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CHEMISTRY AS > * Elements from the sea > Flashcards

* Elements from the sea Flashcards

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1
Q

How does electrolysis break down a substance?

A
  • if you pass an electric current through an ionic substance thats molten or in solution
  • it breaks down into the elements its made of
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2
Q

What is the electrolyte?

A
  • the liquid used in electrolysis that conducts electricity
  • electrolytes contain free ions (ions that are free to move around)
  • the ions are usually the molten or dissolved ionic substance
    • its the free ions in the electrolyte that conducts the electricity
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3
Q

Describe where the ions move to in electrolysis?

A
  • for a complete circuit - there has to be a flow of electrons
  • negative ions (anions) move to the positive electrode (the anode) and lose electrons
  • positive ions (cations) move the the negative electrode (the cathode) and gain electrons
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4
Q

Describe how the electrolysis of an aqueous solution would be carried out

A
  • use wires and clips to connect each electrode to the power supply
  • the electrode connected to the positive pole will be the anode, and the electrode connected to the negative pole will be the cathode
  • usually inert electrodes are used (such as platinum or carbon electrodes) so that they do not react and interfere with the electrolysis
  • place the electrodes into a beaker containing the electrolyte, making sure that the electrodes do not touch each other
  • turn the power supply on
  • depending on what the electrolyte is, the products will form as metals (as a thin layer on the surface of the cathode - known as plating) or as gases (bubbles at the cathode or anode)
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5
Q

what is the purpose of half-equations in electrolysis

A
  • hald-equations show the movement of electrons during a reaction
  • in electrolysis you can write half-equations to show what’s happening at each electrode
  • the half-equation for the anode will show negative ions losing electrons to form atoms
    • e.g. 2Cl-(l) –> Cl2(g) + 2e-
  • the half-equation for the cathode will show positive ions gaining electrons to form atoms
    • e.g. Zn2+(l) + 2e- –> Zn(s)​
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6
Q

How do we predict what is formed at the electrodes from a molten compound electrolyte

A
  • if electrolyte is a molten salt
  • only ions around are the ones that make up the salt
  • so substance will just break up into its elements
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7
Q

How do we predict what is formed at the electrodes from an aqueous compound electrolyte

A
  • in aqueous solutions, you’ll have H+ and OH- ions from the water as well as the ions from the ionic compound
  • the products formed at each electrode depend on the reactivity of the ions, as well as the concentration of the salt solution
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8
Q

What are the rules for what happens at the cathode in an aqueous solution during electrolysis

A
  • Cathode
    • if metal is less reactive than hydrogen (e.g. silver or copper), then the metal will be formed
    • if the metal is more reactive than hydrogen (all group 1 and group 2 metals and aluminium)
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9
Q

What are the rules for what happens at the anode in an aqueous solution during electrolysis

A
  • Anode
    • if the solution doesn’t contain a halide, oxygen will be formed (from hydroxide ions in the water)
    • 4OH-(aq) –> O2(g) + 2H2O(l) + 4e-
    • if the solution is concentrated and contains a halide, then the halogen will be formed
    • if the solution contains a halide but is dilute, oxygen will be formed (from hydroxide ions in the water)
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10
Q

What may happen if metal electrodes are used during electrolysis

A
  • if you use metal electrodes (apart from platinum ones) metal ions can also be made at the anode
  • e.g. in the purification of copper, the anode is made from impure copper and the cathode is made from pure copper
  • at the anode, copper atoms lose electrons and become copper ions, which enter the solution
  • these ions then attracted to the cathode where they gain electrons to become copper atoms again and plate the pure copper cathode
  • Anode: Cu(s) –> Cu2+(aq) + 2e-
  • Cathode: Cu2+(aq) + 2e- –> Cu(s)
  • Pure cathode increases in mass while impure cathode shrinks
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11
Q

How can halogens be extracted by the electrolysis of halide solutions?

A
  • when concentrated aqueous solutions containing halide ions are electrolysed, the halogen element is released at the anode
  • halide ions lose electrons to the electrode and are oxidised to atoms, which combine to form molecules
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12
Q

What is brine?

A
  • brine is a solution of water with a high concentration of salts - mainly sodium chloride, but also some bromine and iodine salts
  • brine occurs naturally in salt lakes or as seawater, or can be made by dissolving rock salt in water
  • industrially, chlorine is made by the electrolysis of brine
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13
Q

Describe the electrolysis of brine to extract chlorine

A
  • at the cathode, two hydrogen ions accept two elecrons to become one hydrogen molecule
    • 2H+(aq) + 2e- –> H2(g)
  • at the anode, two chloride (Cl-) ions lose their electrons and become one chlorine molecule
    • 2Cl-(aq) –> Cl2(g) + 2e-
  • the sodium ions stay in solution because theyre more reactive than hydrogen
  • sodium ions and hydroxide ions (from water) are left behind while hydrogen and chlorine are removed, so sodium hydroxide (NaOH) is left in the solution
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14
Q

What are the conditions for the electrolysis of brine to extract chlorine

A
  • the electrodes are made of an inert material (e.g. carbon, platinum, titanium)
  • the electrolysis cell is constantly fed with a fresh stream of brine
  • the chlorine is collected as a gas
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15
Q

Why must the sodium chloride solution (or the brine) be concentrated to extract chlorine in electrolysis

A
  • can only extract chlorine from concentrated sodium chloride solution
  • in dilute solutions, the chloride ions (Cl-) aren’t discharged - they hang on to their electrons
  • the OH- ions lose electrons instead and the products at the anode are oxygen and water, not chlorine
    • 4OH-(aq) –> 2H2O(l) + O2(g) + 4e-
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16
Q

How is bromine produced using brine

A
  • displacement reaction
  • brine contains bromide ions
  • chlorine is more reactive than bromine
  • so when you bubble chlorine gas through brine, the chlorine will displace the bromine
    • 2Br-(aq) + Cl2(g) –> Br2(g) + 2Cl-(aq)
    • this is a redox reaction
  • bromine produced is then collected, condensed into a liquid and purified
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17
Q

How is iodine produced using brine

A
  • displacement reaction
  • brine contains iodide ions
  • chlorine is more reactive than iodine
  • so when you bubble chlorine gas through brine, the chlorine will displace the iodine
    • 2I-(aq) + Cl2(g) –> I2() + 2Cl-(aq)
    • this is a redox reaction
  • iodine produced is then collected, condensed into a liquid and purified
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18
Q

What does an atoms oxidation state tell you?

A
  • tells you how many electrons it has donated (lost) or accepted (gained) to form an ion or a bond
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19
Q

What is the rule for the oxidation states of uncombined atoms

A
  • oxidation state of 0
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20
Q

What is the rule for the oxidation states of identical atoms (O2)

A
  • oxidation state of 0
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21
Q

What is the rule for the oxidation states of a monatomic ion

A
  • oxidation state = same as its charge
  • Na+ - oxidation state = +1
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22
Q

What is the rule for the oxidation states of a compound ion

A
  • overall oxidation state = same as overall charge
  • sum of atoms oxidation states will equal the overall charge
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23
Q

What is the rule for the oxidation states of a neutral compound

A
  • overall oxidation state = 0
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24
Q

What is the oxidation state of fluorine?

A
  • -1
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25
Q

What is the oxidation state of oxygen?

A
  • -2
  • except when combined with F or the peroxide ion O22-
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26
Q

What is the oxidation state of chlorine?

A
  • -1
  • except when combined with fluorine or oxygen
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27
Q

What is the oxidation state of bromine

A
  • -1
  • except when combined with oxygen, fluorine or chlorine
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28
Q

What is the oxidation state of iodine

A
  • -1
  • except when combined with oxygen, fluorine or chlorine or bromine
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29
Q

What is the oxidation state of hydrogen

A
  • +1
  • except when in a metal hydride NaH
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30
Q

What is the oxidation state of all group 1 elements

A
  • +1
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31
Q

What is the oxidation state of all group 2 elements

A
  • +2
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32
Q

What is the oxidation state of aluminium

A
  • +3
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33
Q

What do roman numerals in a compound mean?

A
  • its oxidation state
  • used if an element can have multiple oxidation states
  • written after the name of the element they correspond to
34
Q

How are ions named

A
  • when naming an ionic compound
  • stick the name of the positive ion in front of the name of the negative ion
    • H2S - hydrogen sulfide
    • LiOH - lithium oxide
  • if the ion ends in -ate that tells you the ion contains oxygen with another elements
    • the non-oxygen element can sometimes have different oxidation states - so roman numerals used
35
Q

What is the systematic name for NO3-

A
  • nitrate(V)
36
Q

What is the systematic name for SO42-

A
  • sulfate(VI)
37
Q

What is the systematic name for SO42-

A
  • sulfate(VI)
38
Q

What is the systematic name for CO32-

A
  • carbonate
39
Q

What is the systematic name for MnO4-

A

manganate(VII)

40
Q

What is the systematic name for OH-

A
  • hydroxide
41
Q

What is the systematic name for NH4+

A
  • ammonium
42
Q

What is the systematic name for HCO3-

A
  • hydrogencarbonate
43
Q

What is the systematic name for S2-

A
  • sulfide
44
Q

What is a redox reaction?

A
  • oxidation is a loss of electrons
  • reduction is a gain of electrons
  • redox is when oxidation and reduction happen simultaneously
45
Q

How are oxidation states effected by lost or gained electrons

A
  • oxidation state of an atom will increase by 1 for each electron lost
  • oxidation state of an atom will decrease by 1 for each electron gained
46
Q

How are half-equations relevant to redox equations

A
  • half-equations show oxidation or reduction
  • you can combine half equations to make full equations for redox reactions
47
Q

What is the reducing agent

A
  • the thing that has been oxidised
  • reducing agent donates electrons
48
Q

What is the oxidising agent

A
  • the thing that has been reduced
  • oxidising agent accepts electrons
49
Q

How is balancing redox equation different to balancing normal equations?

A
  • charges need to be balanced as well as atoms
  • you can do this by using the oxidation states of the reactants and products
50
Q

How do you balance a redox equation?

for example: Au3+ + I- –> Au + I2

A
  • first balance atoms on each side
  • check to see if charges are balanced
    • e.g. Au3+ + 2I- –> Au + I2​
    • LHS: (+3 + (2 x -1) = 1
    • RHS: 0
  • find the change in oxidation states of both elements
    • Au = + 3 to 0
      • change = -3
    • I = -1 to 0
      • change = +1 x 2 (because there are 2 moles of iodine) = +2
  • need to find LCM - so they cancel each other out (+6 and -6)
  • multiply for Au: (-3 by 2) and for I: (+2 by 3)
  • so now - multiply all Au species by 2 and all I species by 3
  • equation = 2Au3+ + 6I- –> 2Au + 3I2
51
Q

Where are the blocks in the periodic table

A
52
Q

what is the purpose of an iodine-sodium thiosulfate titration?

A
  • way of finding the concentration of an oxidising agent
  • the more concentrated an oxidising agent is - the more ions will be oxidised by a certain volume of it
53
Q

Describe how the iodine-sodium thiosulfate titration can find out the concentration of a solution of the oxidising agent potassium iodate(V)

A
  • (KIO3)
  • Stage 1: use a sample of oxidising agent to oxidise as much iodide as possible
  • measure out a certain volume of potassium iodate(V) the oxidising agent (like 25cm3)
  • add this to an excess of acidic potassium iodide solution
    • the iodate(V) ions in the potassium iodate(V) solution will oxidise some of the iodide ions to iodine
    • IO3-(aq) + 5I-(aq) + 6H+(aq) –> 3I2(aq) + 3H2O(l)
  • Stage 2: Find out how many moles of iodine have been produced
  • do this by titrating the resulting solution with sodium thiosulfate
  • the iodine in the solution reacts with thiosulfate ions
    • I2(aq) + 2S2O32-(aq) –> 2I-(aq) + S4O62-(aq)
    • then calculate the moles of iodine
  • Stage 3: Calculate the concentration of the oxidising agent
  • IO3-(aq) + 5I-(aq) + 6H+(aq) –> 3I2(aq) + 3H2O(l)
  • equation shows 1 mole of iodate(V) produces three moles of iodine
  • so if 25.0cm3 of potassium iodate(V) solution produced 6.66 x 10-4 moles of iodine
  • there must have been 6.66 x 10-4 / 3 - of iodate ions
  • then find concentration
54
Q

Describe a method for an iodine-sodium thiosulfare titration

A
55
Q

What errors could occur in a titration?

A
  • using contaminated apparatus could make your results inaccurate - so make sure the burette is very clean, and rinse it out with sodium thiosulfate before you start (because traces of water will dilute the solution = larger titre than normal)
  • read the burette correctly (from bottom of the meniscus, with your eyes level with liquid
  • reduce effect of random errors, you should repeat the experiment until you get at least three results within 0.1cm3 of each other
    • then take average of these results
  • remember to wash flask between repeat experiments or use a new, clean one
56
Q

What errors could occur in a sodium thiosulfate titration specifically?

A
  • the solutions your using will react very slowly with oxygen in the air, so they should be made up as fresh as possible
  • if you add starch solution too early during the titration, the iodine will stick to the starch and wont react as expected with the thiosulfate, making the result unreliable
    • only add the starch when the solution is pale yellow
57
Q

Describe the trend of volatility as you go down group 7

A
  • decreases down the group - because of the increasing strength of the instantaneous dipole - induced dipole bonds - these incease as the size and relative mass of the atoms increase
58
Q

What is volatility?

A
  • measure of how easy it is to vaporise something (turn it from a liquid to a gas).
59
Q

Describe why halogens are more soluble in organic solvents than in water

A
  • halogens exist in their natural state as covalent diatomic molecules
  • because theyre covalent and non-polar they have low solubility in water but dissolve easily in organic solvents - like hexane
60
Q

Describe the trend of reactivity down group 7

A
  • halogen atoms react by gaining an electron in their outer p sub-shell
  • this means they are reduced
  • as theyre reduced they oxidise another substance - so theyre oxidising agents
  • as you go down the group, the atoms become larger, so their outer electrons are further from the nucleus. The outer electrons are also shielded more from the attraction of the positive nucleus because there are more inner electrons. This makes it harder for larger atoms to attract the electron needed to form an ion
  • so larger atoms are les reactive and reactivity descreases down a group
61
Q

How do halogens displace halide ions from solutions

A
  • displacement reaction - type of reaction where one reactant replaces another reactant in a compound
  • the halogens relative oxidising strength can be seen in their displacement reaction with halide ions
  • more reactive halogen will replace a less reactive halide in a solution
  • chlorine diaplaces bromine and iodine
  • bromine displaces iodine
  • iodine cant displace chlorine or bromine
62
Q

What other reaction is the displacement reaction of a halogen-halide reaction

A
  • also a redox reaction
  • thing that is displaced is oxidised
  • thing that does the displacing is reduced
63
Q

What are the ionic equations for the reactions that happen if a halide is displaced with a halogen

A
64
Q

how can you tell if a displacement of a halide by a halogen has taken place

A
  • will be a colour change
  • if bromine is dispalced and bromine formed - reaction mixture will turn orange
  • if iodine is displaced and iodine is formed, the reaction mixture will turn brown
  • you can make these colour changes easier to see by shaking the reaction mixture with an organic solvent like hexane
  • the halogen thats present will dissolve in the organic solvent, which settles out as a distinct layer above the aqueous solution
65
Q

What are hydrogen halides?

A
  • hydrogen and a halogen
  • hydrogen has 1 electron in its outer shell
  • halogens have 7 electrons in their outer electron shells
  • halogens react with hydrogen to form hydrogen haldies
66
Q

How do you make a hydrogen halide?

A
  • adding concentrated acid, to a solid, ionic halide
  • to make hydrogen chloride (HCl) add concentrated phosphoric acid (H3PO4) to sodium chloride
  • all hydrogen halides can be madde this way using an ionic halide and concentrated phosphoric acid
67
Q

What is the issue of using concentrated sulfuric acid to make hydrogen halides?

A
  • H2SO4 - is an oxidising agent, so it cn get involved in redox reactions
  • you can make hydrogen chloride using sulfuric acid
    • NaCl + H2SO4 –> HCl + NaHSO4
    • Cl- + H2SO4 –> HCl + HSO4-
  • you cannot make hydrogen bromide or hydrogen iodide using sulfuric acid
  • when sodium bromide or sodium iodide reacrs with sulfuric acid, the bromide or iodide ions are oxidised to make bromine or iodine gas
  • this is because iodine and bromine are are strong enough reducing agents to reduce sulfuric acid
  • when you add sulfuric acid to sodium bromide, the bromide ions are oxidised to bromine gas and the sulfuric acid is reduced from sulfuric acid to sulfur dioxide
    • 2NaBr + 2H2SO4 –> Na2SO4 + Br + SO2 + H2O
    • H2SO4 + 2H+ + 2Br- –> Br2 + SO2 + 2H2O
  • iodine is such a strong reducing agent that when you add sulfuric acid to sodium iodide it reduces it from sulfuric acid to hydrogen sulfide
    • 8NaI + 5H2SO4 –> 4Na2SO4 + 4I2 + H2S + 4H2O
    • H2SO4 + 8H+ + 8I- –> 4I2 + H2S + 4H2O
68
Q

When heated which hydrogen halides are stable

A
  • when heated, hydrogen fluoride and hydrogen chloride are stable, and won’t split up into hydrogen and halide ions
  • hydrogen bromide will split slightly when heated, and hydrogen iodide even more so
  • this is because of the strength of the hydrogen-halide bonds
  • as you go down group 7 - the strength of the bond that the halide forms gets weaker
  • this is because the halogen atoms get bigger down the group, meaning the bonding electrons are further away from the nucleus and shielded by more inner electrons shells
69
Q

Are hydrogen halides acidic, neutral or alkali?

A
  • hydrogen chloride, hydrogen bromide and hydrogen iodide all disoolve in water to create strong acids
  • when a hydrogen halide is dissolved in water, it dissociates. This just means that the molecule splits apart to form two ions - in this case a hydrogen ion and a halide ion
  • its the hydrogen ions that make the solutions acidic
  • hydrogen fluoride is an exception to the rule - it doesn’t fully dissociate in water (only a few of the moelcules split apart) its still acidic but, its a weak acid
70
Q

How do hydrogen halides react with ammonia?

A
  • ammonia (NH3) is a base, so it can accept a proton to from positively charge ammonium ion (NH4+)
  • the ammonium ion can bond with a negative halide ion, to produce an ammonium halide
71
Q

How do hydrogen halides reduce sulfuric acid

A
  • hydrogen fluoride and hydrogen chloride dont react with sulfuric acid -they dont reduce sulfuric acid - they are not strong enough reducing agents to reduce the sulfur
  • hydrogen bromide and hydrogen iodide do react with sulfuric acid though
  • similar reactions to sodium halide and sulfuric acid
72
Q

How do silver ions react with halide ions

A
  • this can be used to test for halides
  • add dilute nitric acid to remove ions which might interfere with the reaction
  • then add silver nitrate solution (AgNO3(aq))
  • a precipitate of silver halide is formed
  • the colour of the precipitate identifies the halide present in the original solution
73
Q

how do you tell silver halides apart from each other?

A
  • precipitates can look quite similar, so can be difficult to identify a halide
  • add ammonia solution
74
Q

Describe how chlorine is stored and transported

A
  • chlorine is a very dangerous chemical
  • its toxic and corrosive, so must be kept away from skin and eyes
  • its also harmful if its breathed in - it irritates the respiratory system
  • chlorine must be kept away from flammable materials too
  • being an oxidising agent, it can increase fire risks
  • because of these dangers, chlorine has to be stored and transported very carefully
  • its usually kept as a liquid under pressure in small cylindes
75
Q

Why is chlorine still produced despite being a dangerous chemical?

A
  • has some really important uses
  • chlorine is an important part of water treatment
  • adding chlorine to water sterilises it, making it safe to drink or swim in
  • it kills disease-causing microorganisms, such as bacteria
  • if we didn’t treat our drinking water with chlorine, we’d be at risk of getting all sorts of nasty and potentially dangerous infections
  • chlorine also used to make bleach
76
Q

What is a reversible reaction?

A
  • reaction that goes both ways
  • ⇌ used in the reaction
  • as the reactants get used up, the forward reaction slows down - and as more product is formed, the reverse reaction speeds up
  • after a while the forward reaction will be going at the same rate as the backward reaction
  • the amount of reactants and products wont be changing, so it seems like nothing is happening - called dynamic equilibrium
  • dynamic equilibrium can only happen in a closed system - nothing can get out or get in
  • postion of equilibrium tells you how much reactant and how much product youll have at equilibrium
    • if equilibrium lies to left then youll have more reactants than products
    • if equilibrium lies to the right youll have more products than reactants
77
Q

What is Kc

A
  • an equilibrium constant
  • if you know the molar concentration of each substance at equilibrium, you can work out the equilibrium constant, Kc
  • value of Kc will only be true for that particular temperature
78
Q

describe how Kc tells you the position of equilibrium

A
  • Kc >> 1
    • equilibrium lies far to the right - there are much more products than reactants present at equilibrium
  • Kc > 1
    • equilibrium lies to the right - there are slightly more products than reactants present at equilibrium
  • Kc = 1
    • equilibrium lies in the middle - the amount of products and reactants at equilibrium will be the same
  • Kc < 1
    • equilibrium lies to the left - there are slightly more reactants than products present at equilibrium
  • Kc << 1
    • equilibrium lies far to the left - there are much more reactants than products present at equilibrium
79
Q

How does changing the concentration of reactants or products change Kc

A
  • the position of equilibrium will shift to keep Kc constant
80
Q

Hwo does pressure alter equilibrium position

A
  • (changing this only affects the equilibria involving gases)
  • increasing the pressure shifts the equilibrium to the side with fewer gas molecules. this reduces the pressure
  • decreasing the pressure shifts the equilibrium to the side with more gas molecules. this raises the pressure
81
Q

How does temperature alter the equilibrium position

A
  • increasing the temperature means adding heat
  • the equilibrium will shift in the endothermic position (positive ΔH) direction to absorb this heat
  • decreasing the temperature removes heat
  • the equilbrium shifts in the exothermic (negative ΔH) direction to try to replace the heat

CHEMISTRY AS (24 decks)

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