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
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)
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
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
5
Q
What 3 important things can group 2 elements react with?
A
- Water
- Oxygen
- Chlorine
6
Q
What is formed when group 2 metals react with water?
A
• Metal hydroxide
• Hydrogen
(Except beryllium)
7
Q
Give the general equation for group 2 metals reacting with water (where metal = M).
A
M + H2O -> M(OH)2 + H2
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.
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)
10
Q
Describe how magnesium reacts with:
1) Water
2) Steam
A
Water: • Slowly • Produces magnesium hydroxide Steam: • Quickly • Produces magnesium oxide
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.
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
13
Q
What is formed when group 2 metals react with oxygen?
A
Metal oxides
14
Q
Give the general equation for group 2 metals reacting with oxygen (where metal = M).
A
2M + O₂ -> 2MO
15
Q
What is formed when group 2 metals react with chlorine?
A
Metal chlorides
16
Q
Describe the appearance of metal chlorides.
A
White solids
17
Q
Give the general equation for group 2 metals reacting with chlorine (where metal = M).
A
M + Cl2 -> MCl2
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)
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
20
Q
What acidity are group 2 metal hydroxides?
A
Alkaline.
21
Q
What is formed when group 2 oxides react with water?
A
Metal hydroxides (except beryllium).
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.
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
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.
25
What makes metal hydroxide solutions alkaline?
The OH- ions
26
What happens when group 2 metal oxides react with water?
• Metal hydroxide is formed
• This metal hydroxide dissolves
• Alkaline solution is formed
(Except with beryllium. Magnesium is very slow and hardly soluble.)
27
What acidity are group 2 metal oxides?
Alkaline.
28
Give the general equation for group 2 metal oxides reacting with water (where M = metal).
MO (s) + H2O (l) -> M(OH)2 (aq)
29
Give the general equation for SOLID group 2 metal hydroxides reacting with water (where M = metal).
M(OH)2 (s) + H2O (l) -> M(OH)2 (aq)
30
What happens when solid group 2 metal hydroxides are added to water?
They dissolve (if they are soluble) to give an alkaline solution.
31
What is formed when group 2 metal oxides react with dilute acid?
* Salt
| * Water
32
Give the general equation for metal oxides reacting with dilute hydrochloric acid (where M = metal).
MO (s) + 2HCl (aq) -> MCl2 (aq) + H2O (l)
33
What is formed when group 2 metal hydreoxides react with dilute acid?
* Salt
| * Water
34
Give the general equation for metal hydroxides reacting with dilute hydrochloric acid (where M = metal).
M(OH)2 (s) + 2HCl (aq) -> MCl2 (aq) + 2H2O (l)
35
Are there solubility trends for group 2 metal compounds?
Yes, but they depends on the compound anion (e.g. OH-).
36
Describe how the solubility of group 2 metal hydroxides changes going down the group.
It increases.
37
Describe how the solubility of group 2 metal sulfates changes going down the group.
It decreases.
38
Which group 2 compound can be used as a reference for all other group 2 compound solubilities?
* 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.
39
What's more soluble, calcium hydroxide or strontium hydroxide?
Strontium hydroxide
40
What's more soluble, calcium sulfate or strontium sulfate?
Calcium sulfate
41
What are compounds with very low solubilities said to be?
Sparingly soluble
42
What is thermal decomposition?
When a substance breaks down when heated.
43
What is thermal stability?
The amount of heat required to thermally decompose a substance.
44
What is a substance with high thermal stability?
A substance that requires a lot of heat to be thermally decomposed.
45
How does thermal stability of group 1 and 2 compounds change going down the group and why?
• 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
46
How does a cation's charge density affect the thermal stability of its compound?
High charge density -> High polarising power -> Strong distortion of anion -> Thermally unstable compound
47
Compare the thermal stability of group 1 and 2 compounds.
• 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
48
What are the general trends for the thermal stability of group 1 and 2 compounds?
* Thermal stability increases down a group
| * Group 1 compounds are more thermally stable than group 2 compounds
49
Describe what happens when you heat group 1 carbonates.
* 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
50
Describe what happens when you heat group 1 nitrates.
* Decompose to the metal nitrite and oxygen
| * Except LiNO3, which decomposes to form Li2O, NO2 and O2.
51
Describe what happens when you heat group 2 carbonates.
• Decompose to the metal oxide and carbon dioxide
52
Describe what happens when you heat group 2 nitrates.
• Decompose to the metal oxide, nitrogen dioxide and oxygen.
53
Give the general equation for the thermal decomposition of group 1 carbonates (where M = metal).
They do not thermally decompose easily, except Li2CO3.
54
Give the general equation for the thermal decomposition of group 1 nitrates (where M = metal).
2MNO3 (s) -> 2MNO2 (s) + O2 (g)
55
Give the general equation for the thermal decomposition of group 2 carbonates (where M = metal).
MCO3 (s) -> MO (s) + CO2 (g)
56
Give the general equation for the thermal decomposition of group 2 nitrates (where M = metal).
2M(NO3)2 (s) -> 2MO (s) + 4NO2 (g) + O2 (g)
57
Which group 1 or 2 metal is an exception to the thermal decomposition rules and how does it behave?
* Lithium
| * Its carbonate and nitrate decompose like group 2 carbonates and nitrates
58
What are the products of the thermal decomposition of lithium carbonate?
Lithium oxide and carbon dioxide
| Remember: This is an exception!
59
What are the products of the thermal decomposition of lithium nitrate?
Lithium oxide, nitrogen dioxide and oxygen
| Remember: This is an exception!
60
How can you test how easily a nitrate decomposes?
Time how long it takes to produce a certain amount of either:
• Oxygen
• Brown gas (NO2) -> In a fume cupboard
61
How can you test how easily a carbonate decomposes?
Time how long it takes to produce a certain amount of:
| • Carbon dioxide
62
Name the flame colour of lithium.
Red
63
Name the flame colour of sodium.
Orange/Yellow
64
Name the flame colour of potassium.
Lilac
65
Name the flame colour of rubidium.
Red
66
Name the flame colour of caesium.
Blue
67
Name the flame colour of calcium.
Brick red
68
Name the flame colour of strontium.
Crimson
69
Name the flame colour of barium.
Green
70
```
Name the flame colour of these group 1 metals:
• Li
• Na
• K
• Rb
• Cs
```
* Li - Red
* Na - Orange/Yellow
* K - Lilac
* Rb - Red
* Cs - Blue
71
Name the flame colour of these group 2 metals:
• Ca
• Sr
• Ba
* Ca - Brick red
* Sr - Crimson
* Ba - Green
72
Describe how to carry out a flame test.
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
Explain why different elements have distinct flame colours.
* 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
What is the movement of electrons between energy levels called?
Electron transition
75
What are the first 4 halogens?
* Fluorine
* Chlorine
* Bromine
* Iodine
76
How does electronegativity change as you go down the halogens?
It decreases.
77
Describe the colour and state of fluorine at RTP.
* Pale yellow
| * Gas
78
Describe the colour and state of chlorine at RTP.
* Green
| * Gas
79
Describe the colour and state of bromine at RTP.
* Red-brown
| * Liquid
80
Describe the colour and state of iodine at RTP.
* Grey
| * Solid
81
Describe and explain the solubility of halogens in water and organic solvents.
* In water -> Low
* In organic solvents -> High
Because they’re non-polar.
82
Describe the colour of these halogens in water:
• Chlorine
• Bromine
• Iodine
* Chlorine -> Colourless
* Bromine -> Yellow/Orange
* Iodine -> Brown
83
Describe the colour of these halogens in hexane:
• Chlorine
• Bromine
• Iodine
* Chlorine -> Colourless
* Bromine -> Orange/Red
* Iodine -> Pink/Violet
84
What is the colour of fluorine:
| • At RTP
• At RTP -> Pale yellow gas
85
What is the colour of chlorine:
• At RTP
• In water
• In hexane
* At RTP -> Green gas
* In water -> Colourless
* In hexane -> Colourless
86
What is the colour of bromine:
• At RTP
• In water
• In hexane
* At RTP -> Red-brown liquid
* In water -> Yellow/Orange
* In hexane -> Orange/Red
87
What is the colour of iodine:
• At RTP
• In water
• In hexane
* At RTP -> Grey solid
* In water -> Brown
* In hexane -> Pink/Violet
88
Explain how the reactivity of halogens changes as you go down the group.
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
Explain how the melting and boiling points of halogens change as you go down the group.
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
Why is the chemistry of fluorine and astatine difficult to study?
* Fluorine -> Toxic has
| * Astatine -> Radioactive and so it decays quickly
91
What happens when you add a solution of a halogen to a solution of a hydrogen halide?
* If the halogen is more reactive -> It will displace the halide from its solution
* If the halogen is less reactive -> No reaction
92
Which halides will chlorine displace from their solution?
Bromide and iodide
93
Which halides will bromine displace from their solution?
Iodide
94
Which halides will iodine displace from their solution?
None
95
What type of reaction is the displacement reaction between a halogen and a halide?
Redox
96
In a halogen-halide reaction, is the halogen reduced or oxidised?
Reduced
97
In a halogen-halide reaction, is the halide reduced or oxidised?
Oxidised
98
Remember to practise writing out ionic equations for halogen-halide reactions.
Pg 47 of revision guide
99
In a halogen-halide reaction, what can be done to make the colour change more obvious?
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
How do halogens react with group 1 and 2 metals?
They produce halide salts.
101
2Li + F₂ ->
2Li + F₂ -> 2LiF
102
Mg + Cl₂ ->
Mg + Cl₂ -> MgCl₂
103
What sort of reaction do halogens undergo with alkalis?
Disproportionation
104
What is it important to remember about halogen reactions with alkalis?
The products depend on whether the alkali is hot or cold.
105
Give the general equation for a halogen reacting with a cold alkali.
(Use NaOH as the alkali)
X₂ + 2NaOH -> NaOX + NaX + H₂O
106
When halogen reacts with a cold alkali, what happens to the oxidation number of the halogen?
0 (X₂) -> +1 (NaOX) and -1 (NaX)
107
I₂ + 2NaOH ->
| When the alkali is cold
I₂ + 2NaOH -> NaOI + NaI + H₂O
108
Name: NaOI
Sodium iodate(I)
109
Describe the oxidation states that halogens can exist in.
They can exist in many. For example: Br, BrO⁻, BrO₂⁻
110
What is the chemical name and formula for bleach?
* Sodium chlorate(I) solution
| * NaClO
111
Describe how you can prepare bleach.
Mix chlorine gas with cold, dilute aqueous sodium hydroxide.
112
What type of reaction is the preparation of bleach?
Disproportionation
113
Give the equation for the formation of bleach.
2NaOH + Cl₂ -> NaClO + NaCl + H₂O
| Note: The NaOH must be cold
114
What are some uses of sodium chlorate(I)?
It is bleach, so:
• Water treatment
• Bleaching paper and textiles
• Cleaning toilets
115
Give the general equation for a halogen reacting with a hot alkali.
(Use NaOH as the alkali)
3X₂ + 6NaOH -> NaXO₃ + 5NaX + 3H₂O
116
When halogen reacts with a hot alkali, what happens to the oxidation number of the halogen?
0 (X₂) -> +5 (NaXO₃) and -1 (NaX)
117
Cl₂ + NaOH ->
| When the alkali is hot
3Cl₂ + 6NaOH -> NaClO₃ + 5NaCl + 3H₂O
118
Name: NaClO₃
Sodium chlorate(V)
119
What is chlorine used for?
Killing bacteria in water.
120
What type of reaction is mixing chlorine with water!
Disproportionation
121
What is produced when you mix chlorine with water?
* Hydrochloric acid
| * Hypochlorous acid
122
Give the chemical formula for hypochlorous acid.
HClO
123
What is the equation for the reaction between chlorine and water?
Cl₂ + H₂O {-} HCl + HClO
124
When chlorine reacts with water, what happens to the oxidation number of the halogen?
0 (Cl₂) -> -1 (HCl) and +1 (HClO)
125
Do other halogens except chlorine also undergo disproportionation when mixed with water?
Yes
126
What happens to hypochlorous acid in water?
It ionises to make chlorate(I) ions and H₃O⁺.
127
Give the equation for what happen to hypochlorous acid in water.
HClO + H₂O {-} ClO⁻ + H₃O⁺
128
Give the chemical formula for chlorate(I) ions.
ClO⁻
129
What is another name for chlorate(I) ions? (ClO⁻)
Hypochlorite ions
130
Which part of chlorine causes it to kill bacteria?
The chlorate(I) (ClO⁻) ions it Eventually forms.
131
Give the entire pathway for how chlorine kills bacteria in water.
• 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
How can halides act as reducing agents?
They can lose an electron from the outer shell.
133
Explain how the reducing power of halides changes as you go down the group.
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
Do all group 1 halides react the same with sulfuric acid?
* 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
```
Describe the number of times these will react with sulfuric acid:
• KF
• KCl
• KBr
• KI
```
* KF + KCl -> Once
* KBr -> Twice
* KI -> Three times
136
Give the equations for the reactions of KF or KCl with H₂SO₄. Include state symbols.
KF(s) + H₂SO₄(l) -> KHSO₄(s) + HF(g)
KCl(s) + H₂SO₄(l) -> KHSO₄(s) + HCl(g)
137
KF(s) + H₂SO₄(l) -> KHSO₄(s) + HF(g)
KCl(s) + H₂SO₄(l) -> KHSO₄(s) + HCl(g)
What is observed in these reactions?
Misty fumes (as the gas comes into contact with moisture)
138
KF(s) + H₂SO₄(l) -> KHSO₄(s) + HF(g)
KCl(s) + H₂SO₄(l) -> KHSO₄(s) + HCl(g)
Are these redox reactions?
No, because the oxidation numbers stay the same.
139
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?
The F⁻ and Cl⁻ ions are not strong enough reducing agents to reduce the sulfuric acid.
140
Give the equations for the reactions of KBr with H₂SO₄. Include state symbols.
* KBr(s) + H₂SO₄(l) -> KHSO₄(s) + HBr(g)
| * HBr(g) + H₂SO₄(l) -> Br₂(g) + SO₂(g) + 2H₂O(l)
141
* KBr(s) + H₂SO₄(l) -> KHSO₄(s) + HBr(g)
* HBr(g) + H₂SO₄(l) -> Br₂(g) + SO₂(g) + 2H₂O(l)
What is observed in these reactions?
```
First reaction:
• Misty fumes of HBr gas
Second reaction:
• Choking fumes of SO₂
• Orange fumes of Br₂
```
142
* KBr(s) + H₂SO₄(l) -> KHSO₄(s) + HBr(g)
* HBr(g) + H₂SO₄(l) -> Br₂(g) + SO₂(g) + 2H₂O(l)
Are these redox reactions?
First reaction:
• Not redox
Second reaction:
• Redox
143
* 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?
The Br⁻ ions are only strong enough reducing agents to reduce the sulfuric acid once.
144
Give the equations for the reactions of KI with H₂SO₄. Include state symbols.
* 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
* 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?
```
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
* 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?
First reaction:
• Not redox
Second and third reaction:
• Redox
147
Why does KI react with H₂SO₄ three times?
The iodide ions are such powerful reducing agents.
148
What is the indicator of each group 1 halide reacting with H₂SO₄?
* 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
What is an indicator of H₂S?
The smell of rotten eggs.
| It is also toxic.
150
What types of chemical are hydrogen halides?
Acidic gases
151
What happens when hydrogen halides come in contact with water?
* They dissolve
* Misty fumes of acidic gas are produced
* This would turn damp blue litmus paper red
152
What happens when hydrogen halides come in contact with ammonia gas?
They react to give white fumes.
153
What is the equation for ammonia gas reacting with HCl gas?
What changes are seen?
NH₃(g) + HCl(g) -> NH₄Cl(s)
This gives white fumes.
154
Describe the rest for halide ions.
* Add dilute nitric acid
* Add silver nitrate solution
* Check the colour of the precipitate
155
When testing for halides, why do you add dilute nitric acid?
To remove ions which might interfere with the reaction.
156
What type of precipitate is formed in the rest for halide ions?
Silver halide
157
When testing for halides, what is the positive result for fluoride ions?
No precipitate
158
When testing for halides, what is the positive result for chloride ions?
White precipitate
159
When testing for halides, what is the positive result for bromide ions?
Cream precipitate
160
When testing for halides, what is the positive result for iodide ions?
Yellow precipitate
161
When testing for halides, what is the positive result for each of the halides?
* Fluorine -> No precipitate
* Chlorine -> White precipitate
* Bromine -> Cream precipitate
* Iodide -> Yellow precipitate
162
When testing for halides, how can you tell apart the different precipitates if the colour is not obvious?
Add ammonia and see if the precipitate dissolves.
163
After testing for halides, what happens to the silver chloride precipitate when ammonia is added?
It dissolves in DILUTE ammonia solution to give a colourless solution.
164
After testing for halides, what happens to the silver bromide precipitate when ammonia is added?
* Remains unchanged when DILUTE ammonia is added
| * Dissolves when CONCENTRATED ammonia is added to give a colourless solution
165
After testing for halides, what happens to the silver bromide precipitate when ammonia is added?
It does not dissolve, even in concentrated ammonia solution.
166
Give the ionic equation for the test for halides.
Ag⁺(aq) + X⁻(aq) -> AgX(s)
168
What is the formula for hydrogencarbonate?
HCO₃⁻
169
Describe how you can test for carbonates and hydrogencarbonates.
* Add HCl
* Bubble the gas through limewater
* Positive result -> Limewater turns cloudy
170
Give the ionic equation for the test for carbonate ions.
CO₃²⁻ + 2H⁺ -> CO₂ + H₂O
171
Give the ionic equation for the test for hydrogencarbonate ions.
HCO₃²⁻ + H⁺ -> CO₂ + H₂O
172
What is the formula for a sulfate ion?
SO₄²⁻
173
Describe how you can test for sulfates.
* Add dilute HCl
* Add barium chloride solution (BaCl₂)
* Positive result -> White precipitate of barium sulfate
174
Give the ionic equation for the test for sulfates.
Ba²⁺(aq) + SO₄²⁻(aq) -> BaSO₄(s)
175
What is the pH of ammonia gas?
Alkaline
| Note: It accepts H⁺ to become NH₄⁺
176
How can you test for ammonia?
It turns damp red litmus paper blue.
177
Describe how you can test for ammonium ions.
* 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
Give the ionic equation for the test for ammonium ions.
NH₄⁺(aq) + OH⁻(aq) -> NH₃(g) + H₂O(l)
The ammonia gas can then be tested for using damp red litmus paper.
179
Why does litmus paper need to be damp?
So that the gas being tested for can dissolve and make the colour change.
180
When testing for sulfates, why is HCl added before the barium chloride?
It is to get rid of any traces of carbonate ions that would produce a precipitate.
