Group 7 - Halogens

Question 1

Which elements make up Group 7 of the periodic table?

Answer 1

The halogens in Group 7 are:

• Fluorine (F)
• Chlorine (Cl)
• Bromine (Br)
• Iodine (I)
• Astatine (At)

Question 2

What is the molecular formula of halogens?

Answer 2

Halogens form diatomic molecules with the formula X₂ (e.g., F₂, Cl₂), where X represents the element.

Question 3

What are the appearances of halogens at room temperature?

Answer 3

Fluorine (F₂): Pale yellow gas
Chlorine (Cl₂): Green gas
Bromine (Br₂): Red-brown liquid
Iodine (I₂): Grey solid

Question 4

How do the boiling points of halogens change down Group 7?

Answer 4

The boiling points increase down Group 7 because:

• The molecules get larger and have a higher relative mass.
• This results in stronger induced dipole-dipole forces between molecules.
• more energy is required to overcome these forces

Question 5

How does electronegativity change down Group 7?

Answer 5

it decreases because:

• Nuclear charge increases protons but:
• Atomic radius increases so outer electrons are further away from the nucleus
• Electron shielding increases reduces nuclear attraction

Question 6

What do halogens act as?

Answer 6

oxidising agents as they gain an electron, causing another substance to lose an electron

Question 7

How does the reactivity of halogens change down Group 7?

Answer 7

It decreases because:

• Atomic radius increases
• Outer electrons are farther from the nucleus and experience more shielding.g
• The electrostatic attraction between the outer electrons and the nucleus becomes weaker.
• It becomes harder for larger halogens to attract the electron needed to form a negative ion

Question 8

How does the oxidising power of halogens change down Group 7?

Answer 8

Oxidising power decreases down Group 7 because:

• The halogens get larger and experience more shielding.
• This makes it harder for larger halogens to remove electrons
• Fluorine is the strongest oxidising agent, and iodine is the weakest.

Question 9

What happens in halogen displacement reactions?

Answer 9

• A more reactive halogen displaces a less reactive halide ion from solution. - The displacing halogen is reduced to form the halide ion.
  -The displaced halide is oxidised to form the halogen molecule.

Question 10

What is the rule for halogen displacement reactions?

Answer 10

A halogen will displace any halide ion below it in Group 7

Question 11

What are the displacement reactions for halogens and halides?

Answer 11

• Chlorine (Cl₂) displaces bromide (Br⁻) and iodide (I⁻)
• Bromine (Br₂) displaces iodide (I⁻):
• Iodine (I₂) does not displace any halides, so no reaction

Question 12

What colour changes occur during halogen displacement reactions?

Answer 12

• Bromine displaces bromide, turning the solution yellow.
• Iodine displaces iodide, turning the solution orange/brown.
  If no reaction occurs, the solution remains colourless.

Question 13

What are halide ions?

Answer 13

negatively charged forms of halogens with a 1- charge

Question 14

How do halide ions react in redox reactions?

Answer 14

They lose electrons to form neutral halogen molecules in oxidation reactions

Question 15

What happens to the oxidation number of halide ions during oxidation?

Answer 15

It increases from -1 to 0, meaning the halide ion is oxidised

Question 16

Why do halide ions act as reducing agents?

Answer 16

They donate electrons to another substance, causing its reduction

Question 17

How does the reducing power of halide ions change down Group 7?

Answer 17

it increases because:

• Ionic radius increases
• More electron shielding weakens nuclear attraction
• Weaker electrostatic attraction makes it easier to lose electrons

Question 18

Why is fluoride the weakest reducing agent?

Answer 18

• Smallest ionic radius
• Less shielding, making electron loss harder.

Question 19

Why is iodide the strongest reducing agent?

Answer 19

• Largest ionic radius
• More shielding

Question 20

What happens when halide ions react with concentrated sulfuric acid?

Answer 20

• Initially, hydrogen halide gases are formed
• Further reactions depend on the reducing power of the halide ion

Question 21

How do fluoride and chloride react with sulfuric acid?

Answer 21

• Only acid-base reactions occur—no redox.
• HF (g) or HCl (g) gas is formed.

Question 22

Observations and equations of fluoride and chloride with sulfuric acid

Answer 22

Misty white fumes of HF or HCl gas.

• NaF (s) + H₂SO₄ (aq) → NaHSO₄ (aq) + HF (g)
• NaCl (s) + H₂SO₄ (aq) → NaHSO₄ (aq) + HCl (g)

Question 23

Why Do Fluoride and Chloride Not Undergo Redox?

Answer 23

Their reducing power is too weak to reduce sulfuric acid

Question 24

How does bromide react with sulfuric acid?

Answer 24

• initial reaction: forms HBr gas
• Further reaction: Bromide reduces H₂SO₄ to SO₂, forming Br₂ vapour.

Question 25

observations of bromide and sulfuric acid

Answer 25

- Misty white fumes of HBr gas. - Orange bromine vapour. - Choking SO₂ gas.

Question 26

How does iodide react with sulfuric acid?

Answer 26

- Initial reaction: Forms HI gas (acid-base reaction). - Further reactions: Iodide reduces sulfur all the way from +6 to -2, forming SO₂, solid sulfur (S), and H₂S gas.

Question 27

observations and equations of iodide with sulfuric acid

Answer 27

- NaI (s) + H₂SO₄ (aq) → NaHSO₄ (aq) + HI (g) - 2HI (g) + H₂SO₄ (aq) → I₂ (g) + SO₂ (g) + 2H₂O (l) - 6HI (g) + SO₂ (g) → H₂S (g) + 3I₂ (s) + 2H₂O (l) - Misty white fumes of HI gas. - Violet iodine vapour. - Yellow solid sulfur. - Rotten egg smell of H₂S gas

Question 28

What is a disproportionation reaction?

Answer 28

A reaction where the same element is both oxidized and reduced in the same reaction.

Question 29

Give an example of a disproportionation reaction involving chlorine.

Answer 29

When chlorine reacts with water: Cl₂(g) + H₂O(l) → HCl(aq) + HClO(aq) ## Footnote Chlorine is oxidized from 0 to +1 in HClO and reduced from 0 to -1 in HCl.

Question 30

How does chlorine disinfect water?

Answer 30

Chlorine reacts with water in a disproportionation reaction: Cl₂(g) + H₂O(l) → HCl(aq) + HClO(aq) HClO dissociates into ClO⁻ ions, which kill bacteria: HClO → H⁺ + ClO⁻

Question 31

What happens to chlorine in sunlight?

Answer 31

Chlorine reacts with water to form hydrochloric acid and oxygen: Cl₂(g) + 2H₂O(l) → 2HCl(aq) + ½O₂(g) ## Footnote This removes the disinfectant (ClO⁻), so pools need regular chlorine addition.

Question 32

What are the benefits of using chlorine in water treatment?

Answer 32

* Kills disease-causing microorganisms * Prevents reinfection further down the water supply * Removes bad tastes, smells, and discoloration

Question 33

What are the risks of using chlorine in water treatment?

Answer 33

* Chlorine gas causes severe irritation if inhaled * Reacts with organic compounds in water to form carcinogenic chlorinated hydrocarbons

Question 34

What are some alternatives to chlorine for water disinfection?

Answer 34

* Ozone (O₃): Strong disinfectant but expensive and decays quickly * Ultraviolet (UV) light: Damages microbial DNA but less effective in cloudy water

Question 35

How is bleach made using chlorine?

Answer 35

Chlorine reacts with cold, dilute sodium hydroxide to form sodium chlorate(I): Cl₂(g) + 2NaOH(aq) → NaClO(aq) + NaCl(aq) + H₂O(l)

Question 36

Why is the bleach reaction a disproportionation reaction?

Answer 36

Chlorine is oxidized from 0 to +1 in NaClO and reduced from 0 to -1 in NaCl.

Question 37

What is the active ingredient in household bleach?

Answer 37

Sodium chlorate (I), which contains chlorate(I) ions that act as oxidizing agents to kill bacteria.

Question 38

How do you test for carbonate ions?

Answer 38

- Add dilute nitric acid (HNO₃) to the sample. - If carbonate ions are present, effervescence occurs due to CO₂ gas

Question 39

What is the equation for the carbonate test?

Answer 39

CO₃²⁻(aq) + 2H⁺(aq) → CO₂(g) + H₂O(l)

Question 40

How do you confirm the presence of carbon dioxide gas?

Answer 40

Bubble the gas through limewater (Ca(OH)₂ solution). If CO₂ is present, limewater turns cloudy due to the formation of calcium carbonate

Question 41

How do you test for sulfate (SO₄²⁻) ions?

Answer 41

- Add dilute nitric acid (HNO₃) to remove carbonate impurities - Add a few drops of barium chloride or barium nitrate - If sulfate ions are present, a white precipitate of barium sulfate forms.

Question 42

What is the equation for the sulfate test?

Answer 42

Ba²⁺(aq) + SO₄²⁻(aq) → BaSO₄(s)

Question 43

How do you test for halide ions (Cl⁻, Br⁻, I⁻)?

Answer 43

- Add dilute nitric acid (HNO₃) to remove interfering ions. - Add silver nitrate (AgNO₃) solution. - Observe the precipitate color to identify the halide.

Question 44

What are the precipitate colors for halide ions in the silver nitrate test?

Answer 44

- Chloride (Cl⁻): White precipitate (AgCl) - Bromide (Br⁻): Cream precipitate (AgBr) - Iodide (I⁻): Yellow precipitate (AgI)

Question 45

What is the equation for the halide test with silver nitrate?

Answer 45

Ag⁺(aq) + X⁻(aq) → AgX(s) (where X = Cl⁻, Br⁻, I⁻)

Question 46

How do you confirm halide identity using ammonia (NH₃) solution?

Answer 46

- AgCl: Dissolves in dilute NH₃. - AgBr: Insoluble in dilute NH₃ but dissolves in concentrated NH₃. - AgI: Insoluble in both dilute and concentrated NH₃.

Question 47

Why is it important to eliminate interfering ions? (CARBONATE)

Answer 47

- can give a false positive in the sulfate test because BaCO₃ is also a white precipitate. - Add acid first to remove carbonates as CO₂

Question 48

Why is it important to eliminate interfering ions (SULFATE)

Answer 48

- can interfere with the halide test because Ag₂SO₄ forms a white precipitate like AgCl. - Solution: Add Ba²⁺ ions first to remove sulfates as BaSO₄.

Question 49

What is the correct order for testing anions?

Answer 49

- Carbonates - Sulfates - Halides