4.1 groups 1, 2 & 7

Question 1

explain why ionisation energy decreases down group 2.

Answer 1

• each element down group 2 has an additional energy level compared to the element above.
• this additional energy level shields the outer electrons from the attraction of the nucleus.
• the electrostatic attraction between the nucleus and outer electrons decrease down the group, as the additional energy level causes the outer electrons to move further away from the nucleus.

Question 2

ionisation energy decreases down group 2. use this trend to explain the trend in the reactivity of group 2 elements.

Answer 2

• reactivity increases down group 2.
• most group 2 elements react by losing their outermost two electrons.
• the higher the first and second ionisation energies, the more likely group 2 elements are to lose these electrons.

Question 3

group 2 elements react with water, oxygen and chlorine. give the product(s) formed in each of these reactions, including their state symbols.

Answer 3

group 2 element + water → metal hydroxide (aq) + hydrogen (g)

group 2 element + oxygen → oxide (s)

group 2 element + chlorine → chloride (s)

Question 4

the oxides of group 2 elements react readily with water to form metal hydroxides, which dissolve. how do these dissolved OH⁻ ions affect the pH of the resulting solution?

Answer 4

OH⁻ ions make the resulting solution strongly alkaline.

Question 5

give the two oxides of group 2 which do not react with water, and explain why they are unable to react with water to form metal hydroxides.

Answer 5

beryllium oxide - does not react with water due to its small size and high ionisation energy compared to other elements in the group.

magnesium oxide - reacts very slowly with water due to its high lattice enthalpy.

Question 6

why do group 2 oxides form more strongly alkaline solutions down the group?

Answer 6

because the hydroxides become more soluble as you progress down the group.

Question 7

solubility trends of group 2 elements depend on the compound anion. explain the solubility trend of group 2 hydroxides, and how this is different from group 2 sulphates. provide an example for each.

Answer 7

• the solubility of group 2 hydroxides increases down the group as the compounds contain singly charged negative ions, such as OH⁻.
• in contrast, the solubility of group 2 sulphates decreases down the group as the compounds contain doubly charged negative ions, such as SO₄²⁻

Question 8

with reference to the ions involved, explain how distortion can affect the stability of group 2 carbonates and nitrates.

Answer 8

• the carbonate and nitrate ions are anions, and can be made unstable by the presence of a cation.
• the cation polarises the anion, which causes it to distort.
• the greater the distortion, the less stable the compound is.

Question 9

explain why large cations cause less distortion than small cations, and how this can affect the stability of group 2 carbonates and nitrates.

Answer 9

• large cations cause less distortion than small cations as they have a lower charge density.
• the larger the cations, the lower the charge density, so the less distortion caused and the more stable the carbonate / nitrate.

Question 10

give the thermal stability trends of group 1 and 2 compounds.

Answer 10

• group 1 carbonates are thermally stable - they do not decompose.
• group 1 nitrates decompose to form the nitrate and oxygen.
• group 2 carbonates decompose to form the oxide and carbon dioxide.
• group 2 nitrates decompose to form the nitrate, nitrogen dioxide, and oxygen.

Question 11

give one way in which you could test for the decomposition of a nitrate in the lab.

Answer 11

by measuring the time taken for a certain amount of oxygen to be produced, for example, by measuring how long it takes for a glowing splint to be relit.

Question 12

give one way in which you could test for the decomposition of a carbonate in the lab.

Answer 12

by measuring the time taken for a certain amount of carbon dioxide to be produced, for example, by measuring how long it takes for a solution of limewater to turn cloudy.

Question 13

give the flame colours produced during a flame test for the following group 1 elements: Li, Na, K, Rb, Cs.

Answer 13

• lithium - red flame.
• sodium - yellow/orange flame.
• potassium - lilac flame.
• rubidium - red/purple flame.
• caesium - blue/violet flame.

Question 14

explain how to carry out a flame test in the lab.

Answer 14

• mix a small amount of the element to be tested with a few drops of HCl.
• heat a piece of nichrome wire loop in a bunsen burner flame to sterilise the wire.
• dip the wire loop into the compound / acid mixture.
• hold the wire loop in a very hot flame and note the flame colour produced.

Question 14

give the flame colours produced during a flame test for the following group 2 elements: Mg, Ca, Sr, Ba.

Answer 14

• magnesium - bright white flame.
• calcium - red/orange flame.
• strontium - crimson flame.
• barium - pale green flame.

Question 15

with reference to electron activity, explain why group 1 and 2 elements burn with distinctive flame colours.

Answer 15

• the thermal energy absorbed from the flame causes the electrons to move to higher energy levels within the atom.
• these distinctive colours are produced when the electrons fall back down to lower energy levels, releasing the energy previously gained in the form of light.
• the colour of the flame produced is dependent on the wavelength of light released, which is determined by the difference in energy between the higher and lower energy levels within the atom.

Question 16

halogens in their natural state exist as what type of covalent molecules?

Answer 16

covalent diatomic molecules.

Question 17

explain why halogens have low solubility in water.

Answer 17

halogens have low solubility in water because they are non-polar molecules.

Question 18

halogens are oxidising agents. explain what is meant by an oxidising agent.

Answer 18

• an oxidising agent is a substance that oxidises another species by causing it to lose electrons.
• an oxidising agent gains electrons, so itself is reduced.

Question 19

explain why the reactivity of the halogens decreases down the group.

Answer 19

• the atoms of group 7 halogens become larger as you progress down the group.
• the outer electrons move further away from the nucleus, and are shielded from the positive attraction of the nucleus due to an increased number of inner electrons.
• this makes it more difficult for large atoms to attract the electron needed to form an ion, so reactivity decreases down the group.

Question 20

explain why the electronegativity of the halogens also decreases down the group.

Answer 20

• as you progress down group 7, there is an increase in the number of electron shells, and an increase in the distance between the nucleus and the bonding electrons.
• this makes it more difficult for an atom to attract an electron in a covalent bond, so the electronegativity decreases down the group.

Question 21

explain why the melting and boiling points of group 7 elements increase down the group.

Answer 21

• as you progress down the group, there is an increase in electron shells, and therefore electrons.
• this causes an increase in the London forces present between the halogen molecules.
• this increase in London forces means that more energy is required to overcome the strong intermolecular forces present, which causes the melting and boiling points of the halogens to increase down the group.

Question 22

the oxidising power of the halogens can be seen in their displacement reactions with halide ions. explain the trend in the oxidising power of the halogens down the group.

Answer 22

• the oxidising power of the halogens decreases down the group.
• as the halogen atoms become larger, they accept electrons less easily, which causes the oxidising power to become weaker.

Question 23

when a displacement reaction between a halogen and halide occurs, a colour change will be observed. give the colour change that would be observed if bromide is displaced and bromine is formed.

Answer 23

an orange colour change would be observed.

Question 24

give the colour change that would be observed if iodide is replaced and iodine is formed.

Answer 24

a brown colour change would be observed.

Question 25

give one way in which you could make these colour changes easier to see.

Answer 25

• you could make these colour changes easier to see by shaking the reaction mixture with an organic solvent, such as hexane.
• the halogen present will dissolve in the organic solvent, which settles out as a distinct layer above the aqueous solution.

Question 26

halogens can oxidise group 1 and 2 metals to produce what?

Answer 26

halide salts.

Question 27

give the three reactants which halogens can react with, resulting in a disproportionation reaction. provide an example for each.

Answer 27

• cold dilute alkali solutions (NaOH)
• hot dilute alkali solutions (KOH)
• water.

Question 28

halides are reducing agents. explain what is meant by a reducing agent.

Answer 28

• a reducing agent is a substance that reduces another species by causing it to gain electrons.
• a reducing agent loses electrons, so itself is oxidised.

Question 29

explain why the reducing power of the halides increases down the group.

Answer 29

• the reducing power of a halide depends on the attraction between the halide’s nucleus and the outer electrons.
• as you go down the group, this attraction becomes weaker due to the increase in the size of the ions, which results in the outer electrons moving further away from the nucleus.
• there is also a greater shielding effect, due to an increase in inner electron shells.
• the overall result is an increase in the reducing power of the halides as you progress down the group.

Question 30

all halides react with concentrated sulphuric acid to give what as a starting product?

Answer 30

a hydrogen halide.

Question 31

explain why the reaction between potassium chloride and sulphuric acid is not a redox reaction.

Answer 31

• potassium chloride is not a strong enough reducing agent to reduce the sulphuric acid.
• therefore, the oxidation states of both species remain the same, and the reaction is not a redox reaction.

Question 32

hydrogen halides are colourless gases. they can dissolve in water to produce misty fumes of acidic gas. damp litmus paper can be used to test if a hydrogen halide has dissolved. describe the colour change of the litmus paper upon coming into contact with a hydrogen halide.

Answer 32

hydrogen halides dissolved in water turn damp litmus paper from blue to red.

Question 33

the hydrogen halides can also react with ammonia to produce what colour fumes?

Answer 33

white fumes.

Question 34

explain one way in which you could test for the presence of a halide ion in the lab.

Answer 34

• add dilute nitric acid to the reaction mixture. this is to remove any ions which may interfere with the reaction.
• add silver nitrate solution to the reaction mixture. if a halide is present, a precipitate of the silver halide will from.

Question 35

give the colour of the precipitate formed when silver nitrate reacts with the following halides: fluoride, chloride, bromide, iodide.

Answer 35

• fluoride - no precipitate formed.
• chloride - white precipitate formed.
• bromide - cream precipitate formed.
• iodide - yellow precipitate formed.

Question 36

describe what you would observe if you added dilute ammonia solution to the following silver halide precipitates: silver chloride, silver bromide, silver iodide.

Answer 36

silver chloride - precipitate dissolves to give a colourless solution.

silver bromide - precipitate remains unchanged.

silver iodide - precipitate does not dissolve.