Electrochemistry, Redox & Group VII [Oxidation & Reduction, Electrolysis, Electroplating, Halogens, Hydrogen Fuel Cell] Flashcards
(106 cards)
1
Q
Roman numeral to indicate
A
indicate the oxidation
number of an element in a compound
2
Q
define electrolysis - for understanding
A
the breakdown of an ionic compound, in molten or aqueous solution, by the passage of electric current.
- electric current is used to drive a non-spontaneous redox reaction.
current provides voltage that forces electrons to be gained and lost when this would not otherwise be possible.
3
Q
Electrolysis is important in ….
A
in the commercial production of obtaining metals from their ores/ions
and oxidising halogen ions into their elements.
4
Q
redox reactions - DEFINE as
identify -
A
involving simultaneous
oxidation and reduction
redox reactions as reactions involving gain and loss of oxygen
reactions involving
gain and loss of electrons
5
Q
define oxidation
A
gain of oxygen [loss of H2]
(a) loss of electrons [ON PRODUCTS elec ox]
(b) an increase in oxidation number
+ water in reactants ox
6
Q
define reduction
A
loss of oxygen [gain of H2]
(a) gain of electrons [ON REACTANTS elec red]
(b) a decrease in oxidation number
+water in products red
7
Q
Identify redox reactions by changes in oxidation
number using:
4 facts
A
(a) the oxidation number of elements in their
uncombined state is zero
of elements = 0
8
Q
b)
A
(b) the oxidation number of a monatomic ion is
the same as the charge on the ion
+ on polyatomic ion equals the overall charge of ion
9
Q
c)
A
c) the sum of the oxidation numbers in a
compound is zero
10
Q
d)
A
(d) the sum of the oxidation numbers in an ion is
equal to the charge on the ion
11
Q
Define an oxidising agent as
A
a substance that
oxidises another substance and is itself reduced
12
Q
define a reducing agent as
A
a substance that
reduces another substance and is itself oxidised
13
Q
3 types of electrolytes
A
molten
dilute aqueous solution
concentrated aqueous solution
14
Q
cathode ;; electrons are GAINED by the positive ions (CATIONS) and transferred from the CATHODE (-).
Attracts the positive ions (CATIONS)
A
metals or hydrogen are formed at
the cathode
CATHODE;
molten - metal
dilute - less reactive metal/h2
concentrated - less reactive metal/h2
15
Q
anode ;; electrons are LOST from the negative ions (ANIONS) and transferred to the anode (+).
Attracts the negative ions (ANIONS)
A
non-metals (other than
hydrogen) are formed at the anode
ANODE;
molten - non-metal
dilute - oxygen
concentrated - halogen gas/O2
16
Q
in molten electrolytes there are
MOLTEN - NON-METAL AN, METALCAT
A
- only pos and neg ions from the ionic salt that are attracted to the electrodes
e.g. NaCl (l)
anox: non-metal forms. half-eqn: 2Cl- → Cl2 + 2e-
redcat: sodium metal forms. half-eqn: Na+ + e- → Na
17
Q
dilute aqueous electrolytes, there are…
O2 AN, H2/LESS REACTVE CAT
A
- pos and neg ions from the salt ANDD a large amount of H+ and OH- ions from water
18
Q
at the anode (oxid.)
A
negative non-metal ions and OH- ions from water are attracted
Only O2 will form
19
Q
at the cathode (reduc.)
DILUTE SOLUTION
A
Positive metal ions and H+ ions from water are attracted.
Either metal or H2 will form, depending on reactivity of metal
MORE REACTIVE THAN H+ => H2 FORMS!
20
Q
concentrated aqueous electrolytes, there are…
HALOGEN GAS/O2 AN, less reactive metalCAT/H2
A
- a LARGE amount of pos and neg ions from the salt ANDD H+ and OH- ions from water
21
Q
at the anode (oxid.)
CONCENT.
A
- negative non-metal ions and OH- ions from water are attracted
if HALIDE ions present, HALOGEN gas forms.
otherwise O2 forms.
22
Q
at cathode (reduc.)
CONCENTRATED
A
positive metal ions and H+ ions from water attracted
either metal or H2 will form, depending on reactivity of metal.
more reactive => H2 forms
23
Q
redcat
[cathode negative]
A
cathode reduction occurs
pos ions go here ?????
24
Q
anox
A
25
intert electrode
graphite, platinum
26
aqueous & dilute
gases at both electrodes
27
bromine
28
MOLTEN CaI2 electrolysed
at anode? pos
cathode? neg
anode: iodine
cathode: calcium
29
Chemical cells include batteries used in everyday appliances and cars.
they produce _____________ until __________
e.g.: fuel cells
Chemical cells produce a voltage until one of reactants is used up. E.g., fuel cell
30
Spontaneous reactions
a reaction that proceeds on its own in a given set of conditions, without the external addition of energy.
electrolytic cells - non-spontaneous
31
fuel cells - the eg, - they:
are ELECTROCHEMICAL CELLS
fuel + O2 -> electrical energy
- produce voltage continuously, as long as they are supplied with
-> a constant supply of a suitable fuel
-> oxygen e.g. from air
-> fuel is oxidised electrochemically, rather than being burned, so reaction takes place at a lower temp than if it was to be burned. energy released as electrical energy, NOT "heat"
32
hydrogen-oxygen fuel cells are an alt to rechargeable cells & batteries
H2 & O2 used to produce voltage
product? what is reaction?
water
hydrogen + oxygen -> water
2H2 (g) + O2 (g) -> 2H2O (l)
33
half eqns
- REMEMBER
cathode: O2 + 4H+ + 4e- -> 2H2O
anode: H2 -> 2H+ + 2e
34
advantages
- Hydrogen can be produced from water so the process is renewable
- do not produce any pollution: only product is water whereas petrol engines produce carbon dioxide, and oxides of nitrogen
- release more energy per kilogram than either petrol or diesel
- No power is lost in transmission as there are no moving parts, unlike an internal combustion engine
- Quieter so less noise pollution compared to a petrol engine
35
disadvantages
- Hydrogen obtained by methods that involve: 1. The combustion of fossil fuels releases carbon dioxide and other pollutants into the atmosphere;; 2. The electrolysis of water requires large amounts of electricity to produce
- Materials used in producing fuel cells are expensive
- Hydrogen is more difficult and expensive to store compared to petrol as it is very flammable and easily explodes when under pressure
- Fuel cells are affected by low temperatures, becoming less efficient
- There are only a small number of hydrogen filling stations across the country
36
NaCl electrolyte
sodium chloride
sodium + cathode -
chlorine - anode +
MOLTEN
2Na+ + 2e- -> 2Na
2Cl- -> Cl2 + 2e-
bubbles, at cathode solid sodium forms, silver colour
37
products formed at the electrodes? + describe the observations made during the electrolysis of:
(a) molten lead(II) bromide
(b) concentrated aqueous sodium chloride
(c) dilute sulfuric acid
using inert electrodes made of platinum or carbon/ graphite
a)
b)
c)
38
why do we use graphite
- inert electrode (INERT = DOESN'T REACT)
- conducts electricity bc every C bonds to 3 or 4 other Cs so electrons free to carry charge and conduct electricity
39
what is electroplating?
metal objects are electroplated to
...
- a process where the surface of one metal is coated with a layer of a different metal
;; improve their appearance [shiny] and resistance to corrosion
40
electroplating uses...
uses electrolysis to coat one metal on another
OR purify an impure metal
41
metal A at cathode/neg - COATED, object electroplated
metal B at anode/pos - doing the coating;; made from the pure metal that will be plated onto the object
42
electrolyte
AQUEOUS solution
of a soluble salt of the pure metal at the anode
43
metal anode...
less reactive than hydrogen,
placed in electrolyte
containing same metal cation
e.g. copper anode placed in copper sulfate
; copper metal coats cathode
44
Describe how metals are electroplated, e.g., tin and iron
- at anode
Tin atoms lose electrons to form tin ions in solution
The loss of electrons is oxidation
45
at cathode
Tin ions gain electrons to form tin atoms
The gain of electrons is reduction
The tin atoms are deposited on the strip of iron metal, coating it with a layer of tin
46
common metals used for electroplating
copper, chromium, silver, tin
chromium doesnt corrode, resists scratching, can be polished so is attractive shiny
47
Group VII halogens, chlorine,
bromine and iodine,
RMBR- TRENDS
- poisonous
diatomic non-metals with
general trends down the group
(a) increasing density
(b) decreasing reactivity GOING DOWN
ONE AT TOP MOST REACTIVE
48
why are elements at top more reactive w/ halogens
As you go down Group 7 , the number of shells of electrons increases
so the outer electrons are closer to the nucleus so there are stronger electrostatic forces of attraction, which help to attract the extra electron needed
This allows an electron to be attracted more readily, so the higher up the element is in Group 7 then the more reactive it is
49
appearance of the halogens at r.t.p
bc?
chlorine, a pale yellow-green gas
bromine, a red-brown liquid
iodine, a grey-black solid
bc density of the halogens increases as you go down the group:
[in water: chlorine colourness, bromine orange, iodine brown]
50
halogens have...
form halide ions by
7 electrons in their outer shell
- gaining one more electron to complete their outer shells
51
right side, non-metal, 7, diatomic, negative 1 ions
52
Higher up in group, more reactive -
Cl2 + 2KBr -> 2KCl + Br2
Cl2 + 2KI -> 2KCl + I2
Br2 + 2KI ->2KBr + I2
TAKES THE K
Reactivity decreases down group; more reactive halogen displacing less reactive halogen from its halide
fluorine will displace Cl, Br, I
Cl displaces Br and I
Br displaces I
53
electrolysis
occurs because the charged ions in the electrolyte allow electrons (electricity) to flow.
54
where do electrons move
electrons move through the external circuit TOWARDS the cathode.
anode to cathode
55
ion ORDER
- positive
K
Na
Ca
Mg
Al
Zn
Fe
Sn
Pb
HYDROGEN
Cu
Ag
SO FOR CU AND AG IONS ATTRACTED TO CATHODE (NEG), CU AND AG FORMS
for others, hydrogen forms
56
NEGATIVE ions
F
SO4
NO3
Cl
Br
I
OH
57
CONCENTRATED CuSO4
anode, oxida. - SO42- and OH- attracted. SO42- isn't a halide.
HALF EQN AT THE ANODE IN CONCEN.
2OH- -> O2 + 2H+ + 4e-
O2 would be formed at the anode in the oxidation of OH- to O2.
58
at the cathode, REDUC. - CONC. CuSO4
Cu2+ and H+ attracted
Cu less reactive than H
HALF-EQN at CAT => Cu2+ + 2e- -> Cu
Cu would be formed at the cathode in the reduction of Cu2+ to Cu
59
DILUTE H2SO4
at anode, OX
SO42- and OH- attracted
only O2 forms IN DILUTE....!!!! ANODE
HALF-EQN AT ANODE IN DILUTE
2OH- -> O2 + 2H+ + 4e-
O2 would be formed at the anode in the oxidation of OH- to O2.
SAME as b4
60
at cathode - reduc.
ONLY H+ attracted
2H= + 2e- => H2
H2 would be formed at the cathode in the reduction of H+ to H2.
61
mass of negative electrode (cathode) increases
The gain in mass by the negative electrode is the same as the loss in mass by the positive electrode. So the copper deposited on the negative electrode must be the same copper ions that are lost from the positive electrode.
62
electrodes
-simply present to allow a flow of electrons.
-must be able to conduct electricity but NOT react with the electrolytes.
INERT electrodes are used, e.g. graphite OR platinum.
63
IONS must be free to move -
electrolyte must be molten for electrolysis to work. explain
IONS!!!
64
The greater the difference in reactivity between the two metals, the greater the voltage.
RMBR ;;
[chemicals in the cell eventually run out, so the reactions eventually stop; means that the cell can only produce electricity for a limited period of time.]
65
A piece of white paper was coated with silver bromide and exposed to the light. Sections of the
paper were covered as shown in the diagram.
Predict the appearance of the different sections of the paper after exposure to the light and the
removal of the card. Explain your predictions. [4]
prediction:
the ‘not covered’ section will be black;
the ‘covered in thick card’ section will be white/ cream;
the ‘covered in thin card’ section will be grey;
explanation:
the more light, the more silver ions are reduced;
[darker under light]
66
Describe two things that the student would observe when the mixture is added to the dilute sulfuric acid. [2]
bubbles
solution turns blue [bc turns to copper sulfate]
67
Describe how the student can produce pure dry copper from the mixture of copper and
copper(II) sulfate solution. [3]
filter/ centrifuge/decant
wash with (distilled) water
(dry with) filter paper/ tissues
68
The student then adds sodium hydroxide solution to the copper(II) sulfate solution to produce
copper(II) hydroxide.
i) Write an ionic equation for this reaction.
ii) After separating the copper(II) hydroxide from the mixture, the copper(II) hydroxide is heated
strongly. The copper(II) hydroxide decomposes into copper(II) oxide and steam.
(i) Write an equation for the decomposition of copper(II) hydroxide. Include state symbols.
i) Cu2+ + 2OH → Cu(OH)2
ii) Cu(OH)2 (s) → CuO (s) + H2O (g)
69
Explain why the positive ions in the above equations are oxidising agents.
first thought is "because they are reduced" BUT ALSO WRITE
"because they can accept electrons"
70
Which metals in the series above do not react with dilute acids to form hydrogen?
copper and mercury and palladium
71
Describe an experiment which would confi rm the prediction made in (c)(i).
add copper / mercury / metal to (named) acid and no reaction
72
Describe what you would observe when zinc, a reducing agent, is added to this
pink /purple solution
(solution) becomes colourless
73
Why do the oxidation states increase from sodium to silicon?
number of electrons to be lost
74
After Group(IV) the oxidation states are negative and decrease across the period.
Explain why.
gain electrons
number of electrons to be gained is less across period
75
Explain why the concentration of copper ions in the electrolyte remains constant throughout
step 1.
formation of Cu2+ /copper ions at the anode happens at the same rate as;
removal of Cu2+ /copper ions at the cathode
76
Give two changes which would be needed in order to coat nickel onto the object in step 2.
replace (anode of) copper with nickel;
replace electrolyte with nickel(II) sulfate/NiSO4;
77
Give three different properties of transition metals which are not typical of other metals.
(good) catalysts;
variable oxidation numbers;
form coloured compounds
78
Aluminium is extracted by the electrolysis of a molten mixture which contains aluminium oxide,
Al 2O3. This decomposes to form aluminium at the negative electrode and oxygen at the positive
electrode.
Is the reaction exothermic or endothermic? Explain
endothermic AND (electrical) energy supplied;
79
Explain why the mass of the magnesium electrode decreases and the mass of the copper
electrode increases.
[2]
mg and cu electrode
magnesium loses electrons & is oxidised; DECREASES
copper ions gain electrons and are reduced; INCREASES
80
How could you use this cell to determine which is the more reactive metal, magnesium or
manganese? [2]
set up a magnesium/ manganese cell;
the negative electrode (cathode,, is the more reactive) OR the electrode that loses
mass (is more reactive);
81
Suggest why the electrolyte is a paste.
to allow ions to flow
82
At which electrode does oxidation occur? Explain
anode;
electrons lost / electrons
move from this electrode
83
ionic equation for the reaction at anode
hydrogen oxygen fuel cell
H2 → 2H+ + 2e-
84
Give two advantages of a fuel cell over a gasoline-fuelled engine.
lightweight
quieter
fewer working parts/less maintenance
more efficient
H2O is the only product
85
Chlorine is made by the electrolysis of concentrated aqueous sodium chloride.
Describe this electrolysis.
Write ionic equations for the reactions at the electrodes and name
the sodium compound formed. [5]
products, compound formed, half-eqn, half-eqn, product at anode&cathode
hydrogen and chlorine (1)
sodium hydroxide (1)
2H+ + 2e → H2
2Cl- → Cl2 + 2e
Hydrogen at cathode and chlorine at anode (1)
86
Aluminium is obtained from purified alumina, Al2O3, by electrolysis
Describe the extraction of aluminium from alumina.
Include the electrolyte, the electrodes and the reactions at the electrodes. [6 MARKS]
electrolyte alumina/aluminium oxide dissolved in molten cryolite (1)
use cryolite to reduce mp / temperature of electrolyte 900 to 1000°C (1)
electrodes carbon (1)
aluminium formed at cathode/Al
3+ + 3e → Al (1)
oxygen formed at anode/2O2- → O2 + 4e (1)
anode burns/reacts to carbon dioxide/C + O2 → CO2 (1)
87
The reaction between calcium and nitrogen to form calcium nitride is a redox reaction.
In terms of electron transfer, explain why calcium is the reducing agent. [3]
Calcium loses electrons
(these are) gained by nitrogen
calcium/Ca is oxidised (by electron loss) therefore calcium is the
reducing agent
88
A piece of iron was weighed and placed in the apparatus. It was removed at regular
intervals and the clock was paused. The piece of iron was washed, dried, weighed and replaced.
The clock was restarted.
This was continued until the solution was colourless.
The mass of iron was plotted against time. The graph shows the results obtained.
(i) Suggest an explanation for the shape of the graph. [3]
rate of reaction decreases / gradient decreases [1]
because concentration of bromine decreases [1]
reaction stops because all bromine is used up [1]
89
Predict the shape of the graph if a similar piece of iron with a much rougher surface
had been used.
Explain your answer [2]
initial rate greater / gradient greater
because bigger surface area
90
Describe how you could find out if the rate of this reaction depended on the speed of stirring. [2]
increase rate of stirring
measure new rate / compare results
91
Describe how you could test the solution to fi nd out which ion, Fe2+ or Fe3+, is present. [3]
add sodium hydroxide solution / ammonia(aq)
Fe2+ green precipitate
Fe3+ brown precipitate
92
Explain, using the above data, why the forward reaction is exothermic. [2]
More heat given out bond forming than taken in bond breaking
// ;; More heat given out than taken in [1]
–2328 + 945 + 1308 = –75(kJ) [1]
93
oxygen at anode half-eqn for H2SO4 - DILUTE
RMBR.
oxygen - Aqueous tin(II) sulfate
2OH- -> O2 + 2H+ + 4e-
RMBR.
4OH → O2 + 2H2O + 4e
94
Copper can be extracted by the electrolysis of an aqueous solution.
Suggest why the electrolysis of an aqueous solution cannot be used to extract aluminium. [2]
aluminium more reactive than copper / aluminium higher in reactivity series
hydrogen not aluminium formed at cathode
95
The ions which are involved in the electrolysis are Al 3+ and O2–. The products of this
electrolysis are given on the diagram.
Explain how they are formed. Use equations where appropriate. [4 MARKS]
Al 3+ + 3e → Al [1]
2O2- → O2 + 4e [2]
oxygen reacts with carbon (anode) to form carbon dioxide / C + O2 → CO2 [1]
OR;; aluminium ions accept electrons / are reduced [OPT. 1]
oxide ion loses electrons / is oxidised [1]
96
Give two reasons why the oxide is dissolved in cryolite.
reduce temp
better conductivity
97
Why do the carbon anodes need to be replaced frequently?
they burn (away) / react with oxygen / form carbon dioxide
98
The hydrogen ions are from the water.
H2O -> H+ + OH–
Suggest an explanation why the concentration of hydroxide ions increases. [2]
H+ removed [1]
(equilibrium) moves to RHS / forward reaction favoured / more water molecules ionise [1]
99
(i) How could you measure the rate of decomposition of sodium chlorate(I)?
.............................................................................................................................. [1]
(ii) Describe how you could show that the rate of decomposition of sodium chlorate(I) is
a photochemical reaction.
(reference to) volume and time / how long it takes
(ii) carry out experiment with different intensities of light / one in light and one in
dark / repeat experiment in reduced light
measure new rate which would be faster or slower depending on light intensity
100
____ sulfate change to electrolyte
sulfuric acid at end
101
i) dilute aq NaCl - sodium chloride remains [water used up or solution becomes more concentrated] - hydrogen, oxygen
ii) electrolysis of CONCENTRATED (!!) aqueous sodium chloride produces three commercially important chemicals hydrogen, chlorine and sodium hydroxide.
(a) The ions present are Na+ (aq), H+
(aq) ,Cl (aq) and OH (aq).
Explain why the solution changes from sodium chloride to sodium hydroxide. [1]
Na+ and OH- are left
OR Cl removed OH left
SO THIS IS BC Cl NOT PRODUCED IN DILUTE bc at anode OXYGEN FORMS IN DILUTE
so Cl removed in concentrated NaCl, OH left
CL NOT REMOVED IN DILUTE SO NaCl remains.
102
The coating on both of the other two pieces was scratched, exposing the steel.
The piece of steel coated with zinc still did not rust but the copper-coated piece of steel rusted very rapidly.
Explain these observations in terms of the formation of ions and the transfer of electrons. [4]
Zn more reactive than Fe ;
Zn loses / transfers electrons (more readily) and forms (+ve) ions
Fe is more reactive than Cu;
Fe loses / transfers electrons (more readily) and forms (+ve) ions
103
When aqueous lithium bromide is electrolysed, a colourless gas is formed at the negative elctrode and the solution becomes alkaline.
Explain these observations and include an equation [3]
(gas) hydrogen (given off at cathode) ;
lithium hydroxide / LiOH are alkali(ne);
2H+ + 2e → H2
104
The electrolysis of CONCENTRATED aqueous sodium chloride forms three products. They are
hydrogen, CHLORINE and sodium hydroxide. SOOOO CL REMOVED
(i) Explain how these three products are formed. Give ionic equations for the reactions at the
electrodes.
2Cl → Cl2 + 2e-
2H+ + 2e- → H2
hydrogen formed at cathode and chlorine at anode
sodium ions and hydroxide ions left in solution become sodium hydroxide
105
The purification of bauxite uses large amounts of sodium hydroxide.
(i) Describe the chemistry of how sodium hydroxide is made from concentrated aqueous sodium chloride.
The description must include at least one ionic equation. [5]
electrolysis
chlorine formed at anode (positive electrode)
hydrogen formed at cathode (negative electrode)
2Cl → Cl2 + 2e or 2H+ + 2e → H2
sodium
hydroxide left behind/remains in solution;
106
A reading on the voltmeter shows that electrical energy is being produced. Suggest an
explanation for how this energy is produced.
it is a cell; [1]
hydrogen reacts with oxygen; [1]
this reaction produces energy / changes chemical energy to electrical energy
