1
Q
galvanic vs electrolytic cells
A
galvanic cells:
- produces electricity
- spontaneous reactions
- convert chemical energy to electrical
- oxidation anode
- reduction cathode
- anode is negative
- cathode is positive
electrolytic cells
- consume electricity
- have non-spontaneous reaction
- convert electrical energy to chemical energy
- oxidation anode
- reduction cathode
- anode is positive
- cathode is negative
2
Q
ELECTROLYSIS
A
- involves the passage of electrical energy from a power supply such as a battery, through a conducting liquid
- the use of electrical energy drives a non-spontaneous redox reaction
3
Q
electrolysis of molten NaCl
A
- inert electrodes → platinum or graphite
- the electrolyte (conducting liquid) is molten NaCl → no H2O present
- NaCl melts at 801 degrees Celsius
- negative electrode connected to the negative terminal
- positive electrode connected to the positive terminal
- electrolysis takes place in a single container
- products must be kept apart or they will spontaneously react to reform original reactants
4
Q
competition at electrodes
A
- in some electrolytic cells there may be several chemicals present that can react
- water is a potential reactant when aqueous electrolytes are used (look for solution word)
- reactive material used for electrodes may also participate in the reaction
5
Q
standard conditions - factors affecting reactions
A
- electrochemical series based on standard conditions
- most electrolysis reactions aren’t performed at standard conditions
- reactions are affected by electrolyte concentration, gas pressures. current, voltage and electrode types
6
Q
electrolysis of NaCl solution
A
at the anode
- electrons are withdrawn from this electrode
- two possible reductants 2Cl- or H2O(l)
- reactions lower in the electrochemical series are stronger reductants and more likely to occur
- reaction would be the oxidation of water
at the cathode
- power supply pushes electrons towards the negative electrode
- 2 possible oxidants H2O(l) and Na+
- reactions higher in the electrochemical series involves stronger oxidants and are more likely to occur
- water will be reduced and H2 formed
7
Q
electrolysis of NaCl solution in real life
A
- standard conditions
- water react at both anode and cathode
- in practice
- possible for either anode reactions to occurs as they are close to each other on the electrochemical series
- when [NaCl] increases, Cl2(g) can be produces
8
Q
use of electrolysis in industry
A
- chemical industries tend to avoid using electrolysis
- high cost of electrical energy and high energy usage
- but process enables some chemical to be produced that could not be readily produced in any other way
- creating sodium metal and chlorine gas
- downs cell - molten NaCl
- membrane cell - aqueous NaCl
9
Q
why is Fe not oxidised in downs cell
A
- it is connected to the negative terminal of the power supply
- continuous supply of electrons prevent it from oxidising
10
Q
advantages and disadvantages of molten electrolyte
A
advantages
- no interference of water in reactions
disadvantage
- process req more energy
- operates at a high temp
- a mixture of NaCl and CaCl2 decreases the melting temperature
- saves energy costs
11
Q
aqueous electrolyte - membrane cell
A
- preferable
- products of electrolysis
- sodium hydroxide
- chlorine
- hydrogen
- highly concentrated NaCl solution used as an electrolyte
- anode and cathode separated by a semipermeable membrane
- prevents contact between products and only allow Na+ thru
12
Q
membrane cell advantage
A
- NaOH not contaminated w NaCl, Cl2, H2
- enables process to occur between 80-90
- no need to heat electrolyte
- reduced cost of production
13
Q
advantages of use of reactive electrodes
A
- many commercial electrolytic cells use inert electrodes, however reactive electrodes have benefits
- can be used to purify metals
- impure metal used as anode and pure metal deposited at cathode
- production of aluminum for alumina
- can be used to purify metals
14
Q
electroplating cells
A
- results in the application of a thin layer of metal over another surface
- when the cell is in operation the power supply acts as an electron pump, pushing electrons onto the negative electrode + removing electrons from positive electrode
- object being plated is at the cathode - negative
- an electrode of the metal is at the anode - positive
15
Q
faradays first law
A
- amount of substance deposited is measured directly by taking the mass of metal formed at the cathode
- the amount of any substance deposited, evolved or dissolved at an electrode is directly proportional to the quantity of electrical charge passed through the cell
- 1A indicated that 1C (6.24 x 10^18 electrons) of charge flow every secomd
- the more charge that passes through the cell, the more metals form at the cathode
- charge on metal ION determines how much metal is deposited
16
Q
faraday
A
- amount of charge on 1 mole of electron
- 1F = 96500 C
17
Q
faraday’s second law
A
- to produce one mole of a substance, use mole ratio to find the no. moles of electrons that must be consumed
- no. moles of electrons which carry the same no. f faradays correspond to the given charge in coulombs.
18
Q
conditions used in chemical industry
A
- to reduce waste and reduce costs, conditions are carefully selected to ensure acceptable yields are obtained
- optimum conditions used to make production of useful chemicals economically viable and environmentally responsible
- reaction conditions are manipulated to
- maximise product yield
- maximise reaction rate
- minimise costs
- maximise overall efficiency
19
Q
conditions to increase reaction rate
A
- high concentration/pressure
- high temperatures
- high S.A of solids
- use of a catalyst
20
Q
conditions to increase equilibrium yield
A
- pressure depending on no. particles
- temp depending if exo or endo
- additions of excess reactant
- removal of product as it forms
- consideration incl. product yield, reaction rate, safety hazards, environmental concerns, costs
21
Q
primary vs secondary cells
A
primary:
- disposable
- non-rechargeable
- go flat when reaction reaches equilibrium
- eg alkaline cells
- the products slowly migrate away from electrodes or are consumed by side reactions occurring in the cell preventing the cells from being recharged
secondary
- known as rechargeable cells or accumulators
- can be reused many times
- rechargeable
- eg. lithium ion cells
22
Q
secondary cells
A
- can undergo hundreds of recharges
- chemical reaction gets reversed
- recharged by connecting to an electrical power supply
- input of energy drives a non-spontaneous reaction
- the positive terminal of the charger is connected to the cell’s positive electrode
- negative terminal of charger connected to the cell’s negative electrode
- charge of an electrode will remain the same during discharge
and recharge.
23
Q
conditions required for recharging
A
- the power supply must have a potential difference (voltage) greater than that produced by the cell during discharge
- electrodes must not be damaged and products must remain in contact with the electrodes
24
Q
discharge vs recharge
A
- when a secondary cell discharges, it acts as a galvanic cell
- chemical to electrical energy
- spontaneous reaction
- negative terminal - anode
- when a secondary cell recharges, it acts as an electrolytic cell
- electrical to chemical
- non-spontaneous
- positive terminal - anode
25
electrolyser
is a system that uses electricity to break water into hydrogen and oxygen via electrolysis
26
PEM electrolyser vs simple electrolyser
- PEM electrolyser does not operate in an aqueous environment like a simple electrolyser
- It uses a conductive polymer electrolyte.
- increasing the PEM thickness would decrease efficiency of fuel cell -> **due to added resistance to proton flow**
- The overall equation is the same, but the half-equations differ.
- The overall equation for both will be 2H2O l 2H2 + O2. The states will differ.
27
hydrogen gas as a fuel
- is an attractive alternative to carbon based fuels because:
- it is abundant on earth in the form of water and most carbon compounds
- has a high energy density → 1 g releases 141kJ of energy whereas 1g of methane releases 55.6kJ
- combusting hydrogen gas only gives water
28
production of hydrogen gas
- most produced from steam reforming of methane
- CO is converted to CO2
-** hydrogen produced from fossil fuels is known as brown hydrogen**
- as co2 released is harmful for the environment
29
green hydrogen
- H2 produced through the electrolysis of water (RENEWABLE FEEDSTOCK) USING RENEWABLE ENERGY
- overall 2H2O(l) → 2H2(g) + O2(g)
- anode 2H2O(l) → O2(g) + 4H+(aq) +4e-
- cathode 2H+(aq)+2e- → H2(g)
## Footnote
SUSTAINABILITY CHALLENGES WITH USING WATER AS RENWABLE FEEDSTOCK:
* CHALLENGE TO MAINTAIN SUFFICIENT WATER SUPPLY DURING DROUGHT
* LIMITED SUPPLY OF WATER FOR DIRNKING
* LIMITED SUPPLY OF WATER FOR SANITATION, WASHING AND BATHING
30
industrial production of hydrogen
- doesn’t use simple electrolysis
- renewable energy sources don’t provide stable supply of energy req to run process
- acidic or alkaline solutions are corrosive
- hydrogen gas produced is NOT compressed
- industry uses polymer electrolyte membrane (PEM) electrolyser to electrolyse water
31
PEM to electrolyse water
- water is oxidised at the anode
- H+ ions move through conducting polymer to cathode
- H+ ions reduced at cathode to produce COMPRESSED H2 - no need for compressor stage
- electrodes are made from exp metals which allows gases through but not liquids
- metals act as catalysts to improve efficiency of gas production at electrodes
- membrane contains advanced polymers which allow H+ through but not electrons
ANODE
2H2O(l) → 4H+(aq) + O2(g) + 4e−
CATHODE
4H+(aq) + 4e− → 2H2(g)
32
advantages and disadvantages of PEM
- advantages
- high rate of hydrogen production
- disadvantages
- exp bc of expensive metals used as catalysts
33
hydrogen cars
- comparable to conventional vehicles
- hydrogen gas distribution system needs to be established
- hydrogen tank needs to be fitted such that it won't explode in the event of a car crash
34
artificial photosynthesis
- **photoelectrochemical cell:** reactions involving an electrical current generated by the action of light
- sunlight lands on anode
- energy from the sun causes excitation of the metals in the electrode
- H+ migrate from anode to cathode
- H2 produced instead of glucose
- Water oxidation in acid and using catalysts: 2H2O(l) → 4H+ + O2(g) + 4e−
* Proton reduction in the presence of catalysts: 4H+ + 4e− → 2H2(g)
* Overall: 2H2O(l) → 2H2(g) + O2(g)
35
advantages of artificial photosynthesis
- does not create greenhouse gases (liquid water is produced).
* does not require the use of fossil fuels.
36
lead acid batteries
- will need replacing as PbSO4 breaks away from the electrode
- anode reaction during discharge:
Pb(s)+ SO42-(aq) → PbSO4(s)+ 2e-
- cathode reaction during discharge
PbO2(s) + 4H+(aq) + SO42-(aq) + 2e- → PbSO4(s) + 2H2O(l)
- OVERALL:
Pb(s) + PbO2(s) + 4H+(aq) + 2SO42-(aq) → 2PbSO4(s) + 2H2O(l)
37
commercial electrolytic cells
- electrolysis is the only way to produce reactive metals - K, Na, Ca, Li, Al (but not in aq solutions)
- high temp use energy - costs money
- aq electrolysis useful for NaOH and Cl2 production
- products must be kept separate - otherwise will spontaneously react
- electrodes are inert
38
green hydrogen advantage and disadvantage
**advantage**
* the only product of combustion is water
**disadvantage**
* limited infrastructure for production, storage and distribution
* electrolysis process is energy intensive
39
why is it necessary to employ hydrogen compressor
* crucial for storing hydrogen gas efficiently due to its very low
density at standard atmospheric pressure
* Compressing hydrogen reduces its volume, making it feasible to store larger quantities in smaller, more manageable volumes
40
features of electrolytic cells
* distance between electrodes affect current and thus rate of reaction
* increase in current can increase reaction rate
* powersource must have greater EMF than spontaneous galvanic cell would produce
* surgace area of electrode is proportional to rate
* conc of electrolyte remain constant
41
role of membrane in electrolytic cell
* allows H+ ions / protons to pass through in order to **complete the circuit.**
* The membrane prevented a **spontaneous reaction** between the products.
42
how does the operation of a secondary cell change when it is nearing the end of its cell life
* The cell will not be able to be fully recharged.
* The battery does not last as long between recharges.
43
pruposes of the membrane in the electrolyser
* separation of products that would spontaneously react if they came in contact
* completion of the internal circuit
* prevention of the direct flow of electrons / direct reaction of reactants
44
why don't cells last forever?
● Loss of active materials (reactants and products)
● Formation of other chemicals in side reactions that
prevent the cell from working properly
● Impurities in cell materials, including electrodes
➔ These can react with active materials
● Decrease in contact of electrolyte with electrodes
➔ Either through leakage of electrolyte or its transformation
into a non-conductive material
● Corrosion or failure of internal components
45
self-discharge
● Even when a battery/cell is not in use (anode and cathode not connected), side reactions and deterioration still occur
● These reactions can be slowed by storage at cold temperatures
