3.2.3 Group 7 halogens

Describe fluorine, chlorine, bromine and iodine in terms of colour, state and other properties

• fluorine- very pale yellow glass and highly reactive

• chlorine- green gas, reactive and poisonous in high concs

• bromine- red-brown liquid, gives off dense brown/orange poisonous fumes

• iodine- shiny grey solid sublimes to purple gas

What happens to MP and BP down group 7?

• increases down the group

• as molecules become larger, they have more electrons and have larger VDW forces between molecules

• intermolecular forces get larger meaning more energy needed to break forces

What happens to electronegativity down the group?

• electronegativity decreases

• nuclear charge increases but despite this:

• atomic radii increases as number of shells increases

• shielding increases- inner electron shells reduce atom’s ability to attract electrons

What happens to reactivity down group 7?

• reactivity decreases

• atomic radius increases

• more electron shells added meaning outer electron experiences more shielding and is further away from nucleus

• therefore electrostatic attraction between outer electron and nucleus gets weaker

• becomes harder for larger halogens to attract electron

What 4 types of reactions do halogens take part in?

1. displacement reactions (displacing less reactive halide ions from solution)

2. reaction with silver nitrate

3. reaction of halide salts with conc sulfuric acid

4. disproportionation reactions of chlorine

In a displacement reaction, what is the role of the halogen?

acts as a strong oxidising agent that displaces a halogen with a lower oxidising power from one of its compounds

Why halogens act as oxidising agents?

• halogens gain an electron meaning it’s reduced as it’s oxidation number decreases from 0 to -1

• causes another substance to be oxidised

How does oxidising strength change down the group and why?

• oxidising strength decreases down the group

  • oxidising agents are electron acceptors

• cos reactivity decreases as atoms get larger and outer shell is further from nucleus so electrostatic attraction between outer electron and nucleus gets weaker

When will a halogen displace another halide from solution?

If the halide is below it in the periodic table (less reactive)

What will chlorine displace? Write their ionic equations

• bromine (Br^-)

• Iodide (I^-)

What will bromine displace? Write their ionic equation

• Iodide (I^-)

What will iodine displace?

It doesn’t displace fluorine chlorine or bromine

If chlorine was the free halogen present in a solution in a test tube, what colour would show?

very pale green solution (often colourless)

If bromines was the free halogen present in a solution in a test tube, what colour would show?

yellow solution

If iodine was the free halogen present in a solution in a test tube, what colour would show?

brown solution

Fill in this table. In each box, write what colour would show and if a displacement has occurred.

What test is used to identify which halide ion is present?

silver nitrate test- nitric acid is added to the test solution and then silver nitrate is added dropwise

What is observed when F, Cl, Br & I react with the silver nitrate? Write the ionic equations for each

What is the role of nitric acid in the reaction of halide ions with silver nitrate?

to react with any carbonates present to prevent formation of the precipitate Ag₂CO₃ as this would mask the desired observations

What can you do if the colour of the silver halide precipitates look similar?

• treat them with ammonia solution to help differentiate the colours

What is the reducing power of halide ions and what is the trend down group 7?

• reducing power refers to the ability of halide ions to donate electrons to other substances

• reducing ability of halide ions increase down group 7

• increasing ionic radius so electrons are further from nucleus

• outer electrons experience more shielding so electrostatic attraction between outer electron and nucleus gets weaker

• becomes easier for larger halide ions to lose electrons and become oxidised

  • Fluoride is the weakest reducing agent and iodide is the strongest

  • this explains the reactions of halide

What happens when halide ions react with conc H₂SO₄ ?

hydrogen halide gases are initially produced but subsequent reactions depend on the reducing powers of the hydrogen halide formed

What happens when sodium fluoride and sodium chloride react with conc H₂SO₄ ? Write the equations

• white steamy fumes of HF/HCl gas are evolved

• no redox reactions occur only acid-base reactions

• fluoride and chloride have low reducing power and are not strong enough reducing agents to reduce the S in H₂SO₄

• oxidation states of halide and sulfur stay the same

What happens when NaBr reacts with conc H₂SO₄ ? Write the equations and any half equations involved and also state the role of H₂SO₄

• first acid-base reaction gives white misty fumes of HBr gas

• since HBr is a stronger reducing agent it reduces the S in H₂SO₄ from +6 to +4 - reaction produces toxic fumes of SO₂ and orange fumes of Br₂ (redox reaction)

• overall equation: 2NaBr + 2H₂SO₄ → Na₂SO₄ + SO₂ + Br₂ + 2H₂O

• H₂SO₄ acts as acid in first step and then acts as oxidising agent in 2nd step

What happens when NaI reacts with conc H₂SO₄ ? Write the equations and any half equations involved

• first acid-base equation gives white steamy fumes of HI gas

• I^- ions are the strongest halide reducing agents so they can reduce H₂SO₄ to SO₂ , H₂S or S

• the oxidation state of S goes from +6 to +4 to 0 to -2

• black solid and purple fumes of iodine are evolved

• yellow solid of sulfur formed

• H₂S is a toxic gas with a smell of bad eggs

What is a disproportionation reaction?

they occur when a substance is simultaneously oxidised and reduced in the same chemical reaction

Chlorine reacts with water in a disproportionation reaction. Write the equation and explain why this is a disproportionation reaction

Cl_2(g) + H₂O_(l) ⇌ HCl_(aq) + HClO_(aq)

• oxidation of Cl increases from 0 to +1 in HClO and decreases to -1 in HCl

• you get a mixture of chloride and chlorate (I) ions

What happens if you add some universal indicator to the solution in Cl_2(g) + H₂O_(l) ⇌ HCl_(aq) + HClO_(aq)

it will first turn red due to the acidity of both products and then it will turn colourless as the HClO bleaches the colour

What happens when you react chlorine with water in sunlight? Write the equation

2Cl_2(g) + 2H₂O_(l) ⇌ 4H⁺_(aq) + 4Cl^-_{(aq) +} O_2(g)

• chlorine decomposes water to form chloride ions and oxygen

• the greenish colour of chlorine water fades as the Cl reacts and a colourless gas O₂ is produced

What is chlorine used for?

• water treatment to kill bacteria in drinking water and in swimming pools

• chlorate ions kill bacteria (ClO^-)

What are the downsides of using chlorine for water treatment?

• chlorine is toxic

• water contains variety of organic compounds and Cl reacts with these compounds to form chlorinated hydrocarbons eg. chloromethane - many are carcinogenic but these risks are small compared to risks from untreated water

How is bleach made from chlorine? Write the equation and state the type of reaction this is

• mixing chlorine gas with cold, dilute, aqueous NaOH produces sodium chlorate solution

• Cl_2(g) + 2NaOH_(aq) → NaClO_(aq) + NaCl_(aq) + H₂O_(l)

• oxidation number of Cl increases from 0 to +1 in NaClO and decreases to -1 in NaCl

• bleach solution contains chlorate ions that act as oxidising agents to kill bacteria

Name these chlorates/sulfates make sure to add oxidation number

NaClO

NaClO₃

K₂SO₄

K₂SO₃