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

Topic 4 - Inorganic Chemistry and the Periodic Table Flashcards

1
Q

Name the first 5 elements in group 2.

A
  • Beryllium
  • Magnesium
  • Calcium
  • Strontium
  • Barium
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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)

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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

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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

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5
Q

What 3 important things can group 2 elements react with?

A
  • Water
  • Oxygen
  • Chlorine
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6
Q

What is formed when group 2 metals react with water?

A

• Metal hydroxide
• Hydrogen
(Except beryllium)

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7
Q

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

A

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

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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.
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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)

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10
Q

Describe how magnesium reacts with:

1) Water
2) Steam

A 
Water:
• Slowly
• Produces magnesium hydroxide
Steam:
• Quickly
• Produces magnesium oxide
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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.

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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
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13
Q

What is formed when group 2 metals react with oxygen?

A

Metal oxides

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14
Q

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

A

2M + O -> 2MO

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15
Q

What is formed when group 2 metals react with chlorine?

A

Metal chlorides

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16
Q

Describe the appearance of metal chlorides.

A

White solids

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17
Q

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

A

M + Cl2 -> MCl2

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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)
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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
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20
Q

What acidity are group 2 metal hydroxides?

A

Alkaline.

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21
Q

What is formed when group 2 oxides react with water?

A

Metal hydroxides (except beryllium).

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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.

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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

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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.

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25
Q

What makes metal hydroxide solutions alkaline?

A

The OH- ions

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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.)

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27
Q

What acidity are group 2 metal oxides?

A

Alkaline.

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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)

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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)

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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.

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31
Q

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

A
  • Salt

* Water

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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)

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33
Q

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

A
  • Salt

* Water

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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)

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35
Q

Are there solubility trends for group 2 metal compounds?

A

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

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36
Q

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

A

It increases.

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37
Q

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

A

It decreases.

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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.
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39
Q

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

A

Strontium hydroxide

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40
Q

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

A

Calcium sulfate

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41
Q

What are compounds with very low solubilities said to be?

A

Sparingly soluble

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42
Q

What is thermal decomposition?

A

When a substance breaks down when heated.

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43
Q

What is thermal stability?

A

The amount of heat required to thermally decompose a substance.

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44
Q

What is a substance with high thermal stability?

A

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

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45
Q

How does thermal stability of group 1 and 2 compounds decrease 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

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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

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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

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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

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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
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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.

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51
Q

Describe what happens when you heat group 2 carbonates.

A

• Decompose to the metal oxide and carbon dioxide

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52
Q

Describe what happens when you heat group 2 nitrates.

A

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

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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.

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54
Q

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

A

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

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55
Q

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

A

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

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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)

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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

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58
Q

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

A

Lithium oxide and carbon dioxide

Remember: This is an exception!

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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!

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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

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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

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62
Q

Name the flame colour of lithium.

A

Red

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63
Q

Name the flame colour of sodium.

A

Orange/Yellow

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64
Q

Name the flame colour of potassium.

A

Lilac

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65
Q

Name the flame colour of rubidium.

A

Red

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66
Q

Name the flame colour of caesium.

A

Blue

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67
Q

Name the flame colour of calcium.

A

Brick red

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68
Q

Name the flame colour of strontium.

A

Crimson

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69
Q

Name the flame colour of barium.

A

Green

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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
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71
Q

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

A
  • Ca - Brick red
  • Sr - Crimson
  • Ba - Green
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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 bromine:
• 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?

94
Q

Which halides will iodine displace from their solution?

95
Q

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

96
Q

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

97
Q

In a halogen-halide reaction, is the halide reduced or 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.

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?

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.

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?

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 (Organic) (21 decks)

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