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A Level Chemistry > Topic 4: Inorganic Chemistry and the Periodic Table > Flashcards

Topic 4: Inorganic Chemistry and the Periodic Table Flashcards

1
Q

Why does ionisation energy of group 2 elements decrease down the group?

A
  • each element down Group 2 has an extra electron shell compared to one above
  • extra inner shells shield the outer electrons from the attraction of the nucleus
  • outer electrons are also further away from the nucleus, reducing electrostatic attraction between the nucleus and outer electrons
  • the positive charge does increase down the group, but the top two factors override
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2
Q

Does reactivity increase or decrease down group 2?

A
  • reactivity increases down the group
  • as atomic radii increase there is more shielding
  • the nuclear attraction decreases and it is easier to remove outer electrons and so cations form more easily
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3
Q

What do group 2 metals react with water to make? Write a general equation

A
  • M(s) + 2H2O(l) → M(OH)2(aq) + H2(g)
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4
Q

Which group 2 metals react with cold water? And which doesn’t?

A
  • Ca, Sr, Ba react with cold water to form hydroxides
  • Mg needs to react with warm water to form its hydroxide
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5
Q

What would you observe with the formation of a group two hydroxide?

A
  • fizzing (more vigorous down the group)
  • metal dissolving (faster down group)
  • solution heating up (more down group)
  • with calcium: a white precipitate appearing
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6
Q

Describe the reaction between magnesium and steam

A
  • forms magnesium oxide and hydrogen (Mg reacts differently in steam)
  • Mg will burn with a bright white flame
  • Mg(s) + H2O(g) → MgO(s) and H2(g)
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7
Q

What do group 2 metals react with oxygen to make?

A
  • solid white oxides
  • Mg with a bright white flame
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8
Q

What do group 2 metals react with chlorine to make?

A
  • solid white chlorides
  • M(s) + Cl2(g) → MCl2(s)
  • more vigorous down group
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9
Q

Why are group 2 oxides classed as basic oxides?

A
  • the oxide ions accept protons to become hydroxide ions
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10
Q

Describe the reaction of group 2 oxides with water

A
  • MO(s) + H2O(l) → M(OH)2(aq)
  • forms colourless solutions
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11
Q

Which metal oxides are the exceptions to being bases?

A
  • beryllium oxide: doesn’t react with water and is insoluble
  • magnesium oxide: reacts slowly and OH isn’t very soluble
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12
Q

Describe the reactions of group 2 oxides and hydroxides with acids

A
  • neutralisation reaction
  • salt and water
  • white solid to colourless solution
  • exothermic reaction
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13
Q

What is the trend in solubility of group 2 hydroxides?

A
  • group 2 hydroxides become more soluble down the group
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14
Q

What is calcium hydroxide used for?

A
  • Testing for CO2:
  • a saturated solution of calcium hydroxide is limewater
  • CO2 + Ca(OH)2 → CaCO3 + H2O
  • white precipitate forms
  • Neutralising soils in agriculture
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15
Q

How is magnesium hydroxide used?

A
  • used in medicine to neutralise excess stomach acid (HCl)
  • it is safe to use because the very low solubility of magnesium hydroxide means it is only weakly alkaline
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16
Q

What is the trend of solubility of group 2 sulfates?

A
  • group 2 sulfates become less soluble down the group
  • magnesium sulfate is classed as soluble
  • calcium sulfate slightly soluble
  • strontium and barium sulfate insoluble
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17
Q

How do you test for sulfate ions?

A
  • add barium ions (usually barium chloride or nitrate)
  • forms a white precipitate of barium sulfate
  • carbonate ions can also form white precipitate with barium ions, so there must be H+ ions to prevent it from forming
  • add dilute nitric acid or dilute HCl
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18
Q

What is thermal decomposition?

A
  • the use of heat to break down a reactant into more than one product
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19
Q

What is the trend of thermal stability down group 1 and group 2 compounds?

A
  • thermal stability increases down a group
  • carbonate and nitrate ions are large negative ions (anions) and can be made unstable by the presence of a cation
  • the cation polarises the anion, distorting it
  • the greater the distortion the less stable the compound
  • large cations cause less distortion than small cations as they have a lower charge density
  • so further down the group, the larger the cations, the lower the charge density so the less distortion caused and the more stable the carbonate/nitrate compound
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20
Q

Are group 1 or group 2 compounds less thermally stable?

A
  • group 2 compounds are less thermally stable than group 1
  • the greater the charge on the cation, the greater the distortion and the less stable the carbonate/nitrate compound becomes
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21
Q

What do group 1 carbonates decompose to?

A
  • they are thermally stable, so they do not decompose with a Bunsen flame
  • except for lithium carbonate because it is small enough to have a polarising effect
  • Li2CO3(s) → Li2O(s) + CO2(g)
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22
Q

What do group 1 nitrates decompose to?

A
  • form metal nitrite (nitrate(III) salt) and oxygen
  • e.g. 2NaNO3 → 2NaNO2 +O2
  • except lithium nitrate which decomposes like group 2 nitrates to form oxide, oxygen and nitrogen dioxide gas
  • 4LiNO3 → 2Li2O + 4NO2 + O2
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23
Q

What do group 2 carbonates decompose to?

A
  • forms group 2 oxides and CO2
  • e.g. MgCO3(s) → MgO(s) + CO2(g)
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24
Q

What do group 2 nitrates decompose to? State the observations

A
  • form the oxide, nitrogen dioxide and oxygen
  • brown fumes of nitrogen dioxide and white nitrate solid is seen to melt to a colourless solution then resolidify
  • e.g. 2Mg(NO3)2 → 2MgO + 4NO2 + O2
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25
Q

How can you test the thermal stability of nitrates?

A
  • how long it takes until a certain amount of oxygen is produced
  • how long it takes until NO2, a brown gas, is produced
  • needs fume cupboard
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26
Q

How to test thermal stability of carbonates?

A
  • how long it takes for an amount of CO2 to be produced
  • test using lime water which turns cloudy with CO2
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27
Q

What flame colour is lithium, sodium, potassium, rubidium, caesium, calcium, strontium, barium?

A
  • Li: red
  • Na: orange/yellow
  • K: lilac
  • Rb: red
  • Cs: Blue
  • Ca: brick-red
  • Sr: crimson
  • Ba: green
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28
Q

How to do a flame test

A
  • use a nichrome wire as it is unreactive and will not give a flame colour
  • dip in concentrated HCl to sterilise and heat in Bunsen flame
  • dip wire in solid and put in Bunsen flame and observe flame
29
Q

Why do flames have colour?

A
  • energy absorbed from the flame causes electrons to move to higher energy levels (from ground state to excited state)
  • colours are seen as electrons fall back down to lower energy levels, releasing energy in the form of light
  • the difference in energy between higher and lower levels determines wavelength of light released
  • this determines colour of the light
30
Q

What are the colours, states of the first four halogens?

A
  • fluorine: very pale yellow gas
  • chlorine: greenish gas
  • bromine: red liquid, that gives off brown, orange fumes
  • iodine: shiny grey solid sublimes to purple gas
31
Q

What are the physical states at room temperature of the first 4 halogens?

A
  • F2: gas
  • Cl2: gas
  • Br2: liquid
  • I2: solid
32
Q

What is the trend in electronegativity in halogen and why?

A
  • electronegativity decreases down the group
  • the atomic radii increases due to the increasing number of shells
  • nucleus is therefore less able to attract the bonding pair of electrons
33
Q

What are the trends in melting and boiling points of halogens?

A
  • they increase down the group
  • as the molecules become large, they have more electrons and so have larger London forces between the molecules
  • as the intermolecular forces increase, more energy has to be put in to break the forces
34
Q

What is the trend in reactivity down the halogens group?

A
  • reactivity decreases down group 7
  • atoms become larger
  • outer electrons are further from the nucleus
  • shielded more from the attraction of nucleus
  • harder for larger atoms to attract the electron to form an ion
35
Q

Why are halogens oxidising agents?

A
  • they usually react by gaining an electron, so they are reduced
  • they oxidise other substances
36
Q

Describe the solubility in water of halogens

A
  • halogens are non-polar diatomic molecules
  • they have low solubility in water
37
Q

What do halogens dissolve in?

A
  • they dissolve easily in organic compounds (such as cyclohexane)
38
Q

What is the colour of chlorine, bromine and iodine in hexane?

A
  • chlorine: virtually colourless
  • bromine: yellow
  • iodine: purple
39
Q

What is the colour in water of chlorine, bromine and iodine?

A
  • chlorine: virtually colourless (or very pale green)
  • bromine: yellow/orange
  • iodine: brown
40
Q

Which halogen will displace both bromide and iodide ions?

A
  • chlorine
41
Q

What is a displacement reactions between halogens and halides called?

A
  • a redox reaction
  • what is displaced is oxidised, what is displacing is reduced
42
Q

Write down half equations for the displacement reaction between chlorine and potassium bromide

A
  • Cl2 + 2e- → 2Cl-
  • 2Br- → Br2 + 2e-
43
Q

Write down the ionic equation when chlorine water reacts with potassium bromide

A
  • Cl2 + 2Br- → 2Cl- + Br2
44
Q

What is the colour change when bromide is displaced and Br2 is formed?

A
  • colourless to orange mixture formed
45
Q

What is the colour change when iodide is displaced and iodine is formed?

A
  • colourless to brown
46
Q

What would shaking a halogen/halide displacement reaction mixture with an organic solvent help to do?

A
  • the halogen present will dissolve in the organic solvent
  • easier to see the colour changes
47
Q

Describe how halogens react with group 1 and group 2 metals

A
  • forms halide salts
  • most vigorous reactions are elements at the bottom of group 1 and 2 with elements at the top of group 7
  • redox reaction
48
Q

What is a disproportionation reaction?

A
  • an element is simultaneously oxidised and reduced
49
Q

Describe the reaction between chlorine and water and its uses

A
  • Cl2(g) + H2O(l) → HClO(aq) + HCl(aq)
  • disproportionation reaction
  • product is called chloric(I) acid
  • chlorine is used in water treatment to kill bacteria
50
Q

What is the reaction between chlorine (and other halogens) with cold, dilute aqueous sodium hydroxide?

A
  • 2NaOH(aq) + Cl2(g) → NaClO(aq) + NaCl(aq) + H2O(l)
  • disproportionation reaction
  • sodium chlorate is bleach
  • colour of halogen solution fades to colourless
51
Q

What is sodium chlorate used for?

A
  • water treatment
  • bleach paper
  • textiles
52
Q

What is the disproportionation reactions of chlorine (and other halogens) with hot alkalis?

A
  • 3Cl2 + 6NaOH → 5NaCl(aq) + NaClO3(aq) + 3H2O(l)
53
Q

Describe the trend in reducing power of halide ions

A
  • reducing power increases down the group
54
Q

Describe the reaction between NaF and NaCl with H2SO4

A
  • sulfuric acid is not strong enough an oxidising agent to oxidise the chloride and fluoride ins
  • no redox reactions occur, only acid-base reactions
  • white steamy fumes observed
55
Q

Describe the reaction between NaBr and H2SO4

A
  • Br- ions are stronger reducing agents than Cl- and F- and after the initial acid-base reaction reduce the sulfur in H2SO4 from +6 to +4 in SO2
  • white steamy fumes of HBr
  • red fumes of bromine and a colourless acidic gas of SO2
56
Q

Describe the reaction between NaI and H2SO4

A
  • I- ions are the strongest halide reducing agents
  • they reduce sulfur from +6 in H2SO4 to +4 in SO2 to 0 in S to -2 in H2S

Observations:

  • white steamy fumes of HI
  • black solid and purple fumes of iodine
  • colourless, acidic gas of SO2
  • yellow solid of sulphur
  • H2S gas, like rotten eggs
57
Q

What colour are hydrogen halides?

A
  • colourless gases
58
Q

Do hydrogen halides dissolve in water?

A
  • yes
  • they are polar diatomic molecules
59
Q

Describe the reactions of hydrogen halides with water

A
  • they readily react with water to form acidic solutions, which are all colourless
60
Q

What do hydrogen halides react with ammonia gas to give?

A
  • forms salts, all of which are white ionic solids
  • e.g. ammonia and chlorine gas give ammonium chloride
  • white fumes
61
Q

Describe silver nitrate to test for halides

A
  • add dilute nitric acid to remove any carbonates present that may react with silver ions
  • add silver nitrate
  • a precipitate of silver halide is formed
62
Q

What colour precipitate is formed when reacting silver nitrate with chloride?

A
  • white precipitate
63
Q

What colour precipitate is formed when reacting silver nitrate with bromide?

A
  • cream precipitate
64
Q

What colour precipitate is formed when reacting silver nitrate with iodide

A
  • yellow precipitate
65
Q

Why would you add ammonia solution to silver halide?

A
  • AgCl: precipitate dissolves in dilute ammonia solution to give a colourless solution
  • AgBr: Precipitate remains unchanged if dilute ammonia solution is added, but will dissolve in concentrated ammonia solution to give colourless solution
  • AgI: precipitate does not dissolve even in conc ammonia
66
Q

What does dilute HCl test for?

A
  • carbonates and hydrogencarbonates
  • they give off CO2 (and H2O)
  • test CO2 with limewater
67
Q

How do you test for sulfate?

A
  • add dilute HCl to get rid of traces of carbonate, which may produce precipitate
  • add barium chloride solution
  • if a white precipitate of barium sulfate forms, there are sulfates
68
Q

How do you test for ammonium compounds?

A
  • to test whether substance has NH4+, add sodium hydroxide and gently heat
  • if ammonia is given off, then there are ammonium ions
  • test for ammonia using damp piece of red litmus paper, which will turn blue

A Level Chemistry (19 decks)

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