Brainscape's Knowledge GenomeTM

Browse over 1 million classes created by top students, professors, publishers, and experts.


See full index

A LEVEL - Chemistry (Inorganic) > Topic 4 - Inorganic Chemistry and the Periodic Table > Flashcards

Topic 4 - Inorganic Chemistry and the Periodic Table Flashcards

You may prefer our related Brainscape-certified flashcards:
Periodic Table flashcards Chemistry 101 flashcards AP® Chemistry flashcards Organic Chemistry 101 flashcards
1
Q

Name the first 5 elements in group 2.

A
  • Beryllium
  • Magnesium
  • Calcium
  • Strontium
  • Barium
How well did you know this?
1
Not at all
2
3
4
5
Perfectly
2
Q

How does ionisation energy change going down group 2 and why?

A

• Decreases
BECAUSE:
• Extra shielding from extra inner shells
• Outer electron is further from the nucleus
(NOTE: The increased nuclear charge is overriden by the extra shielding)

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
3
Q

In a question asking about how ionisation energy changes going down group 2, what things must you mention?

A

1) Extra shielding
2) Increased distance from nucleus
3) Increased nuclear charge -> BUT cancelled out by shielding

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
4
Q

How does reactivity change going down group 2 and why?

A

• Increases
BECAUSE:
• It is easier to lose the outer two electrons
• Due to the extra shielding and distance from the nucleus -> Extra nuclear charge is overriden

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
5
Q

What 3 important things can group 2 elements react with?

A
  • Water
  • Oxygen
  • Chlorine
How well did you know this?
1
Not at all
2
3
4
5
Perfectly
6
Q

What is formed when group 2 metals react with water?

A

• Metal hydroxide
• Hydrogen
(Except beryllium)

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
7
Q

Give the general equation for group 2 metals reacting with water (where metal = M).

A

M + H2O -> M(OH)2 + H2

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
8
Q

What is an exception to the rule for group 2 metals reating with water?

A
  • Magnesium
  • It reacts like the others with water, except very slowly. This means that it can be made to react with steam instead, which produces MgO and H20.
How well did you know this?
1
Not at all
2
3
4
5
Perfectly
9
Q

Give the equation for:

1) Mg (s) + H20 (l) ->
2) Mg (s) + H20 (g) ->

A

1) Mg (s) + 2H20 (l) -> Mg(OH)2 (aq) + H2 (g)

2) Mg (s) + H20 (g) -> MgO (s) + H2(g)

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
10
Q

Describe how magnesium reacts with:

1) Water
2) Steam

A 
Water:
• Slowly
• Produces magnesium hydroxide
Steam:
• Quickly
• Produces magnesium oxide
How well did you know this?
1
Not at all
2
3
4
5
Perfectly
11
Q

Explain why magnesium reacts differently with water and steam.

A

At the high temperatures required to produce steam, any hydroxide would decompose into its oxide, so this is produced instead of the hydroxide.

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
12
Q

Describe how each of the group 2 metals reacts with water.

Be, Mg, Ca, Sr, Ba

A
  • Be - Doesn’t react
  • Mg - Very slowly
  • Ca - Steadily
  • Sr - Fairly quickly
  • Ba - Rapidly
How well did you know this?
1
Not at all
2
3
4
5
Perfectly
13
Q

What is formed when group 2 metals react with oxygen?

A

Metal oxides

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
14
Q

Give the general equation for group 2 metals reacting with oxygen (where metal = M).

A

2M + O₂ -> 2MO

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
15
Q

What is formed when group 2 metals react with chlorine?

A

Metal chlorides

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
16
Q

Describe the appearance of metal chlorides.

A

White solids

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
17
Q

Give the general equation for group 2 metals reacting with chlorine (where metal = M).

A

M + Cl2 -> MCl2

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
18
Q

Give the equations for calcium with:
• Water
• Oxygen
• Chlorine

A 
WATER:
Ca (s) + 2H20 (l) -> Ca(OH)2 (aq) + H2 (g)
OXYGEN:
2Ca (s) + O2 (g) -> 2CaO (s)
CHLORINE:
Ca (s) + Cl2 (g) -> CaCl2 (s)
How well did you know this?
1
Not at all
2
3
4
5
Perfectly
19
Q

Name the products of group 2 metals reacting with:
• Water
• Oxygen
• Chlorine

A 
WATER:
• Metal hydroxide + hydrogen
OXYGEN:
• Metal oxide
CHLORINE:
• Metal chloride
How well did you know this?
1
Not at all
2
3
4
5
Perfectly
20
Q

What acidity are group 2 metal hydroxides?

A

Alkaline.

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
21
Q

What is formed when group 2 oxides react with water?

A

Metal hydroxides (except beryllium).

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
22
Q

How does the solubility of group 2 metal oxides compare to that of metal hydroxides?

A

It is similar, since the oxide react to form metal hydroxides.

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
23
Q

Which of the group 2 metals is an exception?

A

Beryllium -> Be and BeO don’t react with water

NOTE: Magnesium is similar, but it reacts very slowly

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
24
Q

Describe the strength of the alkaline solution produced as you go down group 2.

A

The solutions become increasingly alkaline, since the metal hydroxide is more soluble.

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
25
Q

What makes metal hydroxide solutions alkaline?

A

The OH- ions

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
26
Q

What happens when group 2 metal oxides react with water?

A

• Metal hydroxide is formed
• This metal hydroxide dissolves
• Alkaline solution is formed
(Except with beryllium. Magnesium is very slow and hardly soluble.)

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
27
Q

What acidity are group 2 metal oxides?

A

Alkaline.

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
28
Q

Give the general equation for group 2 metal oxides reacting with water (where M = metal).

A

MO (s) + H2O (l) -> M(OH)2 (aq)

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
29
Q

Give the general equation for SOLID group 2 metal hydroxides reacting with water (where M = metal).

A

M(OH)2 (s) + H2O (l) -> M(OH)2 (aq)

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
30
Q

What happens when solid group 2 metal hydroxides are added to water?

A

They dissolve (if they are soluble) to give an alkaline solution.

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
31
Q

What is formed when group 2 metal oxides react with dilute acid?

A
  • Salt

* Water

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
32
Q

Give the general equation for metal oxides reacting with dilute hydrochloric acid (where M = metal).

A

MO (s) + 2HCl (aq) -> MCl2 (aq) + H2O (l)

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
33
Q

What is formed when group 2 metal hydreoxides react with dilute acid?

A
  • Salt

* Water

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
34
Q

Give the general equation for metal hydroxides reacting with dilute hydrochloric acid (where M = metal).

A

M(OH)2 (s) + 2HCl (aq) -> MCl2 (aq) + 2H2O (l)

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
35
Q

Are there solubility trends for group 2 metal compounds?

A

Yes, but they depends on the compound anion (e.g. OH-).

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
36
Q

Describe how the solubility of group 2 metal hydroxides changes going down the group.

A

It increases.

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
37
Q

Describe how the solubility of group 2 metal sulfates changes going down the group.

A

It decreases.

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
38
Q

Which group 2 compound can be used as a reference for all other group 2 compound solubilities?

A
  • Barium sulfate is insoluble.
  • It is at the bottom of the group, so metal sulfate solubility must decrease going down the group.
  • Metal hydroxides are the opposite of this, so their solubility must increase going down the group.
How well did you know this?
1
Not at all
2
3
4
5
Perfectly
39
Q

What’s more soluble, calcium hydroxide or strontium hydroxide?

A

Strontium hydroxide

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
40
Q

What’s more soluble, calcium sulfate or strontium sulfate?

A

Calcium sulfate

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
41
Q

What are compounds with very low solubilities said to be?

A

Sparingly soluble

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
42
Q

What is thermal decomposition?

A

When a substance breaks down when heated.

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
43
Q

What is thermal stability?

A

The amount of heat required to thermally decompose a substance.

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
44
Q

What is a substance with high thermal stability?

A

A substance that requires a lot of heat to be thermally decomposed.

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
45
Q

How does thermal stability of group 1 and 2 compounds change going down the group and why?

A

• It increases
BECAUSE:
• The metal ion polarises the anion, distorting it -> This makes the compound unstable
• Going down the group, the metal ion’s charge density decreases -> Polarising power decreases -> Less distortion of anion -> More thermally stable

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
46
Q

How does a cation’s charge density affect the thermal stability of its compound?

A

High charge density -> High polarising power -> Strong distortion of anion -> Thermally unstable compound

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
47
Q

Compare the thermal stability of group 1 and 2 compounds.

A

• Group 1 compounds are more thermally stable.
BECAUSE:
• The metal ion polarises the anion, distorting it -> This makes the compound unstable
• Group 1 compounds have a lower charge density (1+ instead of 2+) -> Polarising power decreases -> Less distortion of anion -> More thermally stable

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
48
Q

What are the general trends for the thermal stability of group 1 and 2 compounds?

A
  • Thermal stability increases down a group

* Group 1 compounds are more thermally stable than group 2 compounds

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
49
Q

Describe what happens when you heat group 1 carbonates.

A
  • They are thermally stable so they don’t decompose when heated by a bunsen burner (except Li2CO3, which decomposes to Li2O and CO2)
  • They do decompose at much higher temperatures
How well did you know this?
1
Not at all
2
3
4
5
Perfectly
50
Q

Describe what happens when you heat group 1 nitrates.

A
  • Decompose to the metal nitrite and oxygen

* Except LiNO3, which decomposes to form Li2O, NO2 and O2.

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
51
Q

Describe what happens when you heat group 2 carbonates.

A

• Decompose to the metal oxide and carbon dioxide

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
52
Q

Describe what happens when you heat group 2 nitrates.

A

• Decompose to the metal oxide, nitrogen dioxide and oxygen.

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
53
Q

Give the general equation for the thermal decomposition of group 1 carbonates (where M = metal).

A

They do not thermally decompose easily, except Li2CO3.

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
54
Q

Give the general equation for the thermal decomposition of group 1 nitrates (where M = metal).

A

2MNO3 (s) -> 2MNO2 (s) + O2 (g)

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
55
Q

Give the general equation for the thermal decomposition of group 2 carbonates (where M = metal).

A

MCO3 (s) -> MO (s) + CO2 (g)

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
56
Q

Give the general equation for the thermal decomposition of group 2 nitrates (where M = metal).

A

2M(NO3)2 (s) -> 2MO (s) + 4NO2 (g) + O2 (g)

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
57
Q

Which group 1 or 2 metal is an exception to the thermal decomposition rules and how does it behave?

A
  • Lithium

* Its carbonate and nitrate decompose like group 2 carbonates and nitrates

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
58
Q

What are the products of the thermal decomposition of lithium carbonate?

A

Lithium oxide and carbon dioxide

Remember: This is an exception!

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
59
Q

What are the products of the thermal decomposition of lithium nitrate?

A

Lithium oxide, nitrogen dioxide and oxygen

Remember: This is an exception!

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
60
Q

How can you test how easily a nitrate decomposes?

A

Time how long it takes to produce a certain amount of either:
• Oxygen
• Brown gas (NO2) -> In a fume cupboard

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
61
Q

How can you test how easily a carbonate decomposes?

A

Time how long it takes to produce a certain amount of:

• Carbon dioxide

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
62
Q

Name the flame colour of lithium.

A

Red

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
63
Q

Name the flame colour of sodium.

A

Orange/Yellow

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
64
Q

Name the flame colour of potassium.

A

Lilac

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
65
Q

Name the flame colour of rubidium.

A

Red

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
66
Q

Name the flame colour of caesium.

A

Blue

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
67
Q

Name the flame colour of calcium.

A

Brick red

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
68
Q

Name the flame colour of strontium.

A

Crimson

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
69
Q

Name the flame colour of barium.

A

Green

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
70
Q 
Name the flame colour of these group 1 metals:
• Li
• Na
• K
• Rb
• Cs
A
  • Li - Red
  • Na - Orange/Yellow
  • K - Lilac
  • Rb - Red
  • Cs - Blue
How well did you know this?
1
Not at all
2
3
4
5
Perfectly
71
Q

Name the flame colour of these group 2 metals:
• Ca
• Sr
• Ba

A
  • Ca - Brick red
  • Sr - Crimson
  • Ba - Green
How well did you know this?
1
Not at all
2
3
4
5
Perfectly
72
Q

Describe how to carry out a flame test.

A

1) Mix a small amount of the compound with a few drops of HCl
2) Heat a piece of platinum or nichrome wire in a hot Bunsen burner flame to clean it
3) Dip the wire into the compound-acid mixture and hold it in the flame.

73
Q

Explain why different elements have distinct flame colours.

A
  • Energy absorbed from the flame causes electrons to be excited and move to higher energy levels
  • As they de-excite to lower energy levels, energy is released as light
  • The difference between the higher and lower levels determines the wavelength of the light released
  • So the colour of each element’s flame is distinct
74
Q

What is the movement of electrons between energy levels called?

A

Electron transition

75
Q

What are the first 4 halogens?

A
  • Fluorine
  • Chlorine
  • Bromine
  • Iodine
76
Q

How does electronegativity change as you go down the halogens?

A

It decreases.

77
Q

Describe the colour and state of fluorine at RTP.

A
  • Pale yellow

* Gas

78
Q

Describe the colour and state of chlorine at RTP.

A
  • Green

* Gas

79
Q

Describe the colour and state of bromine at RTP.

A
  • Red-brown

* Liquid

80
Q

Describe the colour and state of iodine at RTP.

A
  • Grey

* Solid

81
Q

Describe and explain the solubility of halogens in water and organic solvents.

A
  • In water -> Low
  • In organic solvents -> High

Because they’re non-polar.

82
Q

Describe the colour of these halogens in water:
• Chlorine
• Bromine
• Iodine

A
  • Chlorine -> Colourless
  • Bromine -> Yellow/Orange
  • Iodine -> Brown
83
Q

Describe the colour of these halogens in hexane:
• Chlorine
• Bromine
• Iodine

A
  • Chlorine -> Colourless
  • Bromine -> Orange/Red
  • Iodine -> Pink/Violet
84
Q

What is the colour of fluorine:

• At RTP

A

• At RTP -> Pale yellow gas

85
Q

What is the colour of chlorine:
• At RTP
• In water
• In hexane

A
  • At RTP -> Green gas
  • In water -> Colourless
  • In hexane -> Colourless
86
Q

What is the colour of bromine:
• At RTP
• In water
• In hexane

A
  • At RTP -> Red-brown liquid
  • In water -> Yellow/Orange
  • In hexane -> Orange/Red
87
Q

What is the colour of iodine:
• At RTP
• In water
• In hexane

A
  • At RTP -> Grey solid
  • In water -> Brown
  • In hexane -> Pink/Violet
88
Q

Explain how the reactivity of halogens changes as you go down the group.

A

It decreases, because:
• Halogens react by gaining at electron
• As you go down the group, the atoms are larger so the outer electrons are further away
• There is also more shielding due to the extra electron shells

89
Q

Explain how the melting and boiling points of halogens change as you go down the group.

A

It increases, because:
• There is an increase in the number of electron shells, which means the London forces are stronger
• So more energy is required to overcome them, increasing the melting and boiling points

90
Q

Why is the chemistry of fluorine and astatine difficult to study?

A
  • Fluorine -> Toxic has

* Astatine -> Radioactive and so it decays quickly

91
Q

What happens when you add a solution of a halogen to a solution of a hydrogen halide?

A
  • If the halogen is more reactive -> It will displace the halide from its solution
  • If the halogen is less reactive -> No reaction
92
Q

Which halides will chlorine displace from their solution?

A

Bromide and iodide

93
Q

Which halides will bromine displace from their solution?

A

Iodide

94
Q

Which halides will iodine displace from their solution?

A

None

95
Q

What type of reaction is the displacement reaction between a halogen and a halide?

A

Redox

96
Q

In a halogen-halide reaction, is the halogen reduced or oxidised?

A

Reduced

97
Q

In a halogen-halide reaction, is the halide reduced or oxidised?

A

Oxidised

98
Q

Remember to practise writing out ionic equations for halogen-halide reactions.

A

Pg 47 of revision guide

99
Q

In a halogen-halide reaction, what can be done to make the colour change more obvious?

A

Shake the reaction mixture with an organic solvent. The halogen present will dissolve in it and will settle out as a distinct, clearly-coloured layer above the aqueous solution.

100
Q

How do halogens react with group 1 and 2 metals?

A

They produce halide salts.

101
Q

2Li + F₂ ->

A

2Li + F₂ -> 2LiF

102
Q

Mg + Cl₂ ->

A

Mg + Cl₂ -> MgCl₂

103
Q

What sort of reaction do halogens undergo with alkalis?

A

Disproportionation

104
Q

What is it important to remember about halogen reactions with alkalis?

A

The products depend on whether the alkali is hot or cold.

105
Q

Give the general equation for a halogen reacting with a cold alkali.
(Use NaOH as the alkali)

A

X₂ + 2NaOH -> NaOX + NaX + H₂O

106
Q

When halogen reacts with a cold alkali, what happens to the oxidation number of the halogen?

A

0 (X₂) -> +1 (NaOX) and -1 (NaX)

107
Q

I₂ + 2NaOH ->

When the alkali is cold

A

I₂ + 2NaOH -> NaOI + NaI + H₂O

108
Q

Name: NaOI

A

Sodium iodate(I)

109
Q

Describe the oxidation states that halogens can exist in.

A

They can exist in many. For example: Br, BrO⁻, BrO₂⁻

110
Q

What is the chemical name and formula for bleach?

A
  • Sodium chlorate(I) solution

* NaClO

111
Q

Describe how you can prepare bleach.

A

Mix chlorine gas with cold, dilute aqueous sodium hydroxide.

112
Q

What type of reaction is the preparation of bleach?

A

Disproportionation

113
Q

Give the equation for the formation of bleach.

A

2NaOH + Cl₂ -> NaClO + NaCl + H₂O

Note: The NaOH must be cold

114
Q

What are some uses of sodium chlorate(I)?

A

It is bleach, so:
• Water treatment
• Bleaching paper and textiles
• Cleaning toilets

115
Q

Give the general equation for a halogen reacting with a hot alkali.
(Use NaOH as the alkali)

A

3X₂ + 6NaOH -> NaXO₃ + 5NaX + 3H₂O

116
Q

When halogen reacts with a hot alkali, what happens to the oxidation number of the halogen?

A

0 (X₂) -> +5 (NaXO₃) and -1 (NaX)

117
Q

Cl₂ + NaOH ->

When the alkali is hot

A

3Cl₂ + 6NaOH -> NaClO₃ + 5NaCl + 3H₂O

118
Q

Name: NaClO₃

A

Sodium chlorate(V)

119
Q

What is chlorine used for?

A

Killing bacteria in water.

120
Q

What type of reaction is mixing chlorine with water!

A

Disproportionation

121
Q

What is produced when you mix chlorine with water?

A
  • Hydrochloric acid

* Hypochlorous acid

122
Q

Give the chemical formula for hypochlorous acid.

A

HClO

123
Q

What is the equation for the reaction between chlorine and water?

A

Cl₂ + H₂O {-} HCl + HClO

124
Q

When chlorine reacts with water, what happens to the oxidation number of the halogen?

A

0 (Cl₂) -> -1 (HCl) and +1 (HClO)

125
Q

Do other halogens except chlorine also undergo disproportionation when mixed with water?

A

Yes

126
Q

What happens to hypochlorous acid in water?

A

It ionises to make chlorate(I) ions and H₃O⁺.

127
Q

Give the equation for what happen to hypochlorous acid in water.

A

HClO + H₂O {-} ClO⁻ + H₃O⁺

128
Q

Give the chemical formula for chlorate(I) ions.

A

ClO⁻

129
Q

What is another name for chlorate(I) ions? (ClO⁻)

A

Hypochlorite ions

130
Q

Which part of chlorine causes it to kill bacteria?

A

The chlorate(I) (ClO⁻) ions it Eventually forms.

131
Q

Give the entire pathway for how chlorine kills bacteria in water.

A

• Chlorine reacts with the water:
Cl₂ + H₂O {-} HCl + HClO
• The hypochlorous acid dissociates in the water:
HClO + H₂O {-} ClO⁻ + H₃O⁺
• The chlorate(I) (ClO⁻) ions are what kills the bacteria.

132
Q

How can halides act as reducing agents?

A

They can lose an electron from the outer shell.

133
Q

Explain how the reducing power of halides changes as you go down the group.

A

It increases because:
• The ion gets bigger, so the electrons are further away from the positive nucleus
• There is greater shielding from the extra inner shells
• This means the outer electrons are easier to lose

134
Q

Do all group 1 halides react the same with sulfuric acid?

A
  • No, it depends on their reducing power.

* The further down the group you go, the more times the halide ion will react with the sulfuric acid.

135
Q 
Describe the number of times these will react with sulfuric acid:
• KF
• KCl
• KBr
• KI
A
  • KF + KCl -> Once
  • KBr -> Twice
  • KI -> Three times
136
Q

Give the equations for the reactions of KF or KCl with H₂SO₄. Include state symbols.

A

KF(s) + H₂SO₄(l) -> KHSO₄(s) + HF(g)

KCl(s) + H₂SO₄(l) -> KHSO₄(s) + HCl(g)

137
Q

KF(s) + H₂SO₄(l) -> KHSO₄(s) + HF(g)

KCl(s) + H₂SO₄(l) -> KHSO₄(s) + HCl(g)

What is observed in these reactions?

A

Misty fumes (as the gas comes into contact with moisture)

138
Q

KF(s) + H₂SO₄(l) -> KHSO₄(s) + HF(g)

KCl(s) + H₂SO₄(l) -> KHSO₄(s) + HCl(g)

Are these redox reactions?

A

No, because the oxidation numbers stay the same.

139
Q

KF(s) + H₂SO₄(l) -> KHSO₄(s) + HF(g)

KCl(s) + H₂SO₄(l) -> KHSO₄(s) + HCl(g)

Why are there no follow-up reactions to these?

A

The F⁻ and Cl⁻ ions are not strong enough reducing agents to reduce the sulfuric acid.

140
Q

Give the equations for the reactions of KBr with H₂SO₄. Include state symbols.

A
  • KBr(s) + H₂SO₄(l) -> KHSO₄(s) + HBr(g)

* HBr(g) + H₂SO₄(l) -> Br₂(g) + SO₂(g) + 2H₂O(l)

141
Q
  • KBr(s) + H₂SO₄(l) -> KHSO₄(s) + HBr(g)
  • HBr(g) + H₂SO₄(l) -> Br₂(g) + SO₂(g) + 2H₂O(l)

What is observed in these reactions?

A 
First reaction:
• Misty fumes of HBr gas
Second reaction:
• Choking fumes of SO₂
• Orange fumes of Br₂
142
Q
  • KBr(s) + H₂SO₄(l) -> KHSO₄(s) + HBr(g)
  • HBr(g) + H₂SO₄(l) -> Br₂(g) + SO₂(g) + 2H₂O(l)

Are these redox reactions?

A

First reaction:
• Not redox
Second reaction:
• Redox

143
Q
  • KBr(s) + H₂SO₄(l) -> KHSO₄(s) + HBr(g)
  • HBr(g) + H₂SO₄(l) -> Br₂(g) + SO₂(g) + 2H₂O(l)

Why are there no follow-up reactions to these?

A

The Br⁻ ions are only strong enough reducing agents to reduce the sulfuric acid once.

144
Q

Give the equations for the reactions of KI with H₂SO₄. Include state symbols.

A
  • KI(s) + H₂SO₄(l) -> KHSO₄(s) + HI(g)
  • HI(g) + H₂SO₄(l) -> I₂(s) + SO₂(g) + 2H₂O(l)
  • 6HI(g) + SO₂(g) -> H₂S(g) + 3I₂(s) + 2H₂O(l)
145
Q
  • KI(s) + H₂SO₄(l) -> KHSO₄(s) + HI(g)
  • HI(g) + H₂SO₄(l) -> I₂(s) + SO₂(g) + 2H₂O(l)
  • 6HI(g) + SO₂(g) -> H₂S(g) + 3I₂(s) + 2H₂O(l)

What is observed in these reactions?

A 
First reaction:
• Misty fumes of HI gas
Second reaction:
• Choking fumes of SO₂ (check this)
• Grey solid of I₂
Third reaction:
• Smell of rotten eggs of H₂S
• Grey solid of I₂
146
Q
  • KI(s) + H₂SO₄(l) -> KHSO₄(s) + HI(g)
  • HI(g) + H₂SO₄(l) -> I₂(s) + SO₂(g) + 2H₂O(l)
  • 6HI(g) + SO₂(g) -> H₂S(g) + 3I₂(s) + 2H₂O(l)

Are these redox reactions?

A

First reaction:
• Not redox
Second and third reaction:
• Redox

147
Q

Why does KI react with H₂SO₄ three times?

A

The iodide ions are such powerful reducing agents.

148
Q

What is the indicator of each group 1 halide reacting with H₂SO₄?

A
  • KF and KCl -> Misty fumes (although these are seen with all halides)
  • KBr -> Choking fumes of SO₂ and orange fumes of Br₂
  • KI -> Rotten egg smell of H₂S
149
Q

What is an indicator of H₂S?

A

The smell of rotten eggs.

It is also toxic.

150
Q

What types of chemical are hydrogen halides?

A

Acidic gases

151
Q

What happens when hydrogen halides come in contact with water?

A
  • They dissolve
  • Misty fumes of acidic gas are produced
  • This would turn damp blue litmus paper red
152
Q

What happens when hydrogen halides come in contact with ammonia gas?

A

They react to give white fumes.

153
Q

What is the equation for ammonia gas reacting with HCl gas?

What changes are seen?

A

NH₃(g) + HCl(g) -> NH₄Cl(s)

This gives white fumes.

154
Q

Describe the rest for halide ions.

A
  • Add dilute nitric acid
  • Add silver nitrate solution
  • Check the colour of the precipitate
155
Q

When testing for halides, why do you add dilute nitric acid?

A

To remove ions which might interfere with the reaction.

156
Q

What type of precipitate is formed in the rest for halide ions?

A

Silver halide

157
Q

When testing for halides, what is the positive result for fluoride ions?

A

No precipitate

158
Q

When testing for halides, what is the positive result for chloride ions?

A

White precipitate

159
Q

When testing for halides, what is the positive result for bromide ions?

A

Cream precipitate

160
Q

When testing for halides, what is the positive result for iodide ions?

A

Yellow precipitate

161
Q

When testing for halides, what is the positive result for each of the halides?

A
  • Fluorine -> No precipitate
  • Chlorine -> White precipitate
  • Bromine -> Cream precipitate
  • Iodide -> Yellow precipitate
162
Q

When testing for halides, how can you tell apart the different precipitates if the colour is not obvious?

A

Add ammonia and see if the precipitate dissolves.

163
Q

After testing for halides, what happens to the silver chloride precipitate when ammonia is added?

A

It dissolves in DILUTE ammonia solution to give a colourless solution.

164
Q

After testing for halides, what happens to the silver bromide precipitate when ammonia is added?

A
  • Remains unchanged when DILUTE ammonia is added

* Dissolves when CONCENTRATED ammonia is added to give a colourless solution

165
Q

After testing for halides, what happens to the silver bromide precipitate when ammonia is added?

A

It does not dissolve, even in concentrated ammonia solution.

166
Q

Give the ionic equation for the test for halides.

A

Ag⁺(aq) + X⁻(aq) -> AgX(s)

168
Q

What is the formula for hydrogencarbonate?

A

HCO₃⁻

169
Q

Describe how you can test for carbonates and hydrogencarbonates.

A
  • Add HCl
  • Bubble the gas through limewater
  • Positive result -> Limewater turns cloudy
170
Q

Give the ionic equation for the test for carbonate ions.

A

CO₃²⁻ + 2H⁺ -> CO₂ + H₂O

171
Q

Give the ionic equation for the test for hydrogencarbonate ions.

A

HCO₃²⁻ + H⁺ -> CO₂ + H₂O

172
Q

What is the formula for a sulfate ion?

A

SO₄²⁻

173
Q

Describe how you can test for sulfates.

A
  • Add dilute HCl
  • Add barium chloride solution (BaCl₂)
  • Positive result -> White precipitate of barium sulfate
174
Q

Give the ionic equation for the test for sulfates.

A

Ba²⁺(aq) + SO₄²⁻(aq) -> BaSO₄(s)

175
Q

What is the pH of ammonia gas?

A

Alkaline

Note: It accepts H⁺ to become NH₄⁺

176
Q

How can you test for ammonia?

A

It turns damp red litmus paper blue.

177
Q

Describe how you can test for ammonium ions.

A
  • Add sodium hydroxide and gently heat (this should produce ammonia gas if NH₄⁺ is present)
  • Test the gas with damp red litmus paper
  • Positive result -> Damp red litmus paper turns blue
178
Q

Give the ionic equation for the test for ammonium ions.

A

NH₄⁺(aq) + OH⁻(aq) -> NH₃(g) + H₂O(l)

The ammonia gas can then be tested for using damp red litmus paper.

179
Q

Why does litmus paper need to be damp?

A

So that the gas being tested for can dissolve and make the colour change.

180
Q

When testing for sulfates, why is HCl added before the barium chloride?

A

It is to get rid of any traces of carbonate ions that would produce a precipitate.

A LEVEL - Chemistry (Inorganic) (21 decks)

Brainscape helps you reach your goals faster, through stronger study habits.
© 2024 Bold Learning Solutions. Terms and Conditions