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AS Chemistry > RX + PT > Flashcards

RX + PT Flashcards

1
Q

Define oxidation (3)

A
  1. Addition of oxygen
  2. Loss of hydrogen
  3. Loss of electrons
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2
Q

Define reduction (3)

A
  1. Loss of oxygen
  2. Addition of hydrogen
  3. Gain of electrons
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3
Q

Define oxidation number

A

The oxidation state/charge of an element or ionic substance

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4
Q

Oxidation number of uncombined elements

A

0

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5
Q

Oxidation number of a monoatomic ion

A

The charge of the ion

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6
Q

Oxidation number of molecular ions

A

Sum of oxidation numbers is equal to overall charge

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7
Q

Oxidation number of a neutral compound

A

Overall charge is 0, sum of oxidation numbers is 0

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8
Q

Oxidation number of hydrogen

A

+1

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9
Q

Oxidation number of oxygen

A

-2

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10
Q

Oxidation number of halogens

A

-1

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11
Q

Oxidation number of group 1 metals

A

+1

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12
Q

Oxidation number of group 2 metals

A

+2

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13
Q

Oxidation number of oxygen in peroxides

A

-1

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14
Q

Oxidation number of hydrogen in metal hydrides (MHx)

A

-1

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15
Q

When are Roman numerals used?

A

For transition metals with variable oxidation numbers

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16
Q

Oxidation state of S block elements (Group 1 and 2)

A
  1. Lose electrons
  2. Oxidised
  3. Positive oxidation numbers
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17
Q

Oxidation state of P block elements

A
  1. Gain electrons
  2. Reduced
  3. Negative oxidation number
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18
Q

What is an oxidising agent?

A
  1. Gains electrons
  2. Reduced
  3. Increase in oxidation number
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19
Q

What is a reducing agent?

A
  1. Donates/loses electrons
  2. Oxidised
  3. Decrease in oxidation number
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20
Q

Define redox reaction

A

A reaction where reduction and oxidation occurs simultaneously

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21
Q

Define disproportionation reaction

A

A reaction in which an element in a single species is simultaneously oxidised and reduced

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22
Q

Trend in oxidation of metals

A
  1. Lose electrons
  2. Increase in oxidation number
  3. Oxidised
  4. Positive ions formed
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23
Q

Trend in oxidation of non metals

A
  1. Gain electrons
  2. Decrease in oxidation number
  3. Negative ions formed
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24
Q

Steps for ionic half equations and full ionic equations

A
  1. Write half equation for first element
  2. Balance all species excluding oxygen and hydrogen
  3. Balance oxygen by adding H20
  4. Balance hydrogen by adding H+ ions
  5. Balance charges using e- electrons
  6. Write half equation for second element
  7. Balance
  8. Ensure both equations have the same number of electrons
  9. Combine
  10. Cross off any common substances
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25
Q

What is the trend in ionisation energy going down groups 1 and 2?

A

Decrease in ionisation energy
1. More electrons going down the group
2. More quantum shells
3. Increased atomic radius
4. Distance between outermost electron and nucleus is greater
5. Weaker electrostatic attraction between nucleus and outermost electron
6. Less energy is required for the outermost electron to be removed

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26
Q

What is the reactivity trend going down groups 1 and 2?

A

Increase in reactivity
1. More electrons
2. More quantum shells
3. Larger atomic radius
4. Greater distance between outermost electron and nucleus
5. Weaker electrostatic attraction between outermost electron and nucleus
6. Less energy is required to remove the outermost electron

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27
Q

Reaction between group 2 metal with water

A

M(s) + 2H20(l) —> M(OH)2 (aq) + H2(g)
Group 2 metal + water —> Metal hydroxide + hydrogen

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28
Q

Reaction between group 2 metal and oxygen

A

2M(s) + O2(g) —> 2MO(s)
Group 2 metal + oxygen —> metal oxide

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29
Q

Reaction of group 2 metals (Sr and Ba) with excess oxygen

A

M(s) + O2(g) —> MO2

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30
Q

Reactions between group 2 metal and chlorine

A

M(s) + Cl2(g) —> MCl2 (s)
Group 2 metal + chlorine —> metal chloride

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31
Q

Reaction between group 1 metal and oxygen

A

4M + O2 —> 2M2O
Group 1 metal + oxygen —> metal oxide

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32
Q

Reaction between group 1 metal and excess oxygen

A

2M + O2 —> M2O2
Group 1 metal + excess oxygen —> metal peroxide

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33
Q

Reactions of group 1 metals and water

A

Group 1 metal + water —> metal hydroxide + H2

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34
Q

Reactions of group 1 metals and chlorine

A

Group 1 metal + chlorine —> metal chloride

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35
Q

Reaction between group 1/2 oxide and water

A

Group 1/2 oxide + water —> metal hydroxide
Group 1: M2O + H2O —> 2MOH
Group 2: MO + H2O —> M(OH)2

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36
Q

Reaction between group 1/2 oxides and dilute acid

A

Group 1/2 oxide + hydrochloric acid —> metal chloride + water
Group 1: M2O + 2HCl —> 2MCl + H2O
Group 2: MO + 2HCl —> MCl2 + H2O

37
Q

Reactions between group1/2 hydroxides and dilute acid

A

Group 1/2 hydroxide + hydrochloric acid —> metal chloride + water
Group 1: MOH + HCl —> MCl + H20
Group 2: M(OH)2 + 2HCl —> MCl2 + 2H2O

38
Q

Solubility trend of Group 2 hydroxides

A

Increases down the group

39
Q

Solubility trend of Group 2 sulfates

A

Decrease down the group

40
Q

Thermal stability of Group 1 and 2 compounds (carbonates and nitrates)

A
  1. Increases down the group
  2. Atomic radius increases
  3. Cations are bigger and have a smaller charge
  4. Lower charge density
  5. Lower polarising power
  6. Less distortion of carbonate/nitrate ions
  7. More stable compounds formed
41
Q

Why are group 2 compounds less thermally stable then group 1 compounds?

A
  1. Group 2 cations have a higher charge density
  2. More polarising
  3. Less stable
42
Q

Thermal decomposition of Group 1 carbonates

A
  1. All group 1 carbonates are thermally stable
  2. Decompose at higher temperatures
  3. Only lithium carbonate decompose
  4. Li2CO3(s) —> Li2O(s) + CO2(g)
43
Q

Thermal decomposition of group 1 nitrates

A
  1. Group 1 nitrates decompose to form nitrite and oxygen
  2. 2MNO3 —> 2MNO2 + O2
  3. Only LiNO3 decomposes to form Li2O + NO2 + O2
44
Q

Thermal decomposition of group 2 carbonates

A

All group 2 carbonates decompose to form oxide and carbon dioxide
MCO3 —> MO + CO2

45
Q

Thermal decomposition of group 2 nitrates

A

All nitrates decompose to form oxide, nitrogen dioxide and oxygen
2M(NO3)2 —> 2MO + 4NO2 + O2

46
Q

Test for nitrogen dioxide

A
  1. Brown gas
  2. Acidic pH when dissolved in water
47
Q

Test for carbon dioxide

A

Bubble the gas through limewater, if CO2 is present limewater turns cloudy

48
Q

Test for oxygen

A

A glowing splint relights in the presence of oxygen

49
Q

Suggest how characteristic flame colours are formed during a flame test

A
  1. Energy absorbed from the flame causes electrons to become “excited”, causing them to move to a higher energy level
  2. Electron is unstable in this energy level
  3. Drops back down to original energy level
  4. Energy is emitted in the form of visible light wavelength which allows for a colour to be observed
50
Q

Describe the method for a flame test

A
  1. Take a nichrome wire and dip it into HCl
  2. Heat the wire in the flame
  3. Take the clean ncihrome wire and dip it into HCl
  4. Dip into sample
  5. Place wire with sample in flame and observe colour change
51
Q

Lithium colour in flame test

A

Red

52
Q

Sodium colour in flame test

A

Orange/yellow

53
Q

Potassium colour in flame test

A

Lilac

54
Q

Rubidium colour in flame test

A

Red

55
Q

Caesium colour in flame test

A

Blue

56
Q

Calcium colour in flame test

A

Brick red

57
Q

Strontium colour in flame test

A

Crimson

58
Q

Barium colour in flame test

A

Green

59
Q

Test for carbonate ions

A
  1. CO2 is produced
  2. Bubble the gas through limewater
  3. Limewater turns cloudy in the presence of carbon dioxide
60
Q

Test for hydrogencarbonate ions

A
  1. CO2 is produced
  2. Bubble the gas through limewater
  3. If carbon dioxide is present, limewater turns cloudy
61
Q

Test for sulfate ions

A
  1. Add acidified barium chloride
  2. A white precipitate of barium sulfate forms
62
Q

Test for hydroxide ions

A

Hydroxide ions turn damp red litmus paper blue

63
Q

Test for ammonium ions

A
  1. Add sodium hydroxide
  2. Gently warm
  3. Damp red litmus paper turns blue
64
Q

Colour and state of fluorine at room temperature

A

Pale yellow, gas

65
Q

Colour and state of chlorine at room temperature

A

Pale green, gas

66
Q

Colour and state of bromine at room temperature

A

Red/brown, liquid

67
Q

Colour and state of iodine at room temperature

A

Grey/black, solid

68
Q

Colour of chlorine in aqueous solution

A

Pale yellow

69
Q

Colour of bromine in aqueous solution

A

Orange/red

70
Q

Colour of iodine in aqueous solution

A

Brown

71
Q

Colour of chlorine in organic solvent

A

Pale yellow

72
Q

Colour of bromine in organic solution

A

Red

73
Q

Colour of iodine in organic solvent

A

Purple

74
Q

Trend in melting point of halogens

A

Down the group
1. Increase in quantum shells/shielding
2. Increase in number of electrons
3. Stronger London dispersion forces
4. More energy is needed to overcome strong intermolecular forces
5. Melting point and boiling point increases

75
Q

Trend in electronegativity of halogens

A

Down the group
1. Atomic radius increases
2. Outer electrons are further away from the nucleus
3. Incoming electron experiences more shielding from the attraction of the positive nuclear charge (nucleus)
4. More difficult to attract an electron
5. Electronegativity decreases

76
Q

Trend in reactivity of halogens

A

Down the group
1. Atomic radius increases
2. Outer electrons are further away from the nucleus
3. Increased shielding
4. Harder to attract an electron
5. Decreased reactivity

77
Q

Group 1 oxide + halogen, change in oxidation number

A

Group 1 element is oxidised: 0 —> +1
Halogen is reduced: 0 —> -1

78
Q

Group 2 oxide + halogen, change in oxidation number

A

Group 2 element is oxidised: 0 —> +2
Halogen is reduced: 0 —> -2

79
Q

Disproportionation reaction of chlorine with water equation

A

Cl2(g) + H2O(l) —> HCl(aq) + HClO(aq)
Chlorine + water —> hydrochloric acid and hydrochlorous acid

HClO(aq) +H2O(l) —> ClO- (aq) + H3O+(aq)
Hydrochlorous acid + water —> chlorite ions

80
Q

What is chlorine used for?

A

Water treatment

81
Q

Disproportionation of chlorine with cold alkalis equation

A

2NaOH(aq) + Cl2(g) —> NaClO(aq) + NaCl(aq) + H2O(l)
Sodium hydroxide + chlorine —> Sodium chlorate + sodium chloride + water
Oxidation of chlorine: 0 —> +1 in NaClO (sodium chlorate/bleach)
Reduction of chlorine: 0 —> -1 in NaCl (sodium chloride)

82
Q

Disproportionation of chlorine with hot alkalis

A

3Cl2 + 6NaOH —> NaClO3 + 5NaCl + 3H2O
Oxidation of chlorine: 0 —> +5
Reduction of chlorine: 0 —> -1

83
Q

Trend in reducing power of halides

A

Down the group
1. Ionic radius increases
2. Electrons are further away from the nucleus
3. Greater shielding effect
4. Easier to remove and electron

84
Q

Reaction of potassium fluoride and sulfuric acid

A

KF(s) + H2SO4(l) —> KHSO4(s) + HF(g)
Steamy fumes produced
NOT REDOX

85
Q

Reaction of potassium chloride and sulfuric acid

A

KCl(s) + H2SO4(l) —> KHSO4(s) + HCl(g)
Steamy fumes produced
NO REDOX

86
Q

Reaction of potassium bromide with sulfuric acid

A

KBr(s)+H2SO4(l) —> KHSO4(s) +HBr
2HBr(aq) + H2SO4(l) —>Br2(g) + SO2(g) + 2H2O(l)

87
Q

Change in oxidation state of sulfur and bromine

A

Oxidation state of S: +6 —> +4= reduction
Oxidation state of Br: -1 —> 0 = oxidation

88
Q

Reaction of potassium iodide and sulfuric acid

A

KI + H2SO4 —> KHSO4 + HI
2HI + H2SO4 —> I2 + SO2 + 2H2O
6HI + SO2 —> 2H2S + 3I2 + 2H2O

AS Chemistry (5 decks)

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