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A-Level Chemistry [OCR B Salters] (sorry for any spelling mistakes - there are many) > [Year 1] ES > Flashcards

[Year 1] ES Flashcards

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1
Q

How is a reversible reaction shown on a concentration-time graph?

A

Reactants: negative exponential.

Products: Positive exponential.

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2
Q

What happens when a concentration-time graph showing a reversible reaction plateaus?

A

It reaches a dynamic equilibrium.

only in a closed system

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3
Q

What is a dynamic equilibrium?

A

The point at which the rate of the forwards reaction is equal to the rate of the backwards reaction.

The concentration of each remains constant.

(amount NOT same).

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4
Q

How is the equilibrium constant found?

A

nA + nB ⇌ nC + nD

K꜀ = [C]ⁿ[D]ⁿ / [A]ⁿ[B]ⁿ

[Products] / [Reactants]

[x] = concentration of x (in mol dm⁻³)

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5
Q

How does temperature affect the equilibrium constant in a reversible reaction?

A

If temperature change shift equilibrium to the right, K꜀ will increase.

If temperature change shift equilibrium to the left, K꜀ will decrease.

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6
Q

What happens to the equilibrium if you heat an exothermic reversible reaction?

A

Equilibrium shifts in the endothermic direction (to oppose the increase in temperature).

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7
Q

What happens to the equilibrium if you cool an exothermic reversible reaction?

A

Equilibrium shifts in the exothermic direction (to oppose the decrease in temperature).

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8
Q

What does the equilibrium constant tell us about the position of equilibrium?

A

If K꜀&raquo_space; 1,
The equilibrium lies well over to the right (many more products)

If K꜀ > 1,
The equilibrium lies slightly to the right (more products)

If K꜀ = 1,
The equilibrium lies in the middle (same)

If K꜀ < 1,
The equilibrium lies slightly to the left( more reactants)

If K꜀ &laquo_space;1,
The equilibrium lies well over to the left (many more reactants)

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9
Q

How does concentration affect the equilibrium constant in a reversible reaction?

A

No effect on K꜀

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10
Q

What happens to the equilibrium if you increase the pressure of a reversible reaction with more moles of gas in the reactants?

A

Equilibrium shifts in the forward direction (to oppose the increase in pressure).

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11
Q

What happens to the equilibrium if you decrease the pressure of a reversible reaction with more moles of gas in the reactants?

A

Equilibrium shifts in the backwards direction (to oppose the decrease in pressure).

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12
Q 
Summarise the shift in equilibria when we:
increase temperature.
decrease temperature.
increase pressure.
decrease pressure.
A

increase temperature = shift towards endothermic

decrease temperature = shift toward exothermic

increase pressure = shift towards fewer moles

decrease pressure = shift towards more moles

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13
Q

What is electrolysis?

A

The breaking down of a substance using electricity.

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14
Q

Why is electrolysis done using an ionic compound that is either molten or dissolved in solution?

A

Because, there are free moving ions which allow the conduction of electricity.

We also need the flow of electrons.

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15
Q

What are the two electrodes in electrolysis?

A

Anode:
a positive electrode that attracts anions to give up electrons.

Cathode:
a negative electrode that attracts cations to attract electrons.

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16
Q

How would you conduct electrolysis?

A
  1. Connect the power supply to inert + conductive electrodes. (pos end will create anode and the negative end will create cathode)
  2. Dipp electrodes in the electrolyte but they mustn’t touch.
  3. Turn on the power supply.
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17
Q

Why must the electrode in electrolysis be inert and conductive?

A

Conductive: so they can become charged.

Inert: so they don’t react with electrolytes.

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18
Q

What will you see if a solid, liquid and gas is produced at an electrode?

A

Solid: plating of the electrode.

Liquid: colour change (with indicator used) (normally acidic and alkaline products)

Gas: effervescence around the electrode.

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19
Q

What does a half equation show?

A

The movement of electrons.

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20
Q

What can produce at the cathode when the electrolytes are in solution?

A
  • Metal (if it’s less reactive than H)
    Cu²⁺(aq) + 2e⁻ → Cu(s)
  • Hydrogen gas (if metal is more reactive)
    2H₂O(l) + 2e⁻ → 2OH⁻(aq) + H₂(g)
  • Hydrogen gas (if electrolysis of acids)
    2H⁺(aq) + 2e⁻ → H₂(g)
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21
Q

What can produce at the anode when the electrolytes are in solution?

A
  • Halogen (if salt is halide)
    2Cl⁻(aq) → Cl₂(g) + 2e⁻
  • Oxygen (if salt is sulfate or nitrate)
    2H₂O(l) → O₂(g) + 4H⁺(aq) + 4e⁻
  • Oxygen (if hydroxides)
    4OH⁻(aq) → O₂(g) + 2H₂O(l) + 4e⁻
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22
Q

How is copper purified?

A

Electrolysis with a pure copper cathode and impure copper anode.

Copper at anode loses electrons to form copper ions:
Cu(s) → Cu²⁺(aq) + 2e⁻
Thus anode becomes lighter (as it wears away).

Copper ions dissolve in solution.

Copper is unreactive and so is attracted at the cathode and gains electrons:
Cu²⁺(aq) + 2e⁻ → Cu(s)
This plates the cathode (which becomes heavier).

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23
Q

How is chlorine extracted from brine?

A

Electrolysis of brine (concentrated NaCl solution) with inert, conductive electrodes.

  • Hydrogen effervesces at the cathode
  • Chlorine effervesces at anode.
  • Sodium ions are left behind in the brine with react with hydroxide ions to form NaOH
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24
Q

Why must the brine solution be concentrated?

A
  • The brine ( NaCl(aq) ) must be concentrated for halide ions to form at the anode.
  • Otherwise, chlorine won’t release its electrons, and Oh ions will making oxygen and water instead.
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25
Q

How is bromine extracted from brine?

A
  • By using a more reactive halide to displace it from NaBr(aq):
    2Br⁻(aq) + Cl₂(g) → Br₂(g) + 2Cl⁻(aq)
  • This is then condensed and purified into a liquid.
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26
Q

How is Iodine extracted from brine?

A
  • By using a more reactive halide to displace it from NaI(aq):
    2I⁻(aq) + Cl₂(g) → I₂(aq) + 2Cl⁻(aq)
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27
Q

When are electrons transfered?

A

When reduction and oxidation occurs.

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28
Q

What is reduction?

A

The gain of electrons.

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29
Q

What is oxidation?

A

The loss of electrons.

30
Q

What is a reducing agent?

A

Something that donates electrons.

31
Q

What is an oxidising agent?

A

Something that attracts electrons.

32
Q

What are the rules for oxidation states?

A
  • Elements have an oxidation state of 0.
  • Ions have an oxidation state equal to its charge.
  • Group 1 are always +1
  • Group 2 are always +2
  • Aluminium is always +3
  • Hydrogen is +1 (unless in hydrides where it is -1)
  • Chlorine is -1 (except with F and O)
  • Fluorine is always -1
  • oxygen is always -2 (except in peroxides where it’s -1 and in OF₂ where is +2)
33
Q

How is Iron(III) Oxide different from Iron(II) Oxide?

A

It oxidation states.

Iron(III) has an iron with an oxidation state of +3
Fe₂O₃

Iron(II) has an iron with an oxidation state of +2
FeO

34
Q

What is the systematic name for ClO₂⁻ ?

A

Chlorate(III)

35
Q

What are the formulas for these ions:

Carbonate
Hydroxide
Sulfate / Sulfate(VI)
Ammonium
Hydrogencarbonate
Manganate(VII)
Nitrate / Nitrate(V)
Sulfide
A

Carbonate = CO₃²⁻

Hydroxide = OH⁻

Sulfate / Sulfate(VI) = SO₄²⁻

Ammonium = NH₄⁺

Hydrogencarbonate = HCO₃⁻

Manganate(VII) = MnO₄⁻

Nitrate / Nitrate(V) = NO₃⁻

Sulfide = S²⁻

36
Q

What is reduction and oxidation in terms of oxidation numbers?

A

Reduction = decrease in oxidation number.

Oxidation = increase in oxidation number.

37
Q

Show a full ionic equation for the following REDOX reaction.

Cu → Cu²⁺ + 2e⁻
O₂ + 4e⁻ → 2O²⁻

A
  • Balance electrons then cancel from both sides.
    Cu → Cu²⁺ + 2e⁻
    0.5O₂ + 2e⁻ → O²⁻
  • Combine.
    Cu + 0.5O₂ → Cu²⁺ + O²⁻
38
Q

How would you work out the following question:

Balance the following reaction: Fe³⁺ + I⁻ → Fe + I₂

A

Fe³⁺ + I⁻ → Fe + I₂

  1. Balence the atoms.
    Fe³⁺ + 2I⁻ → Fe + I₂
  2. Check the charges on both sides.
    Reactants = 3 +(2x-1) = +1
    Products = 0
  3. Work out the change in oxidation states for each molecule/element.
    Fe: +3 to 0 = -3
    I₂: 2x-1 =-2 to 0 = +2
  4. Multiply each to cancel each other out.
    -3 x 2 = -6
    +2 x 3 = 6
  5. Multipy these by the elements/molecules in the equation.
    2Fe³⁺ + 6I⁻ → 2Fe + 3I₂
39
Q

How can you find out the concentration of an oxidising agent?

A

Iodine-Sodium Thiosulfate Titration

40
Q

How would you undergo the experiment to find the concentration of an oxidising agent?

A
  1. Use the oxidising agent (KIO₃) to oxidise iodide ions to iodine.
    - Measure 25cm³ of Potassium iodate(V) (KIO₃)
    - This produced IO₃⁻ ions.
    - Add excess Potassium iodide solution (KI)
    IO₃⁻(aq) + 5I⁻(aq) + 6H⁺(aq) → 3H₂O(l) + 3I₂(aq)
    - The more concentrated the IO₃⁻ the more I⁻ will oxidise from KI to I₂.
    (So we are measuring the strength of IO₃⁻)
  2. Carry out a titration to work out moles of iodine produced in step 1.
    - Add solution from step 1 into a conical flask.
    - Titrate in sodium thiosulfate (Na₂S₂O₃) until pale yellow.
    - (colour change is difficult to see) Just as it turns yellow, add 2cm³ of starch. This will turn blue if iodine is still present.
    - Keep titrating until blue has disappeared.
    - Note volume of sodium thiosulfate added.
    - Use volume to work out moles of sodium thiosulfate added via Moles = Conc x Vol (dm³).
    - Write out reaction between iodine and thiosulfate ions:
    I₂(aq) + 2S₂O₃²⁻(aq) → S₄O₆²⁻(aq) + 2I⁻(aq)
    - Use ratios to work out moles of I₂ with moles of S₂O₃²⁻ ions.
  3. Use the moles of iodine in step 2 to work out the concentration of IO₃⁻.
    - Write equation from step 1:
    IO₃⁻(aq) + 5I⁻(aq) + 6H⁺(aq) → 3H₂O(l) + 3I₂(aq)
    - Use rations to work out the moles of IO₃⁻ ions with moles of I₂.
    - Use moles of IO₃⁻ to work out concentration via Moles = Conc x Vol (being the 25cm³ from step 1 in dm³)
41
Q

When undergoing the experiment to find the concentration of an oxidising agent, what practical techniques can be used to ensure reliable results?

A
  • Rinse burette out with sodium thiosulfate before titrating. To get rid of trace amounts of water (that will dilute the titre).
  • Read from the bottom of the meniscus at eye level.
  • Repeat titration until you get 3 concordant results and find average.
  • Wash conical flask between repeat titrations.
  • Use freshly made solutions (as they can react with oxygen).
  • Don’t add starch to early (as it will cause iodine to react with starch instead of thiosulfate solution).
42
Q

What colour and state is Fluorine at r.t.p?

A

Pale Yellow Gas

43
Q

What colour and state is Chlorine at r.t.p?

A

Pale Green Gas

44
Q

What colour and state is Bromine at r.t.p?

A

Brown-Orange Liquid

45
Q

What colour and state is Iodine at r.t.p?

A

Grey Solid

46
Q

What happens to the boiling point of halogens as you go down the group and why?

A
  • As you go down, the boiling point increases.

- Due to the id-id bonds increasing due to the increasing size and relative mass of the atom.

47
Q

What is electronegativity?

A

The ability for an atom to attract electrons towards itself in a covalent bond.

48
Q

What happens to the electronegativity of halogens as you go down the group and why?

A
  • As you go down, the electronegativity decreases.
  • Due to atoms getting larger and the distance between the positive nucleus and bonding electrons increasing.
  • There is also more shielding between the positive nucleus and bonding electrons.
49
Q

How can we observe halogen displacement reactions?

A

Adding an organic solvent (like hexane) to see a colour change.

  • Halogens present will dissolve in organic solvent layer above the aqueous layer.
50
Q

Why do displacement reactions occur?

A

Because more reactive halogens will displace less reactive halogens.

51
Q

What happens to the reactivity of halogens as you go down the group and why?

A
  • As you go down, the reactivity decreases.
  • As for a reaction to occur electrons must be gained.
  • As the radius of the atom gets bigger, it gets harder to attract electrons.
  • This is because the outer electrons are further from the positive nucleus, so there is a weaker electrostatic attraction between negative outer electrons and the positive nucleus.
  • There is also more shielding, which interfered with the electrostatic attraction force making it weaker.
  • And hence less reactive and less oxidising.
52
Q

What would we see when testing the addition of chlorine water with:

KCL
KBr
KI

A

KCl: (no reaction)
Org Layer = colourless
Aq Layer = colourless

KBr: (Cl₂ + 2Br⁻ → 2Cl⁻ + Br₂)
Org Layer = Orange
Aq Layer = Yellow

KI: (Cl₂ + 2I⁻ → 2Cl⁻ + I₂)
Org Layer = Purple
Aq Layer = Brown

53
Q

What would we see when testing the addition of bromine water with:

KCl
KBr
KI

A

KCl: (no reaction)
Org Layer = Orange
Aq Layer = Yellow

KBr: (no reaction)
Org Layer = Orange
Aq Layer = Yellow

KI: (Br₂ + 2I⁻ → 2Br⁻ + I₂)
Org Layer = Purple
Aq Layer = Brown

54
Q

What would we see when testing the addition of Iodine solution with:

KCl
KBr
KI

A

KCl: (no reaction)
Org Layer = Purple
Aq Layer = Brown

KBr: (no reaction)
Org Layer = Purple
Aq Layer = Brown

KI: (no reaction)
Org Layer = Purple
Aq Layer = Brown

55
Q

How are hydrogen halides formed?

A

Adding a concentrated acid to an ionic halide compound.

e.g phosphoric acid

56
Q

How can we prove halide ions as reducing agents?

A
  • Reaction with sulfuric acid.

- Reaction with silver nitrate solution.

57
Q

Do halide ions as reducing agents get better or worse down the group and why?

A
  • As you go down, halide ions become better reducing agents.
  • To be a reducing agent they must readily give off an electron (oxidise).
  • As you go down, the radius of the atom gets bigger, it gets harder to attract the outer electrons.
  • This is because the outer electrons are further from the positive nucleus, so there is a weaker electrostatic attraction between negative outer electrons and the positive nucleus.
  • There is also more shielding, which interfered with the electrostatic attraction force making it weaker.
  • And hence halide ions will oxidise more readily and thus become more powerful reducing agents.
58
Q

Why might we use phosphoric acid over sulfuric acid when producing hydrogen halides?

A

Sulfuric acid is an oxidising agent and thus we get sulfur-based impurities from it.

59
Q

How does the halide ion test with sulfuric acid work work?

A
  • React three sodium sulfates with sulfuric acid.
  • In the first part of the reaction each of them formes sodium sulfate and their hydrogen halide.
  • Hydrogen Chloride is observed as a white misty fume.
  • The other two go on to reduce sulfur in sulfuric acid to form sulfur dioxide.
  • Here Bromine gas is produced and you can observe it as an orange vapour.
  • After that Iodide ions from HI reduce sulfur in sulfuric acid further to form sulfur.
  • It then reduces the sulfur in sulfuric acid further to form hydrogen sulphide gas which can be recognised by a rotten egg smell.
60
Q

What happens to the stability of hydrogen halides when heated as you go down the group and why?

A
  • As you go down, the stability decreases.
  • HF and HCl are stable and won’t split when heated.
  • HBr splits partially.
  • HI splits more easily.

This is because as you go down, the atom gets larger.

This means that bonding electrons are further from the nucleus and shielded more.

This weakens the attractive force and hence weakens the bond enthalpy.

61
Q

What happens when hydrogen halides dissolve in water?

A

They form acidic solutions.

They also react with water in the air to form white misty fumes.

They dissociate when dissolved in water.

62
Q

What does dissociate mean?

A

To break apart

63
Q

What happens to hydrogen halides that react with ammonia gas?

A

We can observe ammonium halides as white fumes.

64
Q

How does the test for halide ions with silver nitrate work?

A
  • Add nitric acid to halide ion solution.
  • Then add silver nitrate solution.
  • The colour of the precipitate will help you identify the halide ion.
Chlorine = Cloudy White
Bromine = Cloudy Cream
Iodine = Cloudy Yellow
  • Dilute ammonia is added, chloride precipitate will dissolve (becoming colourless).
  • Concentrated ammonia is added to the remaining two, bromide precipitate will dissolve (becoming colourless).
  • Iodide precipitate will not disolve in ammonia.
65
Q

What are the uses of chlorine in everyday life?

A

Added to drinking water to sterilise it as it kills harmful micro-organisms which allows us to drink it safely and swim in it.

Chlorine used in bleach.

66
Q

What are the risks associated with chlorine?

A
  • Toxic and corrosive so must be kept away from skin and eyes.
  • Oxidising agent, so must be kept away from flammable materials.
  • Must be transported and stored as a liquid (as its more economical), so must be under high pressure.
67
Q

What is atom economy?

A

How efficient a reaction is.

68
Q

How can you work out the %atom economy?

A

%Atom Economy = (Mr of desired products / Mr of all products) x 100

69
Q

How would you work out the following question?

Iron oxide (Fe₂O₃) can be reduced using Carbon (coke) to make pure Iron and Carbon Dioxide. Calculate the atom economy in the extraction of Iron.

A
  1. Write out equation
    Fe₂O₃ + 1.5C → 2Fe + 1.5CO₂
  2. Work out Mr of the desired product.
    = 2 x Fe = 2 x 55.8 = 111.6
  3. Work out Mr of the products.
    = 111.6 + (1.5 x CO₂) = 111.6 + (1.5 x 44.0) = 111.6 + 66.0 = 177.6
  4. Plug into the atom economy equation.
    111.6/177.6 = 0.628
    x 100 = 62.8%
70
Q

Why is the atom economy important?

A
  • High atom economies mean that raw materials are used more efficiently. This is more sustainable.
  • High atom economies mean less waste and so benefits the environment.
  • High atom economies mean less by-product so less time and money spent separating these from the desirable product.

A-Level Chemistry [OCR B Salters] (sorry for any spelling mistakes - there are many) (14 decks)

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