
Intense white light
What colour does calcium burn in oxygen?
Brick red flame
What colour does strontium burn in oxygen?
Dark red flame
What colour does barium burn in oxygen?
Pale green flame
What compound do all group 2 metals form when they react with chlorine?
MCl2
Beryllium forms a covalent, anhydrous chloride, all others produce ionic chlorides
Beryllium’s reaction with water
There is no reaction
Magnesium’s reaction with water
Magnesium reacts very slowly with cold water, but burns when heated in steam
Mg (s) + H2O (g) -> MgO (s) + H2 (g)
Calcium, strontium and barium’s reactions with water
React rapidly with cold water
Ba (s) + 2H2O (g) -> Ba(OH)2 (aq) + H2 (g)
How does the reactivity of group 2 elements change down the group?
It increases
Barium reacting with oxygen
2Ba (s) + O2 (g) -> 2BaO (s)
Beryllium oxide with water
No reaction
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
Calcium oxide with water
CaO (s) + H2O (l) -> Ca(OH)2 (s)
Reacts very exothermically to form an alkaline suspension of calcium hydroxide (limewater)
Strontium and barium with water
Sr^2+ + 2OH- -> Sr(OH)2 (aq) (same for barium)
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)
Calcium hydroxide and nitric acid
Ca(OH)2 (aq) + 2HNO3 (aq) -> Ca(NO3)2 (aq) + 2H2O (l)
2OH- (aq) + 2H+ (aq) -> 2H2O (l)
Beryllium oxide and hydrochloric acid
BeO (s) + 2HCl (aq) -> BeCl2 (aq) + H2O (l)
BeO (s) + 2H+ -> Be^2+ (aq) + H2O (l)
Why does beryllium oxide react differently?
It is an amphoteric oxide and so it reacts as both a base and an acid
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
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
What is the trend in solubility of group 2 hydroxides?
As you go down the group, the solubility of the group 2 hydroxides increases
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
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)
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
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
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
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
What is the trend of thermal stability in group 1 and 2 nitrates?
Down the group, they become more thermally stable
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
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
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
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
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
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
