ELECTROLYSIS

electrolysis
Electrolysis: General Principles
• When an electric current is passed through a molten
ionic compound the compound decomposes or
breaks down
• The process also occurs for aqueous solutions of
ionic compounds
• Covalent compounds cannot conduct electricity hence
they do not undergo electrolysis
• Ionic compounds in the solid state cannot conduct
electricity either since they have no free ions that
can move and carry the charge
Key terms used in a simple
electrolytic cell
• Electrode is a rod of metal or graphite through which
an electric current flows into or out of an electrolyte
• Electrolyte is the ionic compound in a molten or
dissolved solution that conducts the electricity
• Anode is the positive electrode of an electrolysis cell
• Anion is a negatively charged ion which is attracted to
the anode
• Cathode is the negative electrode of an electrolysis cell
• Cation is a positively charged ion which is attracted to
the cathode
Metals and hydrogen form positively
charged ions and so either
a metal or hydrogen gas is formed at
the cathode
Non-metals form negatively charged
ions and so non-metals (except
hydrogen) are formed at the anode

,
Electrolysis: Charge Transfer
• During electrolysis, current needs to flow around the circuit
• In order for this to occur, charge must be transferred around the circuit
(current is a measure of the rate of flow of charge) by charge carriers
• The power supply provides the cathode with a supply of electrons,
causing it to become negatively charged
• Positive ions (cations) in the electrolyte move towards the cathode
where they gain electrons
• Negative ions (anions) in the electrolyte move towards the anode
where they lose electrons
• The electrons move from the anode back towards the power supply
• So, in a complete circuit:
• Electrons are the charge carriers in the external circuit
• Ions are the charge carriers in the electrolyte
Electrolysis of Molten Compounds
• A binary ionic compound is one consisting of just two elements joined
together by ionic bonding
• When these compounds undergo electrolysis they always produce their
corresponding elements
• To predict the products made at each electrode, first identify the ions
• The positive ion will migrate towards the cathode and
the negative ion will migrate towards the anode
• Therefore, the cathode product will always be the metal, and the
product formed at the anode will always be the non-metal
 Example of lead (II) bromide
 Method:

• Add lead(II) bromide into a beaker and heat it so it will

turn molten, allowing ions to be free to move and
conduct an electric charge
• Add two graphite rods as the electrodes and connect
this to a power pack or battery
• Turn on the power pack or battery and allow electrolysis
to take place
• Negative bromide ions move to the positive electrode
(anode) and each loses one electron to form bromine
molecules. There is bubbling at the anode as
brown bromine gas is given off
Electrolysis of Aqueous Sodium Chloride & Dilute Sulfuric Acid

 Aqueous sodium chloride

• Brine is a concentrated solution of aqueous sodium chloride
• It can be electrolysed using inert electrodes made from platinum or
carbon/graphite
• When electrolysed, it produces bubbles of gas at both electrodes as
chlorine and hydrogen are produced, leaving behind sodium
hydroxide solution
• These substances all have important industrial uses:
• Chlorine is used to make bleach
• Hydrogen is used to make margarine
• Sodium hydroxide is used to make soap and detergents
Electrolysis continued
 Product at the Negative Electrode:
• The H+ ions are discharged at the cathode as they are less
reactive than sodium ions
• The H+ ions gain electrons to form hydrogen gas
 Product at the Positive Electrode:
• The Cl– ions are discharged at the anode
• They lose electrons and chlorine gas forms
• The Na+ and OH– ions remain behind and form the NaOH solution
Dilute sulfuric acid
 Dilute sulfuric acid
• Dilute sulfuric acid can be electrolysed using inert electrodes made from
platinum or carbon/graphite
• Bubbles of gas are seen at both electrodes
 Product at the Negative Electrode
• H+ ions are attracted to the cathode, gain electrons and form hydrogen
gas
 Product at the Positive Electrode
• OH- ions are attracted to the anode, lose electrons and form oxygen
gas and water
 Determining what gas is produced
• If the gas produced at the anode relights a glowing splint dipped into a
sample of the gas then the gas is oxygen
.
If the gas produced at the anode bleaches
damp litmus paper then the gas is chlorine
If the gas produced at the cathode burns with a 'pop' when a
sample is lit with a lighted splint then the gas is hydrogen
Electrolysis of Aqueous Solutions

• Aqueous solutions will always have water present (H2O)

• In the electrolysis of aqueous solutions, the water molecules dissociate producing
H+ and OH– ions:
 H2O ⇌ H+ + OH–
• These ions are also involved in the process and their chemistry must be
considered
• We now have an electrolyte that contains ions from the compound plus ions from
the water
• Which ions get discharged and at which electrode depends on the relative
reactivity of the elements involved
• Concentrated and dilute solutions of the same compound give different products
• For anions, the more concentrated ion will tend to get discharged over a more
dilute ion
Positive Electrode (anode)
• Negatively charged OH– ions and non-metal ions are attracted to
the positive electrode
• If halide ions (Cl-, Br-, I-) and OH- are present then the halide ion
is discharged at the anode, loses electrons and forms a halogen
(chlorine, bromine or iodine)
• If no halide ions are present, then OH- is discharged at the
anode, loses electrons and forms oxygen gas
• In both cases the other negative ion remains in solution
• The concentration of the solution also affects which ion is
discharged:
• If a concentrated halide solution is being electrolysed,
the halogen forms at the anode
• If a dilute halide solution is being electrolysed, oxygen is
formed
Negative Electrode (cathode

• Positively charged H+ and metal ions are attracted to the negative

electrode but only one will gain electrons
• Either hydrogen gas or metal will be produced
• If the metal is above hydrogen in the reactivity series, then
hydrogen will be produced and bubbling will be seen at the cathode
• This is because the ions of the more reactive metal will remain in the
solution, causing the ions of the least reactive metal to be discharged
• Therefore, at the cathode, hydrogen gas will be produced unless the
positive ions from the ionic compound are less reactive than
hydrogen, in which case the metal is produced
Aqueous Copper Sulfate

• Aqueous copper sulfate contains the following ions:

• Cu2+, SO42-, H+ and OH-
 Product at the Cathode
• Cu2+ and H+ will both be attracted to the cathode but the less
reactive ion will be discharged
• In this case, copper is less reactive than hydrogen
• Copper ions are discharged at the cathode, gain electrons and
are reduced to form copper metal
• The half equation for the reaction at the electrode is:
 Cu2+ + 2e- → Cu
Product at the Anode

• SO42- and OH- are both attracted to the anode

• OH- ions lose electrons more readily than SO42-
• OH- lose electrons and are oxidised to form oxygen gas
• The half equation for the reaction at the anode is
 4OH– ⟶ O2 + 2H2O + 4e–
Ionic Half Equations
• In electrochemistry we are mostly concerned with the transfer of electrons,
hence the definitions of oxidation and reduction are applied in terms of
electron loss or gain rather than the addition or removal of oxygen
• Oxidation is when a substance loses electrons and reduction is when a
substance gains electrons
• As the ions come into contact with the electrode, electrons are either lost or
gained and they form neutral substances
• These are then discharged as products at the electrodes
• At the anode, negatively charged ions lose electrons and are thus oxidised
• At the cathode, the positively charged ions gain electrons and are thus
reduced
• Ionic half equations show the oxidation and reduction of the ions involved
• It is important to make sure the charges are balanced
APPLICATION OF ELECTROLYSIS
• Electroplating is a process where the surface of one metal
is coated with a layer of a different metal
• The anode is made from the pure metal you want to coat your
object with
• The cathode is the object to be electroplated
• The electrolyte is an aqueous solution of a soluble salt of the pure
metal at the anode

• Example: coating a strip of iron metal with tin:

.
• At the anode: Tin atoms lose electrons to form tin ions in solution
• At the cathode: Tin ions gain electrons to form tin atoms which
deposit on the strip of iron metal, coating it with a layer of tin
 Uses of electroplating
• Electroplating is done to make metals more resistant to corrosion
or damage
• e.g, chromium and nickel plating
• It is also done to improve the appearance of metals,
• e.g. coating cutlery and jewellery with silver
 Hydrogen Fuel Cells

• A fuel is a substance which releases energy when burned

• Hydrogen is used as a fuel in rocket engines and in fuel cells to
power some cars
• A fuel cell is an electrochemical cell in which a
fuel donates electrons at one electrode and
oxygen gains electrons at the other electrode
• The hydrogen-oxygen fuel cell produces electricity by combining
both elements, releasing energy and water
• The overall equation for the reaction within a hydrogen fuel cell is:
 hydrogen + oxygen → water
.

• The diagram below shows the setup of a hydrogen

fuel cell
• The air entering provides the oxygen
• The fuel entering is hydrogen
• The only chemical product made is water
Advantages & Disadvantages of
Hydrogen Fuel Cells

 Advantages
• They do not produce any pollution: the only product is water
whereas petrol engines produce carbon dioxide, and oxides of
nitrogen
• They 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
disadvantages
• 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
• Hydrogen is often obtained by methods that involve the
combustion of fossil fuels, therefore releasing carbon dioxide
and other pollutants into the atmosphere