Group 1 and 2 elements

Question 1

Define first ionisation energy and state and explain the trend in first ionisation energies down groups 1 and 2.

Answer 1

Ionisation energy is the energy needed to remove an electron from one mole of an atom in its gaseous state.

First ionisation energies decrease as you go down groups 1 and 2.

This is because:

1. As you go down the group, the number of electron shells increases so there is more shielding between the outermost electron and the nucleus.
2. This means the electron shells are further away from the nucleus, have less attraction towards the nucleus and have more energy.
3. This makes the energy required to remove them decrease.

Question 2

Give the products and observations for reactions of group 1 elements with:

1. Oxygen
2. Chlorine (gas)
3. Water

Answer 2

1. Oxygen
   Forms Metal oxide.
   They tarnish when exposed to oxygen, which means they form a dull dark layer on their surface when oxidised.
   The reaction is vigorous and cannot be done in the lab.
   Reactivity increases down the group.

React in excess oxygen to form peroxides.

2. Chlorine (gas)
   Forms Metal chloride.
   Reactivity increases down the group.
3. Water
   Forms Metal hydroxide and Hydrogen gas.
   Effervescence is seen (due to formation of Hydrogen gas) and the metal dissolves in the water, a colourless solution forms (with lithium).
   Reactivity increases down the group.

Question 3

Give the products and observations for reactions of group 2 elements with:

1. Oxygen
2. Chlorine (gas)
3. Water (cold and steam)

Answer 3

1. Oxygen
   Forms Metal oxide.
   With magnesium- bright flame and formation of a white solid.

*Metal needs to be heated in order to take place because it is very slow without heating.
Strontium and barium can form peroxide (MO2 instead of 2MO)

2. Chlorine (gas)
   Forms Metal chloride.
   Reactivity increases down the group.
3. Water (cold)
   Forms Metal hydroxide and hydrogen.
   With magnesium slow, with barium and strontium very vigorous with effervescence.
   Metal hydroxide is an alkaline solution, so g2 elements are called alkaline earth metals.

Steam
Forms Metal oxide and hydrogen.
Magnesium burns with a bright flame and forms magnesium oxide, a white powder as well as hydrogen gas.

Question 4

Why is hydrogen burned as it leaves the tube in the reaction of magnesium with steam?

Answer 4

This is because hydrogen is highly flammable, and needs to be burned to prevent it from escaping into the laboratory.

Question 5

State and explain the trend in reactivity down groups 1 and 2.

Answer 5

Reactivity increases down the group.

This is because:

1. The increased shielding down the group causes the outermost electrons to be further away from the nucleus.
2. The outermost electrons have more energy as you go down the group, and this makes them easier to lose.

Question 6

State the trend in solubility of:

1. Group 2 hydroxides
2. Group 2 Sulfates

Answer 6

1. Group 2 hydroxides
   Solubility increases down the group.
   This means the alkaline solutions formed have a higher pH as you go down the group.

Eg. Magnesium hydroxide is the least soluble so it forms the lowest alkaline pH out of the rest.

2. Group 2 sulfates
   Solubility decreases down the group.

Barium sulfate is the least soluble.
So barium chloride is used to test for sulfate ions. It forms barium sulfate, which is a white precipitate.

Question 7

Describe the uses of Barium ions in:

1. Testing for sulfate ions
2. Barium meals

Answer 7

1. Testing for sulfate ions
   Barium nitrate/chloride can be used to test for sulfate ions. This forms a white precipitate.

To prevent Barium carbonate from forming, H+ ions must be added to prevent this. So dilute hydrochloric/nitric acid are added.

Equation:
Ba(NO3)2(aq) + Na2SO4(aq) —> BaSO4(s)+ 2NaNO3 (aq)

Ionic equation:
Ba2+ (aq) + SO4 2- (aq) —> BaSO4(s)

2. Barium meals
   Barium meals contain barium sulfate, which is non-poisonous to patients because it is insoluble.

Barium ions are present but they are not free to move so they cannot be of any harm.

Due to the dense white solid, soft tissues can be seen more clearly in X-ray images for patients that are given barium meals before the X-ray.

Question 8

Define thermal stability and explain the trend in thermal stability of the nitrates and carbonates of elements of group 1.

Answer 8

Thermal stability is the term that indicates how stable a compound is when it is heated. (whether it decomposes or not)

Thermal stability in group 1 increases down the group.

This is because:

1. As you go down the group, reactivity of the elements increases.
2. This means that the elements form more stable compounds so more heat energy is required for them to undergo thermal decomposition.

At the top form less stable compounds because:

1. Ionic radius increases down the group for the same overall charge, which means smaller ions have a higher charge density.
2. Smaller ions are better able to polarize the negative carbonate and nitrate ions, the more polarized the negative ions are, the less heat energy is needed to separate the ions.
3. So smaller ions form less stable compounds.

Trend:
Lithium carbonate and Lithium nitrate decompose greatly. Nitrates of any other element below it decomposes less.

Carbonates of any other element below it do not decompose.

Question 9

Why do carbonates and nitrates decompose and not melt?

Answer 9

1. The larger, more complex nitrate ion can change into smaller more stable nitrite or oxide ions by decomposing or releasing oxygen and/or nitrogen dioxide gas.
2. The larger, more complex carbonate ion can change into smaller more stable oxide ions by decomposing or releasing carbon dioxide gas.
3. A smaller cation will form a more stable carbonate/nitrate and a bigger cation will form less stable. The bigger the cation, the more likely the carbonate/nitrate compound is to decompose.

Question 10

Explain the trend in thermal stability of the nitrates and carbonates of elements of group 2.

Answer 10

Same explanation as before, except group 2 elements always decompose. This is because they have a higher charge than group 1 elements. Lithium decomposes the same as them because it is very small.

Question 11

State the equations and observations for decomposition of Lithium nitrate, Magnesium nitrate and Caesium nitrate.

Answer 11

1. Lithium nitrate
   Decomposes like group 2 elements.

4LiNO3 —> 2Li2O + 4NO2 + O2

2. Magnesium nitrate

2Mg(NO3)2 —> 2MgO + 4NO2 + O2

3. Caesium nitrate
   (less decomposition)

2CsNO3 —> 2CsNO2 + O2

General:
More decomposition
Metal nitrate —> Metal oxide + Nitrogen dioxide + Oxygen

Observation: Brown fumes (NO2)

Less decomposition
Metal nitrate —> Metal nitrite + Oxygen

Observation: No brown fumes

Question 12

State the equations and observations for decomposition of Lithium carbonate, Magnesium carbonate and Caesium carbonate.

Answer 12

1. Lithium carbonate
   More decomposition

Li2CO3 —> Li2O + CO2

2. Magnesium carbonate
   More decomposition

MgCO3 —> MgO + CO2

Observation: Colourless gas given off (CO2)
Carbonate and oxide are both white solids.

Not much of an observation, CO2 test can be done.

3. Caesium carbonate
   No decomposition except at very high temperatures.

Question 13

Describe how a flame test is done.

Answer 13

1. Light a bunsen burner in a fume cupboard.
2. Add a few drops of HCl to the metal so it can dissolve. This is done to form metal chlorides, and as chlorides are more volatile they will give better results.
3. Dip a clean nichrome wire into the solution to obtain a sample.
4. Hold the end of the wire to the flame and observe the colour.

Question 14

State the colours formed by the metals and explain the formation of characteristic flame colours by Group 1 and 2 compounds in terms of electron transitions

Answer 14

Group 1
Lithium: Red
Sodium: Yellow/orange
Potassium: Lilac
Rubidium: Red/purple
Caesium: Blue/violet

Group 2
Beryllium: no colour
Magnesium: no colour
Calcium: Brick red
Strontium: Crimson red
Barium: Apple green

Explanation:

• Electrons can absorb energy and move to a higher energy level.
• When it moves to a higher energy level, it is described as its ‘excited state’. Electrons move to an excited state during a flame test.
• Immediately after this, the electrons return back to their ground state. This releases energy.
• Some of this energy is in the form of light.
• The colour produced corresponds to the wavelength of this light emitted.
• Electrons that do not release energy in the form of visible light will not give a result in the flame test, eg beryllium.

Question 15

Describe the tests for CO2 and Oxygen.

Answer 15

1. CO2

It is passed through limewater/ saturated aqueous Ca(OH)2.
It turns milky/cloudy and a white precipitate forms as a result if CO2 is present.

CO2 + Ca(OH)2 —> CaCO3 + H2O

2. Oxygen

Bring a glowing splint close to a test tube with the unknown gas. If oxygen is present, the splint will relight.

Question 16

LO:
Know the reactions of:

1. Carbonate ions, CO3 2-, and hydrogencarbonate ions, HCO3- , using an aqueous acid to form carbon dioxide (and testing the gas with limewater)
2. Ammonium ions, NH4+, using sodium hydroxide solution and warming to form ammonia (and testing with litmus and HCl fumes)

Answer 16

1. When an aqueous acid such as HCl is added to them, effervescence can be seen and CO2 gas is released.

Reaction:

XCO3 + 2HCl —> XCl2 + CO2 + H2O

2. Ammonium ions do not produce a colour in the flame test. So, to test for ammonium ions NaOH is added to the metal.

Ionic equation:
NH4+ + OH- —> NH3 + H2O

Then, it is warmed which releases ammonia gas.
This gas can be tested by using damp red litmus paper (turns blue).

HCl fumes can also be passed which react with NH3 to form a white ammonium chloride smoke.

Reaction:
NH3 + HCl —> NH4Cl