Unit 2: Section 2 - Group 2 and Group 7 elements

Question 1

How does radius change down group 2?

Answer 1

Increases - extra electron shells are added

Question 2

How does first ionisation energy change down group 2?

Answer 2

Decreases
1. Each element has an extra electron shell as you go down
2. This extra inner shell shields the outer electrons from the attraction of the nucleus
3. The electrons are also further away from the nucleus, which greatly reduces the nucleus’ attraction
4. Both of these factors make it easier to remove outer electrons, resulting in a lower first ionisation energy

Question 3

How does reactivity change down group 2?

Answer 3

Increases
1. As you go down the group, first ionisation energy decreases due to the increasing atomic radius and shielding effect
2. When group 2 electrons react, they form +2 ions by losing electrons
3. The lower the first ionisation energy, the more reactive the element, as it’s easier to lose outer shell electrons

Question 4

How do melting points change down group 2?

Answer 4

Decreases
1. They have typical metallic structures with positive ions in a crystal structure surrounded by delocalised electrons
2. Down the group the metal ions get bigger but the number of delocalised electrons per atom doesn’t change and neither does the charge on the ion
3. The larger the ion and the further away the delocalised electrons are from the positive nuclei, the less attraction they feel
4. It takes less energy to break the bonds
5. Mg is an anamoly becsause the crystal structure changes

Question 5

What happens when group 2 elements react with water?

Answer 5

They are oxidised from a state of 0 to +2 froming M2+ ions - react with water to form a metal hydride and hydrogen

Question 6

What are the solubility trends in group 2?

Answer 6

• Compounds that contain singly charged negative ions increase n solubility down the group
• Compounds that contain doubly charged negative ion decease in solubility down the group
  OH- hydrOxides are mOre sOluble dOwn
  SO42- suLfates are Less soLuble down
  Magnesium is sparingly soluble and most sulfates are soluble in water but BaSO4 is insoluble

Question 7

What is the test for sulfate ions?

Answer 7

Acidify the solution with hydrochloric acid to remove any sulfates or carbonates which will also produce a white precipitate
Add acidified potassium dichromate (BaCl2) and sulfate ions are present, a white precipitate of barium sulfate is formed

Question 8

What are the uses of Group 2 elements?

Answer 8

Neutralising acids
1. Calcium hydroxides (Slaked lime Ca(OH)2) - agriculture to neutralise acid soils
2. Magnesium hydroxide (Mg(OH)2) - indigestion tablets as an antacid which removes excess stomach acid

Question 9

What is the equation for neutralisation for group 2 elements?

Answer 9

H+(aq) + OH-(aq) -> H2O(l)

Question 10

What is barium sulfate used for?

Answer 10

1. It is opaque to X-rays so is used in barium meals to help diagnose problems with the oesophagus, stomach or intestines
2. A patient swallows the barium meal, which is a suspension of barium sulfate
3. The barium sulfate coats the tissues, making them show up on the X-rays, showing the structure of the organs

Question 11

How is magnesium used in the extraction of titanium from its ore?

Answer 11

1. The main titanium ore, titanium (IV) oxide (TiO2) is first converted to titanium (IV) chloride (TiCl4) by heating it with carbon in a stream of chlorine gas
2. The titanium chloride is then purified by fractional distillation, before being reduced by magnesium in a furnace of about 1000°C

Question 12

What is the equation for the reduction of titanium with magnesium?

Answer 12

TiCl4(g) + 2MgCl(l) -> Ti(s) + 2MgCl2(l)

Question 13

How can you remove sulfur dioxide?

Answer 13

1. Burning fossil fuels to produce electricity also produces sulfur dioxide, which pollutes the atmosphere
2. The acidic sulfur dioxide can be removed from flue gases by reacting with an alkali - wet scrubbing
3. Powdered calcium oxide (lime, CaO) and calcium carbonate (limestome, CaCO3) can both be used
4. A slurry is made by mixing the calcium oxide or calcium carbonate with water which is then sprayed onto the flue gases
5. The sulfur dioxide reacts with the alkaline slurry and produces a solid waste product - calcium sulfate

Question 14

What are the equations for the reactions between calcium oxide/calcium carbonate with sufur dioxide?

Answer 14

CaO(s) + 2H2(l) +SO2(g) -> CaSO3(s) + 2H2(l)
CaCO3(s) + 2H2O (l) + SO2 (g) -> CaSO3(s) + 2H2O(l) + CO2(g)

Question 15

What are the main properties of the first 4 halogens?

Answer 15

Fluorine - F2 pale yellow gas
Chlorine - Cl2 green gas
Bromine - Br2 red-brown liquid
Iodine - I2 grey solid

Question 16

What is the trend in boiling points down group 7?

Answer 16

Increases - the size and relative mass of the molecules increases therefore the Van der Waals forces also increase in strength

Question 17

What is the trend in electronegativity down group 7?

Answer 17

Decreases - they are all highly electronegative but larger atoms attract electrons less than smaller ones as the electrons are further from the nucleus and are shielded by more electrons

Question 18

When does a halogen displace a halide?

Answer 18

When the halide ion is less reactive - a halogen will displace a halide from solution if the halide is below it in the periodic table

Question 19

What is the trend in oxidising ability in the halogens?

Answer 19

Less oxidising down the group as they are less reactive

Question 20

What are the displacement equations for the halogens?

Answer 20

Cl2(aq) + 2Br-(aq) -> 2Cl-(aq) + Br2(aq)
Cl - displace Br and I
Br - displace I
I - no reaction

Question 21

What can be used to help identify which halogen is present in a solution?

Answer 21

Displacement reactions
https://www.savemyexams.com/gcse/chemistry_combined-science/aqa/18/revision-notes/1-atomic-structure–the-periodic-table/1-2-the-periodic-table/1-2-6-group-7-the-halogens/
Learn the summary table of the displacememnt reactions of the halogens

Question 22

How is bleach made?

Answer 22

By mixing chlorine gas with cold, dilute, aqueous sodium hydroxide which forms sodium chlorate(I) soultion
2NaOH(aq) + Cl2(g) -> NaClO + NaCl(aq) + H2O(l)
Oxidation states: 0 +1 -1
Disproportionation reaction

Question 23

What are the uses of sodium chlorate(I) solution?

Answer 23

Water treatment, to bleach paper and textiles, cleaning

Question 24

How is chlorine used to treat water?

Answer 24

A disproportionation reaction occurs
End up with a mixture of chloride anf chlorate(I) ions
Cl2(g) + H2O(l) <-> 2H+(aq) + Cl-(aq) + ClO-(aq)
In sunlight chlorine can also decompose water to form chloride ions and oxygen
2Cl2(aq) + 2H2O(l) <-> 4H+(aq) + 4Cl-(aq) + O2(g)

Question 25

What are the advamtages of chlorine in water treatment?

Answer 25

• Kills disease-causing microorganisms
• Some chlorine persists in the water and prevents reinfection further down the supply
• Prevents the growth of algae, aliminationg bad tastes and smells, and removes discolouration caused by organic compounds

Question 26

What are the risks of using chlorine to treat water?

Answer 26

• Chlorine gas is very harmful if breathed in - irritates the respiratory system
• Liquid chlorine on the skin or eyes causes severe chemical burns
• Chlorine reacts with a variety of organic compounds in the water to form chlorinated hydrocarbons which are carcinogenic - Small risk compared to the risks of untreated water

Question 27

How does the trend in the reducing power of the halides change down the group?

Answer 27

Greater reducing power
1. the ions get bigger, so the electrons are further away from the positive nucleus
2. Extra inner electron shells, so a greater shielding effect
3. The further down, the easier it is to lose electrons

Question 28

What are the reactions between NaF or NaCl with H2SO4?

Answer 28

NaF(s) + H2SO4(l) -> NaHSO4(s) + HF(g)
NaCl + H2SO4(l) -> NaHSO4(s) + HCl)
1. HF or HCl is formed, which form wmf as the gas comes into contact with the moisture in the air
2. HF and HCl aren’t strong enough reducing agents to reduce the sulfuric acid
3. Not a redox reaction

Question 29

What is the reaction of NaBr with H2SO4?

Answer 29

NaBr(s) + H2SO4(l) -> NaHSO4(s) + HBr(g)
1. wmf of bromine gas
2. The HBr is a stronger reducing agent and reacts with H2SO4 in a redox reaction
2HBr(g) + H2SO4(l) -> Br2(g) + SO2(g) + 2H2O(l)
3. The reaction produces choking fumes of SO2 and orange fumes of Br2

Question 30

What is the reaction of NaI with H2SO4?

Answer 30

NaI(s) + H2SO4(l) -> NaHSO4(s) + HI(g)
2HBr(g) + H2SO4(l) -> Br2(g) + SO2(g) + 2H2O(l)
1. As HI is a very strong reducing agent, it keeps going and reduces the SO2 to H2S
2. Soild iodine is also formed by this reaction
6HI(g) + SO2(g) -> S(s) + 3I2(s) + 4H2O(l)
3. Sulfur is a yellow precipitate and 3I2 is a black solid
8HI(g) + H2SO(l) -> 4I2(s) + H2S(g) + 4H2O(l)
5. 3. H2S gas is toxic and smells of rotten eggs

Question 31

What is the test for halides?

Answer 31

1. Add dilute nitric acid to remove carbonate ions
2. Add a few drops of silver nitrate solution AgNO3(aq) - a precipitate is formed
   Ag+(aq) + X-(aq) -> AgX(s)
   Fluoride - no precipitate
   Chloride - white precipitate
   Bromide - cream precipitate
   Iodide - yellow precipitate
3. Add ammonia solution
   Chloride - dissolves in dilute NH3(aq)
   Bromide - dissolves in conc NH3(aq)
   Iodide - insoluble in conc NH3(aq)

Question 32

How do you test for Group 2 ions?

Answer 32

Flame Test
1. Dip a nichrome wire loop in conc HCl
2. Dip the loop into an unknown compound
3. Hold the loop in the clear blue part of the bunsen burner falme
4. Observe the colour change
Calcium - brick red
Strontium - red
Barium - pale green
Dilut NaOH
1. Add dropwise to a test tube containing the metal ion solution
2. Observe the precipitate
3. Add excess NaOH until it’s in excess
Magnesium - slight white precipitate - white precipitate
Calcium - slight white precipitate - slight white precipitate
Strontium - slight white precipitate - slight white precipitate
Barium - no change - no change

Question 33

How to test for ammonium ions?

Answer 33

1. Add hydroxide ions to a solution containing ammonium ions and gently heat which will react to form ammonia gas and water
   NH4+(aq) + OH-(aq) -> NH3(g) + H2O(l)
2. If ammonia gas is present, damp (gas can dissolve) red litmus paper will turn blue

Question 34

How can you test for sulfates?

Answer 34

Add dilute HCl to remove carbonates, and add barium chloride solution
Ba2+(aq) + SO42-(aq) -> BaSO4(s)
A white precipitate will form

Question 35

How to test for hydroxides?

Answer 35

1. Dip a piece of red litmus paper into the solution
2. If hydroxide ions are present, the paper will turn blue

Question 36

How to test for halides?

Answer 36

1. Add dilute nitric acid
2. Add silver nitrate to form silver halide
   Chloride - white precipitate
   Bromide - cream precipitate
   Iodide - yellow precipitate

Question 37

How to test for carbonates?

Answer 37

1. Add HCl, and the solution will fiz
   CO32-(s) + 2H+(aq) -> CO2(g) + H2O(l)
2. Test for carbon dioxide ny bubbling through limewater which will turn cloudy