topic 4 - inorganic chemistry and the periodic table Flashcards
(120 cards)
1
Q
Why does ionisation energy decrease down Group 2?
A
Ionisation energy decreases down a group due to increasing atomic radius and electron shielding, reducing the effect of electrostatic forces of attraction.
2
Q
How does ionisation energy decrease down Group 2?
A
The number of shells increases down the group, increasing the distance between the nucleus and the outermost shell, meaning they are held more weakly, increasing atomic radius.
Increased nuclear charge is overridden by shielding and increased atomic radius [detailed exp. in topic 2]
3
Q
Why does reactivity increase down Group 2?
A
Reactivity increases due to increasing atomic radius and electron shielding, which decreases nuclear attraction, making it easier to remove two outer electrons and form cations.
4
Q
Why do melting points decrease down Group 2?
A
Metallic bonding weakens as ionic radius increases, increasing the distance between positive ions and delocalised electrons, weakening electrostatic forces and less energy needed to overcome forces
5
Q
What do Group 2 metals form when they react with water?
A
A metal hydroxide and hydrogen gas
6
Q
What is the equation for the reaction of a Group 2 metal with water?
A
M(s) + 2H2O(l) → M(OH)2 (aq) + H2(g)
7
Q
How does magnesium react with water?
A
Magnesium reacts with warm water, while other Group 2 metals react with cold water with increasing vigour down the group.
8
Q
What observations can be made when Group 2 metals react with water?
A
Alkaline solution
- Effervescence
- Metal dissolving
- Solution heating up
- White precipitate appearing with calcium (less likely down the group).
9
Q
What do Group 2 metals form when they react with steam?
A
A metal oxide and hydrogen gas.
10
Q
What is the equation for the reaction of a Group 2 metal with steam?
A
M(s) + H2O(g) → MO(s) + H2(g)
11
Q
What is observed when magnesium reacts with steam?
A
Magnesium burns with a bright white flame, meaning the reaction is much faster than with water.
12
Q
What do Group 2 metals form when they burn in oxygen?
A
Solid white oxides.
13
Q
What is the equation for the reaction of a Group 2 metal with oxygen?
A
2M(s) + O2(g) → 2MO(s)
14
Q
What is observed when magnesium burns in oxygen?
A
Magnesium burns with a bright white flame.
15
Q
Why does magnesium ribbon need to be cleaned before reactions?
A
Magnesium can react slowly with oxygen without a flame, forming a thin layer of magnesium oxide that must be removed by emery paper to avoid false results since magnesium and magnesium oxide react at different rates
16
Q
What do Group 2 metals form when they react with chlorine?
A
Solid white chlorides.
17
Q
What is the equation for the reaction of a Group 2 metal with chlorine?
A
M(s) + Cl2(g) → MCl2(s)
18
Q
How do Group 2 oxides react with water?
A
They readily react to form metal hydroxides, which dissolve to form strongly alkaline solutions due to the OH- ions
19
Q
Which Group 2 oxide does not react with water?
A
Beryllium oxide as it is insoluble
20
Q
Why is magnesium hydroxide not strongly alkaline?
A
It is sparingly soluble, and reacts slowly with water producing fewer free OH- ions, leading to a lower pH (~9).
21
Q
What is magnesium hydroxide used for in medicine?
A
It is used to neutralise excess stomach acid and treat constipation.
22
Q
Why is magnesium hydroxide preferable to calcium carbonate in medicine?
A
It does not produce carbon dioxide gas.
23
Q
What do Group 2 oxides and hydroxides do in dilute acids?
A
They neutralise the acid to form solutions of the corresponding salts.
24
Q
What is the equation for the reaction of a Group 2 oxide with water?
A
MO(s) + H2O(l) → M(OH)2 (aq)
25
What is the equation for the reaction of a Group 2 oxide with dilute acid?
MO(s) + 2HCl(aq) → MCl2(aq) + 2H2O(l)
26
What is the equation for the reaction of a Group 2 hydroxide with water?
M(OH)2(s) + 2H2O(l) → M(OH)2(aq)
27
What is the equation for the reaction of a Group 2 hydroxide with dilute acid?
M(OH)2(aq) + 2HCl(aq) → MCl2(aq) + 2H2O(l)
28
What is the solubility trend of Group 2 hydroxides?
They become more soluble going down the group. This allows them to produce more OH- ions, leading to a higher pH (~12), strongly alkaline solution.
The production of metal hydroxide is seen as a white precipitate, and its appearance will be limited if more soluble
29
What is the equation for the precipitation reaction of a Group 2 hydroxide?
M(NO3)2(aq) + 2NaOH(aq) → M(OH)2(s) + 2NaNO3(aq)
30
What is the solubility trend of Group 2 sulfates?
They become less soluble going down the group.
The production of metal sulfate is seen as a white precipitate, and its appearance will be limited if more soluble
31
What is the equation for the precipitation reaction of a Group 2 sulfate?
M(NO3)2(aq) + Na2SO4(aq) → MSO4(s) + 2NaNO3(aq)
32
Why is barium sulfate insoluble in water?
It forms a barrier on the metal surface, preventing further reaction.
33
How does charge affect solubility trends in Group 2 compounds?
Compounds with singly charged anions become more soluble down the group, whereas those with doubly charged anions become less soluble.
34
What is thermal decomposition?
It is when a reactant breaks down into more than one product when heated. The more thermally stable a substance is, the more heat it will take to break it down.
35
How does thermal stability change down Group 2?
Thermal stability increases down the group.
36
Why do carbonate and nitrate ions become unstable?
The carbonate and nitrate ions are large negative anions and can be made unstable by the presence of a positively charged cation. The cation polarises the anion, distorting it.
The greater the distortion, the less stable the compound.
37
Why does this trend of thermal stability occur down the group?
Going down the group, the larger the cations (ionic radius), the lower charge density (same charge 2+) so the less distortion (than small cations) as smaller polarising power of cation. The more stable the carbonate/nitrate compound as the C-O or N-O bond is weakened less so it less easily breaks down since it requires more energy.
38
Why are compounds with Group 2 cations less stable than Group 1 cations?
This is because the +1 charge on Group 1 metals (compared to the +2 charge on Group 2 metals) means they don’t have a big enough charge density to polarise the anion. Lithium is an exception as its ion is small enough to have a polarising effect.
39
What is the equation for the thermal decomposition of a Group 2 carbonate?
MCO3(s) → MO(s) + CO2(g)
40
What is the equation for the thermal decomposition of a Group 2 nitrate?
2M(NO3)2(s) → 2MO(s) + 4NO2(g) + O2(g)
41
Why does the thermal decomposition of Group 1 carbonates not occur?
Thermally stable – heating with a Bunsen burner to make them decompose is not sufficient as they do so at higher temperatures
An exception is Li2CO3, which decomposes to
Li2O and CO2
42
What is the equation for the thermal decomposition of a Group 1 nitrate?
2MNO3(s) → 2MNO2(s) + O2(g)
An exception is LiNO3 which decomposes to
form Li2O, NO2 and O2
43
How can you test the thermal stability of nitrates?
Measure time taken until a fixed volume of oxygen (relights a glowing splint) and brown gas (NO2) is produced; white nitrate solid is seen to melt into a colourless solution and resolidify
44
How can you test the thermal stability of carbonates?
Measure time taken for a fixed volume of carbon dioxide is produced, which is how long it takes for limewater to turn cloudy due to CO2 production; repeat for different carbonates using the same moles of carbonate/same volume of limewater/same Bunsen flame and height of tube above the flame
45
Why use a nichrome or platinum wire in flame tests?
They are unreactive and do not interfere with flame colours
46
Why dip the wire in hydrochloric acid before a flame test and place it in the hottest part (roaring/blue) of the Bunsen flame?
To clean it and ensure no residual colour interferes, so colour of flame not obscured. Also HCl is chosen to produce the metal chloride salt as more volatile than sulphites or nitrates
47
Why grind a sample before a flame test?
It allows the powder to stick to the damp wire and larger reaction surface area.
48
How do you observe the flame colour produced?
Dip the wire into the solid and put it in the hottest part of the Bunsen flame to observe the flame colour
If the colour is faint, you can dip the wire briefly into hydrochloric acid before returning it to the flame. This will give a short, intense burst of colour
49
Why do Group 1 and 2 compounds burn with different flame colours?
The colour is produced by electrons getting promoted to higher energy levels. The excited electrons then drop down to lower energy levels, emitting energy in the form of visible light with the wavelength of the observed light. The difference in energy between the higher and lower levels determines of the light released, determining the colour.
For each element, energy gaps and energy required to excite an electron vary
The wavelength of light depends on the element that the cation is formed from
50
Why might some elements not show a visible flame colour?
Some elements emit photons with energy outside the visible part of the electromagnetic spectrum.
51
What are the results of the flame tests for some metals?
- lithium: scarlet red
- sodium: yellow
- potassium: lilac
- rubidium: red
- caesium: blue
- magnesium: no flame colour; bright white flame as energy emitted of a wavelength outside visible spectrum
- calcium: brick red
- strontium: red
- barium: apple green
52
What are the physical states of halogens at room temperature?
Fluorine: pale yellow gas, highly reactive; Chlorine: pale green gas, poisonous in high concentrations; Bromine: red-orange liquid, dense brown/orange poisonous fumes; Iodine: shiny grey solid, sublimes to purple gas.
53
Why does ionisation energy decrease down the halogen group?
Ionisation energy decreases down a group due to an increasing atomic radius and electron shielding which reduces the effect of the electrostatic forces of attraction
54
Why does reactivity decrease down the halogen group?
Reactivity decreases due to increasing atomic radius and electron shielding, which decreases nuclear attraction, nuclear attraction decreases and it is harder to attract and electron and form anions – oxidising agent, where its oxidising power decreases down the group.
55
Why does electronegativity decrease down the halogen group?
Increased atomic radius, electron shielding and nuclear change which reduces the attraction for bonding electrons.
56
Why do halogen melting points increase down the group?
Larger molecules have more electrons as more electron shells, increasing London forces between molecules, requiring more energy to break.
57
What state do halogens naturally exist in?
Covalent diatomic molecules.
58
Why do halogens have low solubility in water?
They are non-polar, decreases down the group due to increased London forces and no hydrogen bonding (incompatible IMFs)
59
What are the colours of halogens in water?
- chlorine: virtually colourless
- bromine: yellow/orange
- iodine: brown
60
How does fluorine react with water?
2F2 + H2O → 4HF + O2, irreversible
61
How does chlorine react with water?
Cl₂ (g) + H₂O(l) ⇋ HCl(aq) + HClO(aq), a disproportionation reaction.
62
What happens when universal indicator is added to chlorine water?
Initially red due to acidity, then colourless as HClO bleaches the solution.
63
Why is chlorine used in water treatment?
Chlorine and hypochlorous acid effectively kill bacteria and viruses.
64
What is the equation for the ionisation of hypochlorous acid?
HClO (aq) + H₂O(l) ⇋ ClO⁻(aq) + H₃O⁺(aq)
65
How does bromine react with water?
Br₂ (g) + H₂O(l) ⇋ HBr(aq) + HOBr(aq), a disproportionation reaction with equilibrium further left than chlorine.
66
How soluble is iodine in water?
Insoluble/slightly soluble, forming I₃⁻ in the presence of I⁻ ions, I2 + I- ⇋ I3-
67
Why do halogens dissolve well in organic solvents like hexane?
Similar intermolecular forces (London forces).
68
How do you observe changes in colour in displacement reactions of halogens with halide ions?
The changes in colour are easier to see by shaking the [succcessful] reaction mixture with an organic solvent. The halogen that is present (solution formed) will dissolve in the organic
solvent, which settles out as a distinct layer above
69
What are the colours of halogens in hexane?
- chlorine: very pale green
- bromine: orange/red
- iodine: pink/violet
70
Why is studying fluorine and astatine difficult?
Fluorine is toxic, and astatine is highly radioactive and decays quickly.
71
What happens in halogen displacement reactions?
A more reactive halogen displaces a less reactive halide ion in solution. Allows observations of halogens’ relative oxidising power
72
What colour changes indicate displacement in halogen reactions?
* Metal halide solutions are colourless
* Colours of halogen water as expected
Chlorine is the most reactive halogen in the table and can displace both bromide and iodide ions, forming Br₂ (orange solution) and I₂ (brown solution) in displacement reactions:
Cl₂(aq) + 2Br⁻(aq) → 2Cl⁻(aq) + Br₂(aq) and Cl₂(aq) + 2I⁻(aq) → 2Cl⁻(aq) + I₂(aq).
Bromine is less reactive and can only displace iodide ions, forming I₂ (brown solution): Br₂(aq) + 2I⁻ (aq) → 2Br⁻(aq) + I₂(aq).
Iodine is the least reactive and does not react with any halide solution.
73
What do halogens form when reacting with Group 1 and 2 metals?
Halide salts.
74
What are the equations for halogen-metal reactions?
Group 1: 2M(s) + X₂(g) → 2MX(s); Group 2: M(s) + X₂(g) → MX₂(s).
75
What is the equation for the reaction of chlorine with cold, dilute alkali?
Cl2 (g) + 2NaOH(aq) → NaClO(aq) + NaCl(aq) + H2O(l)
76
What is sodium hypochlorite commonly known as?
Bleach, which has disinfectant properties because of the ClO- ion. It should never be used in conjunction with acidic cleaning products because the ClO- ion reacts with H+ ions to liberate the Cl2 gas.
77
Why is the chlorate(V) ion toxic?
It is a very strong oxidising agent.
78
What is the equation for the reaction of a halogen with hot, dilute alkali?
3X2 + 6NaOH → NaXO3 + 5NaX + 3H2O
79
What happens to the oxidation state of the halogen in reaction with hot, dilute alkali?
The halogen that is oxidised goes to a higher oxidation state.
80
What can a halide ion act as?
A reducing agent, where the reducing power increases going doing the group due to increased shielding effect and ionic radius.
81
Why can't acid-base reactions be redox reactions?
The oxidation states of the halide and sulfur do not change.
82
What is the equation for the reaction of chloride ions with sulfuric acid?
H2SO4 + KCl → KHSO4 + HCl
83
What is observed when chloride ions react with sulfuric acid?
Misty/steamy fumes of HCl.
84
Why does the reaction of chloride ions with sulfuric acid not proceed further?
HCl is a weak reducing agent and not strong enough to reduce sulfuric acid.
85
What role does sulfuric acid play in the reaction with chloride ions?
It acts as an acid.
86
What is the equation for the reaction of bromide ions with sulfuric acid?
H2SO4 + KBr → KHSO4 + HBr
87
What is observed when bromide ions react with sulfuric acid?
Misty/steamy fumes of HBr.
88
Why does the reaction of bromide ions with sulfuric acid proceed further?
HBr is a strong enough reducing agent to react with sulfuric acid in a redox reaction.
89
What role does sulfuric acid play in the second step of bromide reaction?
It acts as an oxidising agent.
90
What is the equation for the redox reaction of bromide ions with sulfuric acid?
H2SO4 + 2H+ + 2Br- → SO2 + 2H2O + Br2
91
What is observed when bromide ions undergo a redox reaction with sulfuric acid?
Red-brown fumes of Br2 and a colourless, pungent gas (SO2).
92
What is the equation for the reaction of iodide ions with sulfuric acid?
H2SO4 + KI → KHSO4 + HI
93
What is observed when iodide ions react with sulfuric acid?
Misty/steamy fumes of HI.
94
Why does the reaction of iodide ions with sulfuric acid proceed further multiple times?
HI is a stronger reducing agent than HBr and continues to reduce sulfuric acid in multiple redox steps.
95
What role does sulfuric acid play in the reactions with iodide ions?
It acts as an oxidising agent.
96
What is the equation for the redox reaction of iodide ions forming iodine?
H2SO4 + 2H+ + 2I- → SO2 + 2H2O + I2
97
What is observed when iodide ions react to form iodine?
Purple fumes of I2.
98
What is the equation for the redox reaction of iodide ions forming sulfur?
H2SO4 + 6H+ + 6I- → S + 4H2O + 3I2
99
What is observed when iodide ions react to form sulfur?
A yellow solid (sulfur).
100
What is the equation for the redox reaction of iodide ions forming hydrogen sulfide?
H2SO4 + 8H+ + 8I- → H2S + 4H2O + 4I2
101
What is observed when iodide ions react to form hydrogen sulfide?
A colourless gas with a rotten egg smell (H2S).
102
What is the reaction of hydrogen halides with water?
Hydrogen halides are colourless gases. They react with water to produce misty fumes and dissolve to form acidic solutions: HX(g) + (aq) → H+ (aq) + X- (aq) and HX(g) + H2O(l) → H3O+ (aq) + X- (aq).
103
What is the reaction of hydrogen halides with ammonia?
They react with ammonia gas to form white fumes: HX(g) + NH3(g) → NH4X(s), which is used as a test for hydrogen halides.
104
How are hydrogen halides made?
They are made by reacting solid sodium halide salts with phosphoric acid: NaCl(s) + H3PO4(l) → NaH2PO4(s) + HCl(g). White steamy fumes form as HCl dissolves in moisture in the air.
105
Why is phosphoric acid used instead of sulfuric acid to make hydrogen halides?
Phosphoric acid is not an oxidising agent and does not cause redox reactions with HBr and HI, making it more suitable than sulfuric acid for producing hydrogen halides.
106
What is the test for halide ions?
Silver ions react with halide ions to form a precipitate: Ag+ (aq) + X- (aq) → AgX (s). First, dilute nitric acid is added to remove carbonate ions that could form a pale precipitate and interfere with results. This is followed by silver nitrate solution, so a precipitate of the silver halide is formed. Hydrochloric acid cannot be used as it contains chloride ions, which would form a precipitate.
107
What are the results of the halide test with silver nitrate?
Chloride (Cl-) forms a white precipitate (AgCl), which is soluble in dilute ammonia. Bromide (Br-) forms a cream precipitate (AgBr), which is insoluble in dilute ammonia but soluble in concentrated ammonia. Iodide (I-) forms a primrose yellow precipitate (AgI), which is insoluble in both dilute and concentrated ammonia.
108
What are the equations for the halide precipitation reactions and solubility tests?
Chloride: Cl-(aq) + Ag+(aq) → AgCl(s), AgCl(s) + 2NH3(g) → Ag(NH3)2+(aq) + Cl-(aq). Bromide: Br-(aq) + Ag+(aq) → AgBr(s). Iodide: I-(aq) + Ag+(aq) → AgI(s).
109
Why might halide precipitates require further testing?
The precipitates can look similar, making visual identification difficult. Adding ammonia solution helps distinguish them: AgCl dissolves in dilute ammonia, AgBr dissolves in concentrated ammonia, and AgI is insoluble in both.
110
What is the test for carbonate ions (CO32-) and hydrogencarbonate ions (HCO3-)?
The presence of carbon dioxide is observed through bubbling the gas produced through a test tube of limewater.
Ca(OH)2 (aq) + CO2 (g) → CaCO3 (s) + H2O(l)
111
What is the chemical equation for the reaction when carbonate ions (CO32-) are tested?
2H+(aq) + CO32-(s) → CO2(g) + H2O(l)
112
What is the chemical equation for the reaction when hydrogencarbonate ions (HCO3-) are tested?
H+(aq) + HCO3-(s) → CO2(g) + H2O(l)
113
What reagent is used to test for carbonate or hydrogencarbonate ions (CO32- or HCO3-)?
Dilute acid (HCl(aq)), added to a solid sample.
114
What is the result of the test for carbonate or hydrogencarbonate ions (CO32- or HCO3-)?
Carbon dioxide is released, producing bubbles or effervescence. The gas turns limewater cloudy.
115
What is the test for sulfate ions (SO42-)?
Add dilute hydrochloric acid (HCl) remove carbonate ions that could form a pale precipitate and interfere with results, then add barium chloride solution (BaCl2). Sulfuric acid can’t be used to acidify the mixture because it would form a precipitate
116
What is the result of the test for sulfate ions (SO42-)?
A white precipitate of barium sulfate (BaSO4) forms, which does not dissolve in dilute acid.
117
What is the chemical equation for the reaction when sulfate ions (SO42-) are tested?
Ba2+(aq) + SO42-(aq) → BaSO4(s)
118
What reagent is used to test for ammonium compounds (NH4+)?
Aqueous sodium hydroxide (NaOH), warmed.
119
What is the result of the test for ammonium compounds (NH4+)?
Ammonia gas is released. Gas produced dissolves in the water on a damp piece of red litmus paper and turns it blue (blue litmus paper stays blue), since it is alkaline
120
What is the chemical equation for the reaction when ammonium ions (NH4+) are tested with sodium hydroxide (NaOH)?
NH4+(aq) + OH-(aq) → NH3(g) + H2O(l)
