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AS Chemistry > RX + PT > Flashcards

RX + PT Flashcards

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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
What is the trend in ionisation energy going down groups 1 and 2?
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
26
What is the reactivity trend going down groups 1 and 2?
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
27
Reaction between group 2 metal with water
M(s) + 2H20(l) —> M(OH)2 (aq) + H2(g) Group 2 metal + water —> Metal hydroxide + hydrogen
28
Reaction between group 2 metal and oxygen
2M(s) + O2(g) —> 2MO(s) Group 2 metal + oxygen —> metal oxide
29
Reaction of group 2 metals (Sr and Ba) with excess oxygen
M(s) + O2(g) —> MO2
30
Reactions between group 2 metal and chlorine
M(s) + Cl2(g) —> MCl2 (s) Group 2 metal + chlorine —> metal chloride
31
Reaction between group 1 metal and oxygen
4M + O2 —> 2M2O Group 1 metal + oxygen —> metal oxide
32
Reaction between group 1 metal and excess oxygen
2M + O2 —> M2O2 Group 1 metal + excess oxygen —> metal peroxide
33
Reactions of group 1 metals and water
Group 1 metal + water —> metal hydroxide + H2
34
Reactions of group 1 metals and chlorine
Group 1 metal + chlorine —> metal chloride
35
Reaction between group 1/2 oxide and water
Group 1/2 oxide + water —> metal hydroxide Group 1: M2O + H2O —> 2MOH Group 2: MO + H2O —> M(OH)2
36
Reaction between group 1/2 oxides and dilute acid
Group 1/2 oxide + hydrochloric acid —> metal chloride + water Group 1: M2O + 2HCl —> 2MCl + H2O Group 2: MO + 2HCl —> MCl2 + H2O
37
Reactions between group1/2 hydroxides and dilute acid
Group 1/2 hydroxide + hydrochloric acid —> metal chloride + water Group 1: MOH + HCl —> MCl + H20 Group 2: M(OH)2 + 2HCl —> MCl2 + 2H2O
38
Solubility trend of Group 2 hydroxides
Increases down the group
39
Solubility trend of Group 2 sulfates
Decrease down the group
40
Thermal stability of Group 1 and 2 compounds (carbonates and nitrates)
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
Why are group 2 compounds less thermally stable then group 1 compounds?
1. Group 2 cations have a higher charge density 2. More polarising 3. Less stable
42
Thermal decomposition of Group 1 carbonates
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
Thermal decomposition of group 1 nitrates
1. Group 1 nitrates decompose to form nitrite and oxygen 2. 2MNO3 —> 2MNO2 + O2 3. Only LiNO3 decomposes to form Li2O + NO2 + O2
44
Thermal decomposition of group 2 carbonates
All group 2 carbonates decompose to form oxide and carbon dioxide MCO3 —> MO + CO2
45
Thermal decomposition of group 2 nitrates
All nitrates decompose to form oxide, nitrogen dioxide and oxygen 2M(NO3)2 —> 2MO + 4NO2 + O2
46
Test for nitrogen dioxide
1. Brown gas 2. Acidic pH when dissolved in water
47
Test for carbon dioxide
Bubble the gas through limewater, if CO2 is present limewater turns cloudy
48
Test for oxygen
A glowing splint relights in the presence of oxygen
49
Suggest how characteristic flame colours are formed during a flame test
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
Describe the method for a flame test
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
Lithium colour in flame test
Red
52
Sodium colour in flame test
Orange/yellow
53
Potassium colour in flame test
Lilac
54
Rubidium colour in flame test
Red
55
Caesium colour in flame test
Blue
56
Calcium colour in flame test
Brick red
57
Strontium colour in flame test
Crimson
58
Barium colour in flame test
Green
59
Test for carbonate ions
1. CO2 is produced 2. Bubble the gas through limewater 3. Limewater turns cloudy in the presence of carbon dioxide
60
Test for hydrogencarbonate ions
1. CO2 is produced 2. Bubble the gas through limewater 3. If carbon dioxide is present, limewater turns cloudy
61
Test for sulfate ions
1. Add acidified barium chloride 2. A white precipitate of barium sulfate forms
62
Test for hydroxide ions
Hydroxide ions turn damp red litmus paper blue
63
Test for ammonium ions
1. Add sodium hydroxide 2. Gently warm 3. Damp red litmus paper turns blue
64
Colour and state of fluorine at room temperature
Pale yellow, gas
65
Colour and state of chlorine at room temperature
Pale green, gas
66
Colour and state of bromine at room temperature
Red/brown, liquid
67
Colour and state of iodine at room temperature
Grey/black, solid
68
Colour of chlorine in aqueous solution
Pale yellow
69
Colour of bromine in aqueous solution
Orange/red
70
Colour of iodine in aqueous solution
Brown
71
Colour of chlorine in organic solvent
Pale yellow
72
Colour of bromine in organic solution
Red
73
Colour of iodine in organic solvent
Purple
74
Trend in melting point of halogens
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
Trend in electronegativity of halogens
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
Trend in reactivity of halogens
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
Group 1 oxide + halogen, change in oxidation number
Group 1 element is oxidised: 0 —> +1 Halogen is reduced: 0 —> -1
78
Group 2 oxide + halogen, change in oxidation number
Group 2 element is oxidised: 0 —> +2 Halogen is reduced: 0 —> -2
79
Disproportionation reaction of chlorine with water equation
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
What is chlorine used for?
Water treatment
81
Disproportionation of chlorine with cold alkalis equation
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
Disproportionation of chlorine with hot alkalis
3Cl2 + 6NaOH —> NaClO3 + 5NaCl + 3H2O Oxidation of chlorine: 0 —> +5 Reduction of chlorine: 0 —> -1
83
Trend in reducing power of halides
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
Reaction of potassium fluoride and sulfuric acid
KF(s) + H2SO4(l) —> KHSO4(s) + HF(g) Steamy fumes produced NOT REDOX
85
Reaction of potassium chloride and sulfuric acid
KCl(s) + H2SO4(l) —> KHSO4(s) + HCl(g) Steamy fumes produced NO REDOX
86
Reaction of potassium bromide with sulfuric acid
KBr(s)+H2SO4(l) —> KHSO4(s) +HBr 2HBr(aq) + H2SO4(l) —>Br2(g) + SO2(g) + 2H2O(l)
87
Change in oxidation state of sulfur and bromine
Oxidation state of S: +6 —> +4= reduction Oxidation state of Br: -1 —> 0 = oxidation
88
Reaction of potassium iodide and sulfuric acid
KI + H2SO4 —> KHSO4 + HI 2HI + H2SO4 —> I2 + SO2 + 2H2O 6HI + SO2 —> 2H2S + 3I2 + 2H2O

AS Chemistry (5 decks)

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