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Halogens. Flashcards

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1
Q

Give the key physical properties of Fluorine. (F2)

A

Fluorine (F2): very pale yellow gas (almost see through). It is highly reactive.

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2
Q

Give the key physical properties of Chlorine. (Cl2)

A

Chlorine (Cl2): greenish reactive gas that is poisonous when in high concentrations. Also green in solution.

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3
Q

Give the key physical properties of Bromine. (Br2)

A

Bromine (Br2) : Brown liquid, that gives off dense brown/orange poisonous fumes from its liquid state.

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4
Q

Give the key physical properties of Iodine. (I2)

A

Iodine (I2) : shiny grey solid that sublimes to purple gas.

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5
Q

State and explain the trend in melting and boiling point for the halogen group.

A

Increase down the group.
As the molecules become larger they have more electrons and so have larger van der waals forces between the molecules. As the intermolecular forces get larger more energy has to be put into break the forces. This increases the melting and boiling points.

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6
Q

State and explain the trend in electronegativity in the halogen group. Define electronegativity in your answer.

A

Electronegativity is “the relative tendency of an atom in a molecule to attract electrons in a covalent bond towards itself”.
As one goes down the group the electronegativity of the elements decreases.
As one goes down the group the atomic radii increases due to the increasing number of shells. The nucleus is therefore less able to attract the bonding pair of electrons.

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7
Q

In what situation will a halogen displace a halide ion from its compound. And hence what is a halogens ability to displace a halide in its compound dependent on.

A

A halogen that is a strong oxidising agent will displace a halogen that has a lower oxidising power from one of its compounds. In other words the more reactive one will displace the less reactive one and hence if it is above it vertically in the periodic table it will displace it.

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8
Q

Define oxidising agent and state the trend in oxidising ability within the halogen group. In what situation is oxidising ability relevant.

A

Oxidising agents are electron acceptors.
The oxidising strength decreases down the group as the HALOGEN ATOMS get less sticky and less powerful attracters of electrons as their charge density decreases.
Oxidising ability is relevant to halogen atoms not halide ions since they will accept not donate electrons to complete a full outer shell. 7e- —> 8e-
This will form the negative ion that lets the halogen boot out the other halogen.
Hence oxidising ability should be considered in regards to halogen displacement reactions only.

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9
Q

Summarise which halogens will displace which.

A

Chlorine will displace both bromide and iodide ions; bromine will displace iodide ions and iodide ions won’t displace anything so will just remain in solution.

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10
Q

How can you work out which displacement if any has taken place.

A

If you learn the three colours of halogen in solution (can be different from their normal colour) you can observe the colour change and know that the new colour of solution will correspond to the. halogen free in solution this been displaced.
Equally if there is no observed colour change then the added halogen was unable to displace the halide in the compound.
Each halogen will have its own “free in solution colour”.

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11
Q

Give the colour of Chlorine in solution.

A

Chlorine =very pale green solution (often colourless)

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12
Q

Give the colour of Bromine in solution.

A

Bromine = yellow solution

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13
Q

Give the colour of iodine in solution.

A

Iodine = brown solution (sometimes black solid present)

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14
Q

Give the colours of the resulting following solutions.

1) Adding chlorine to sodium Iodide.
2) Adding chlorine to sodium Bromide.
3) Adding bromine to potassium Iodide.
4) Adding bromine to sodium Chloride.
5) Adding iodine to sodium Bromide.
6) Adding chlorine to sodium chloride.

A

1) Brown solution
2) Yellow solution
3) Brown solution
4) Yellow solution no recation
5) Brown solution no reaction
6) Very pale green solution - no reaction.

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15
Q

Write two half equations and hence the overall ionic equation for the reaction that occurs when chlorine is mixed with sodium bromide.

A 
Half equations. 
2Br - (aq)--->Br2 (aq)+ 2e-
Cl2 (aq)+2e---->2Cl- (aq)
Simplest ionic equation. 
Cl2(aq) + 2Br – (aq) ---> 2Cl – (aq) + Br2(aq)
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16
Q

Describe the reaction that is used to determine which Halide ION is present in solution. Give reactions to help your answer. (6 marks)

A

The test solution is made acidic with nitric acid, and then silver nitrate solution is added dropwise.
The role of nitric acid is to react with any carbonates present to prevent formation of the precipitate Ag2CO3. This would mask the desired observations
2HNO3 + Na2CO3 —> 2NaNO3 + H2O + CO2

Observations. 
Fluorides produce no precipitate as are soluble in water. 
Chlorides produce a white precipitate 
Ag+(aq) + Cl- (aq) --> AgCl(s)
Bromides produce a cream precipitate 
Ag+(aq) + Br- (aq)--->AgBr(s)
Iodides produce a pale yellow precipitate 
Ag+(aq) + I- (aq)---> AgI(s)

Treatment with ammonia.

The silver halide precipitates can be treated with ammonia solution to help differentiate between them if the colours look similar:

Silver chloride dissolves in dilute ammonia to form a complex ion
AgCl(s) + 2NH3(aq) —>[Ag(NH3)2]+ (aq) + Cl- (aq)
Colourless solution

Silver bromide dissolves in concentrated ammonia to form a complex ion
AgBr(s) + 2NH3(aq) —>[Ag(NH3)2]+ (aq) + Br - (aq)
Colourless solution
Silver iodide does not react with ammonia – it is too insoluble.

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17
Q

In regards to what is reducing power relevant and define and explain the trend in this property in group 7.

A

Reducing power is relevant to HALIDES (not halogens) and their reactions, particularly with sulphuric acid.
Definition: A reducing agent donates electrons.
The reducing power of the halides increases down group 7 They have a greater tendency to donate electrons.
This is because as the ions get bigger it is easier for the outer electrons to be given away as the pull from the nucleus on them becomes smaller.

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18
Q

How can the increasing power as reducing agents of halides be demonstrated. Give an example of the type of reagent that would be used.

A

This can be clearly demonstrated in the various reactions of the solid halides with concentrated sulfuric acid.
Could use the solid. sodium halides e.g NaF , NaCl

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19
Q

Explain the type of reaction that occurs when sodium fluoride or sodium chloride (F- or Cl-) reacts with concentrated sulphuric acid.

A

F- and Cl- ions are not strong enough reducing agents to reduce the S in H2SO4. No redox reactions occur. Only acid-base reactions occur.

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20
Q

State the role of sulphuric acid when it reacts with NaF or NaCl.

A

These are acid –base reactions and not redox reactions. H2SO4 plays the role of an acid (proton donor).

21
Q

Give the balanced symbol equations for the reactions of NaF and NaCl with concentrated sulphuric acid respectively and state also the observation and what compound will cause this observation in either case.

A

1) NaF(s) + H2SO4(l)—>NaHSO4(s) + HF(g) Observations: White steamy fumes of HF are evolved.NaCl(s) +

2) H2SO4(l)—> NaHSO4(s) + HCl(g)
Observations: White steamy fumes of HCl are evolved.

22
Q

What is the ONLY reaction product that occurs by a redox reaction when sodium bromide is reacted with concentrated sulphuric acid.

A

Sulfur dioxide ONLY.

23
Q

Give the acid base step that will occur first, before redox, when you react NaBr with concentrated Sulphuric acid. State the observation for this step.

A

Acid-base step:

NaBr(s) + H2SO4(l)–> NaHSO4(s) + HBr(g). Observations: White steamy fumes of HBr are evolved.

24
Q

Explain why there will be different reactions with sulphuric acid if it is reacted with sodium bromide or sodium chloride and state the nature of this other reaction in terms of what changes.

A

Br- ions are stronger reducing agents than Cl- and F- and after the initial acid-base
reaction reduce the sulfur in H2SO4 from +6 to + 4 in SO2 .

25
Q

By first forming any relevant half equations write an equation for the redox step (reaction) that takes place when NaBr is reacted with concentrated sulphuric acid. Give the property of the main reaction product.

A

Ox 1⁄2 equation: 2Br- —> Br2 + 2e-
Re1⁄2equation H2SO4 +2H+ +2e- —> SO2 +2H2O

Redox step ( reaction) :
2H+. + Br-  +H2SO4 ---> Br2(g) + SO2(g) + 2H2O(l)

colourless, acidic gas = SO2 that will etch glass.

26
Q

State the role of the concentrated H2SO4 when it is reacted NaBr or NaI in redox reactions.

A

Concentrated H2SO4 acts as an oxidising agent in the second redox step (reaction).

27
Q

Give the reaction products that are formed from a reaction of NaI and concentrated sulphuric acid form 1) Acid base step.
2) The redox step.
State any observations made for these products also.

A

Via acid base: HI - White steamy fumes of HI are evolved.

Via Redox:
Sulfur dioxide - colourless, acidic gas
Sulfur - A yellow solid
Hydrogen sulfide - a gas with a bad egg smell

Note: Black solid and purple fumes of Iodine are also evolved as a bi-product from the redox reactions.

28
Q

Define Disproportionation.

A

Disproportionation is the name for a reaction where an element simultaneously oxidises and reduces. Chlorine is both simultaneously reducing and oxidising in its standard reaction with water in absence of sunlight.

29
Q

Give the standard reaction of of chlorine with water and give the oxidation states. Also give two really specific features for this reaction you must include.

A

Features: Reversible reaction sign.
Include the oxidation state of Cl2 as being 0 as asked for oxidation states to show disproportionation - don’t just leave it out.
Cl2 + H20 Reversible HClO + HCl
0 +1 -1

30
Q

Study the following equilibrium established when chlorine reacts with water in the absence of sunlight.
Cl2 + H20 Reversible HClO + HCl
Give the name of the product formed and state what would be seen if some universal indicator where to be added to the solution.

A

Product formed is Chloric (I) acid - where the Roman numeral shows the oxidation state.

If some universal indicator is added to the solution it will first turn red due to the acidity of both reaction products. It will then turn colourless as the HClO bleaches the colour.

31
Q

State where this reaction would take place:
Cl2 + H20 Reversible HClO + HCl
State also the use of Chloric (I) acid.

A

This reversible reaction takes place when chlorine is used to purify water for drinking and in swimming baths , to prevent life threatening diseases.
Chloric(I) acid is an oxidising agent that kills bacteria by oxidation and it is also a bleach.

32
Q

Give the reaction that will occur if chlorine is bubbled through water in the presence of bright sunlight. What type of reaction is this.

A

2Cl2 + 2H2O NON Reversible 4HCl + O2

This reaction is non reversible.

33
Q

What will happen if an equilibrium mixture of chlorine water is left in sunlight. State also what happens to the colour.

A

2Cl2 + 2H2O NON Reversible Into 4HCl + O2
The same reaction occurs to an equilibrium mixture of chlorine water when standing in sunlight. The greenish colour of chlorine water fades as the Cl2 reacts and a colourless gas (O2) is produced.

34
Q

What will cause a pale green colour in a solution.

A

The greenish colour of these solutions is due to the Cl2.

35
Q

Give the reaction for a halogen (any of Cl2, and Br2, I2) reacting with cold sodium hydroxide state also what would happen to the colour of the particular halogen solution.
Give also the significance of the products.

A

Cl2 (aq) + 2 NaOH (aq) NON Reversible NaCl (aq) + NaClO (aq) + H2O (l)
The colour of the halogen solution will fade to colourless.
The mixture of NaCl and NaClO is used as bleach and to disinfect/ kill bacteria.

36
Q

What is used as bleach to disinfect and kill bacteria

A

The mixture of NaCl and NaClO is used as bleach and to disinfect/ kill bacteria.

37
Q

Why is chlorine used in water treatment to kill bacteria.

A

The benefits to health of water treatment by chlorine outweigh its toxic effects. It is also only used in small amounts.

38
Q

Give an alternative method to the direct chlorination of swimming pools , give the balanced symbol equation for this.

A

Addition of sodium or calcium chlorate(I) . Remember oxidation state.
NaClO(s) + H20 Reversible into NaOH-(aq) + HClO.
Remember the reversible reaction and naming chlorates and sulphates with roman numerals.

39
Q

State when using the alternative method to direct chlorination pools are kept slightly acidic and state how this is done.

A

NaClO(s) + H20 Reversible into NaOH-(aq) + HClO.
The above equilibrium is established and in alkaline solution the equilibrium will shift to the left removing the HClO as ClO- ions. To prevent this swimming pools are kept very slightly acidic . This is carefully monitored so that the water will not get too acidic to corrode metal components and affect the swimmers.

Get the state symbols right.

40
Q

Give a summary of the different chemicals produced by the recations of chlorine in regards to the effects on killing bacteria.

A

Chlorine is what is added to water and its products are what do the killing of bacteria and disinfecting.
Chloric acid HClO from standard reaction will kill bacteria via oxidation , it is also a bleach.
The mixture of NaCl and NaClO produced from chlorines reaction with cold sodium hydroxide is what is used as bleach and to disinfect/ kill bacteria.

41
Q 
What are sulphates and chlorates and how should they be named. 
Hence name: 
NaClO 
NaClO3 
K2SO4 
 K2SO3
A

In IUPAC convention the various forms of sulfur and chlorine compounds where oxygen is combined are all called sulfates and chlorates with relevant oxidation number given in roman numerals. If asked to name these compounds remember to add the oxidation number.

NaClO: sodium chlorate(I)
NaClO3: sodium chlorate(V)
K2SO4 potassium sulfate(VI)\
K2SO3 potassium sulfate(IV)

42
Q

Give the equation for the reaction of sodium iodine with sulphuric acid to produce Hydrogen iodide and give the relevant observations. Name the product.

A

NaI(s) + H2SO4(l) —> NaHSO4(s) + HI(g)
White steamy fumes of HI are evolved.
Sodium Hydrogen sulphate.

43
Q

Give the reaction of HI with sulphuric acid to produce SO2 and state the specific observations this would lead to. Give the oxidation states of the relevant compounds.

A

2H++ 2I- +H2SO4—->I2(s) + SO2(g) + 2H2O(l)A colourless, acidic gas SO2 that will etch glass.
Black solid and purple fumes of Iodine are also evolved
( as with all reactions , apart from the acid base one, that involve the iodide ion).

44
Q

Give the reaction of HI with sulphuric acid to produce H2S and state the specific observation this would lead to. Give the oxidation states of relevant compounds.

A

8 H+ +8I- + H2SO4—>4 I2(s) + H2S(g) + 4 H2O(l).
H2S (Hydrogen sulfide), a gas with a bad egg smell. Black solid and purple fumes of Iodine are also evolved
( as with all reactions , apart from the acid base one, that involve the iodide ion).

45
Q

When H2SO4 reacts with metal halides what are the two types of reaction and hence what are the two possible roles that the sulphuric acid can take. What is the deciding factor as to the type of reaction that occurs.

A

Acid Base: It plays the role of an acid so a PROTON DONOR.
For Redox the sulphuric acid plays the role of an oxidising agent since the halides are oxidised into their respective halogens and the the sulphuric acid is itself reduced.

Dependent on reducing power of the halogen.

46
Q

Give the four halides that can react with sulphuric acid and state the number of products in each case and name these products.

A

Acid base:
Flouride: 1 Acid base possible: HF
Chloride: 1 acid base possible: HCl
Redox:
Bromide: 1 acid base and 1 redox: HBr, SO2
Iodide: 1 acid Base and 3 possible redox: HI , SO2 , S , H2S.

47
Q

What do are the different possible oxidation states that the iodide Ion can reduce the sulphur in sulphuric acid to. State the oxidation state of Sulphur in Sulphuric acid.

A 
Original sulphuric acid is:  +6 
From Highest to lowest:
SO2-  =.    +4
S         =      0
H2S.   =      -2
48
Q

What is the other way that an ionic equation (Redox equation) can be written - subtle difference.

A

They can be written with combining the halogen and the Ion so instead of 2H+ and 2I- being in the reactants they could be combed as 2HI.

49
Q

Think about the weird overall type equation for the ones obove and why you write them.

A

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Chemistry A- Level (18 decks)

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