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Electrochemistry

Chemistry Notes

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Electrical devices and circuit symbols

  • Cell: a source of electricity
  • Battery: a collection of cells
  • Switch: used to stop the flow of current
  • Voltmeter: used to measure voltage
  • Ammeter: used to measure current
  • Resistor: ensures that a suitable amount of current is flowing
  • Variable resistor: a resistor whose rating can be changed
  • Electrodes: plates which carry electricity into the liquid
  • Cathode: electrode connected to the negative terminal of the battery
  • Anode: electrode connected to the positive terminal of the battery
  • Cation: positively charged ion that travels to the cathode
  • Anion: negatively charged ion that travels to the anode

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Electrical devices and circuit symbols

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Electrodes

  • Usually made of materials such as graphite, platinum, copper and sometimes mercury

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Conductors and non-conductors

  • A conductor is a substance which conducts electricity but is not chemically changed during the conduction
    • Conduction of electricity through the metal is possible because of the freely moving valence electrons within the metal’s structure
  • A non-conductor is a substance which does not allow the passage of electricity
    • Sometimes they are used to protect something from electricity – insulators
    • Solid substances which act as non-conductors have their valence electrons in fixed positions so they are not free to move
    • All non-metals except graphite are non-conductors

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Electrolytes, weak-electrolyte and non-electrolyte

  • Electrolytes: compounds which when molten or dissolved in water conduct electricity
    • All acids, alkalis and salts are electrolytes
    • They conduct electricity by the movement of ions between electrodes
    • Solid ionic substances do not conduct electricity as their ions are held together in fixed positions by strong forces
  • Weak electrolytes
    • Weak acids (vinegar) and weak alkalis (ammonia solution)
    • Conduct only few ions
  • Non-electrolytes: a liquid which does not allow the passage of electricity
    • Distilled water, alcohol, turpentine, oil, paraffin and other organic solvents
    • These are covalent substances which do not contain ions

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Electrolytes, weak-electrolyte and non-electrolyte

Non-electrolytes

Weak electrolytes

Strong electrolytes

Organic liquids or solutions

Weak acids and alkalis

Strong acids, alkalis and salt solutions

Ethanol, C2H5OH (l)

Limewater, Ca(OH)2 (aq)

Aqueous sulphuric acid, H2SO4 (aq)

Tetrachloromethane, CCl4 (l)

Ammonia solution, NH3 (aq)

Aqueous nitric acid, HNO3 (aq)

Trichloromethane, CHCl3 (l)

Aqueous ethanoic acid CH3COOH (aq)

Aqueous hydrochloric acid, HCl (aq)

Pure water, H20 (l)

Aqueous sulphurous acid, H2SO3 (aq)

Aqueous potassium hydroxide, KOH (aq)

Sugar solution, C12H22O11 (aq)

Aqueous carbonic acid, H2CO3 (aq)

Aqueous sodium hydroxide, NaOH (aq)

Molten sulphur, S (l)

Copper(II) sulphate solution, CuSO4 (aq)

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Distinguishing electrolytes

  • The brightness of the circuit bulb indicates the degree of ionisation
  • Accordingly, with a strong electrolyte, the bulb shines brightly
  • With weak electrolytes it is dim
  • With non-electrolytes it does not light up at all

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Electrolysis: decomposing using passage of electricity

  • When electricity is passed through an electrolyte, chemical decomposition occurs
  • This involves splitting up of the electrolyte
  • All electrolytes are ionic, which means they are composed of positively and negatively charged ions
  • On passing an electric current through the electrolyte, these ions migrate towards the oppositely charged electrode
  • Negatively charged ions move towards the positive anode and are therefore called anions
  • Positively charged ions move towards the negative cathode and are called cations

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Electrolysis

  • Most anions are non-metal ions, such as oxide (O2-), chloride (Cl-), bromide (Br-)
  • Cations are metal ions, such as copper (Cu2+), silver (Ag+), lead (Pb2+), hydrogen (H+)

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Electrolysis

  • At the anode
    • negative ions lose their electron(s) to the anode, which is very ready to accept electrons because it is positively charge - This means it has a lack of electrons
  • At the cathode
    • positive ions gain electron(s) from the cathode, which has an excess of electrons and therefore an overall negative charge – this release of ions at the electrode results in the chemical decomposition of the electrolyte
    • It also allows electrons to travel from the cathode to the anode
  • Therefore movement of ions during electrolysis allows conduction of electricity

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Electrolysis of Lead (II) Bromide (molten only)

  • This is made up of lead(II) ions, Pb2+ and bromide ions Br-
  • Chemical formula: PbBr2

At the anode

  • When electricity is flowing, brown fumes of bromine gas are seen as the anode
  • Br- -> Br + e-
  • Br + Br -> Br2 (g)

At the cathode

  • When electricity is flowing, a silvery deposit of lead metal forms on the cathode
  • Pb2+ + 2e- -> Pb (l) (molten lead deposited)
  • The lead(II) ions are positive so they move to the negative cathode, where each ion gains 2 electrons to from a lead atom
  • Any reaction at a cathode involved is a gain in electrons (reduction)
  • PbBr2 (l) -> Pb (l) + Br2 (g)

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Electrolysis of concentrated sodium chloride solution (brine)

  • The electrolytic cell used for electrolysis of concentrated sodium chloride solution is designed to collect gaseous products at both electrodes
  • The cathode can be platinum or carbon but the anode must be carbon to resist attack by chlorine
  • Ions present
    • From sodium chloride: Na+, Cl-
    • From water: H+, OH-

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Electrolysis of concentrated sodium chloride solution (brine)

  • At the cathode
    • The sodium and hydrogen ions move to the cathode
    • As the hydrogen ion is lower in the reactivity series than the sodium ions, it accepts electrons more easily
    • The hydrogen ions are discharged
      • H+ (aq) + e- -> H
    • Hydrogen atoms join in pairs to give molecules
      • H + H -> H2 (g)
  • At the anode
    • Both the chloride ions and hydrogen ions migrate to the anode
    • The chloride ions are preferentially discharged because of their higher concentration
      • Cl- (aq) – e- -> Cl
    • Chlorine atoms join in pairs to give molecules
      • Cl + Cl -> Cl2 (g)

As the hydrogen ions and chloride ions are discharged, sodium ions and hydroxide ions remain in the solution -> becomes sodium hydroxide

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Electrolysis of aqueous sulphuric acid

  • As sulphuric acid is aqueous, it is composed not only of hydrogen ions and sulphate ions, but also hydroxide ions from the water
  • When we have more than one type of ion moving to an electrode, selective discharge takes place
  • This means that the ion which can lose or gain electrons with the greatest ease is discharged, and the other ions, which are harder to discharge, remain in solution
  • At anode
    • Hydroxide is easier to discharge (compared to sulphate ions) so oxygen gas is given off at the anode
    • 4OH- (aq) -> O2 (g) + 2H2O (l) + 4e-
  • At cathode
    • Each hydrogen ion gains an electron to become a hydrogen atom
    • Two of these newly formed atoms then combine to form a hydrogen gas molecule
    • H+ (aq) + e- -> H
    • H + H -> H2 (g)

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Electrolysis of aqueous sulphuric acid

  • With electrolysis of aqueous solutions of dilute acids of alkalis, the volume of hydrogen given off at the cathode is roughly twice that of the oxygen gas at the anode
  • Accordingly, the elements of water are lost and as the electrolysis continues, the concentration of the acid or alkali increases
  • Essentially, the electrolysis of aqueous sulphuric acid is the electrolysis of water, with hydrogen and oxygen gas being given off in a ration 2:1

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Preferential discharge

Cations

Anions

K+ (aq)

SO42- (aq)

Na+ (aq)

NO3- (aq)

Ca2+ (aq)

Cl- (aq)

Mg2+ (aq)

Br- (aq)

Zn2+ (aq)

I- (aq)

Fe2 (aq)

OH- (aq)

Pb2+ (aq)

H+ (aq)

Cu2+ (aq)

Ag+ (aq)

Difficulty of discharge decreases

  • The ions at the top require large amounts of energy to be discharged
  • Down the table, they become easier to discharge

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Factors affecting electrolysis

  • 2 main factors
    • Concentration
    • Type of electrode

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Concentration

Anode

Cathode

Dilute hydrochloric acid

Oxygen gas

Hydrogen gas

Concentrated hydrochloric acid

Chlorine gas

Hydrogen gas

Concentration

  • If the concentration of a particular ion is high, then this can alter the preferential discharge
  • Eg if dilute hydrochloric acid is electrolysed, hydrogen gas is given off at the cathode and oxygen gas at the anode
  • However, when hydrochloric acid is electrolysed, hydrogen gas is still given off at the cathode, but chlorine gas comes off at the anode
  • This is because although the chloride ion is harder to discharge than the hydroxide ion, its high concentration makes it more likely to be given off
  • Cl- (aq) -> Cl + e-
  • Cl + Cl -> Cl2 (g)

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Type of electrode

  • Eg: electrolysis of aqueous copper(II) sulphate solution
  • If we use carbon electrodes, they are inert electrodes and do not affect the electrolysis
  • Therefore, at the anode, we have a choice of sulphate or hydroxide ions
  • The hydroxide ions are easier to discharge, so oxygen gas is given at the anode
    • 4OH- (aq) -> O2 (q) + 2H2O (l) + 4e-
  • At the cathode, we have a choice of copper or hydrogen ions
  • The copper ions are easier to discharge, so we see a pink deposit of copper metal on the carbon electrode
    • Cu2+ (aq) + 2e- -> Cu (s)

Anode

Cathode

Copper(II) sulphate with carbon electrode

Oxygen gas

Copper deposited

Copper(II) sulphate with copper electrode

Copper anode dissolves

Copper deposited

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Industrial applications: purification of copper

  • Industrial purifications
    • Extraction of metals from ores, esp. aluminium
    • Purification of metals, esp copper
  • Refining or purification of copper
    • The impure copper is made the anode and a thin, pure copper plate is used as a cathode
    • The electrolyte is usually acidified copper(II) sulphate solution
    • When electricity flows, the copper dissolves from the impure anode and goes into solution as copper ions
    • Impurities in the copper do not dissolve, and instead fall off the anode as an anode sludge
    • At the cathode, the copper ions are deposited as pure copper metal
    • A major use of copper is in electrical wiring

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Industrial applications: manufacture of sodium hydroxide

  • Produced by electrolysis of concentrated seawater
  • The chlorine and hydrogen gases which are also produced during this electrolysis are both commercially useful

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Industrial applications: electroplating

  • To form a thin protective coating of a metal on the surface of another which is likely to corrode
  • Eg: galvanising of iron - Iron is electroplated with a layer of zinc
  • Also, stainless steel is electroplated first with nickel and then with a very thin layer of the more expensive but more attractive metal chromium
  • In general, to electroplate, we need to make the cathode the object for plating
  • The anode is then made of the metal we wish to plate with
  • The electrolyte needs to be a solution of a salt of this metal
  • Anode: Cu (s) -> Cu2+ (aq) + 2e- (copper dissolves from the anode)
  • Cathode: Cu2+ (aq) +2e- -> Cu (s) (Copper deposits on the object)

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Industrial applications: electroplating

Metal

Object

Zinc (galvanise)

Dustbins, buckets

Chromium

Car bumpers, bicycle handle bars

Silver

Watches, bracelets

Copper

Saucepans

Nickel (EPNS)

Cutlery (electroplated nickel steel_

Gold

Jewellery, watches

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Cells and batteries

  • A device which converts chemical energy into electrical energy is called a cell, or battery
  • It consists of a pair of dissimilar metals in an electrolyte
  • The more reactive metal, magnesium, dissolves in the dilute acid and becomes magnesium ions, thereby producing electrons
  • As electrons are produced, the magnesium acts as the negative electrode
  • Mg (s) -> Mg2+ (aq) + 2e-
  • These electrons then travel to the copper electrode and produce bubbles of hydrogen gas at this electrode.
  • As electrons are taken in, the copper is the positive electrode
  • 2H+ + 2e- -> H2 (g)

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Cells and batteries

  • This production and movement of electrons is electricity, so electrical energy has been generated and the bulb lights up
  • Overall, the chemical reaction can be represented by the ionic equation
  • Mg (s) + 2H+ (aq) -> Mg2+ (aq) + H2 (g)
  • How bright the bulb is, depends on the difference in the reactivities of the two metals
  • Magnesium and copper are far apart in the reactivity series, so the bulb is bright
  • However, as the magnesium dissolves rapidly in the acid, the bulb does not stay alight for long
  • Also, the copper electrode gets covered with bubbles of hydrogen gas, which prevent the electricity from flowing
  • This is called polarisation

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Reactivity series

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Dry cells

  • If the electrolyte in the cells is a paste, it is called a dry cell
  • Eg in battery operated clocks, torches, games, shavers etc

  • It consists of a carbon rod which is inside a porous container of powdered carbon and manganese(IV) oxide
  • The electrolyte is a paste of ammonium chloride
  • The other electrode is the zinc casing of the cell itself
    • This zinc casing acts as the negative terminal as it produces electrons
  • Zinc -> zinc ions + electrons
  • Zn -> Zn2+ + 2e-

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Dry cells

  • Electrolyte: ammonium chloride paste
  • Positive terminal: carbon electrode
  • Negative terminal: zinc electrode

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Dry cells

  • The positive terminal is the carbon rod, whose surface is increased by surrounding it with powdered carbon
  • Also present is manganese(IV) oxide, which acts as a depolarising mixture, helping to oxidise the hydrogen gas produced when the ammonium ions from the electrolyte take up the electrons produced from the zinc electrode
  • Ammonium ions + electrons -> ammonia + hydrogen
  • 2NH4+ (aq) + 2e- -> 2NH3 (g) + H2 (g)