Electron transfer reactions

oxidation

• the loss of electrons from a substance

• or the gain of oxygen

reduction

• the gain of electrons

• or the loss of oxygen

oxidising agent

• oxidises another substance by causing it to lose electrons, but gets reduced itself

• oxidation number of oxidising agent decreases

  • change in oxidation number is equal to the number of electrons involved in the half equation

reducing agent

• reduces another substance by causing it to gain electrons, but gets oxidised itself

• oxidation number of reducing agent increases

  • change in oxidation number is equal to the number of electrons involved in the half equation

how to balance a redox equation

• write the unbalanced equation and identify which atoms change in oxidation number

• deduce oxidation number changes

• balance the oxidation number changes

• balance the charges

• balance the atoms

redox titrations

• an oxidising agent is titrated against a reducing agent

• most transition metal ions naturally change colour when changing oxidation state, so indicators rarely needed

ease of oxidation

• relative ease of oxidation increases down the group for group 1 & 2 metals

• reactions become more vigorous

reduction of halogens

• halogens are oxidising agents, as they remove an electron from the metal they react with, and themselves gain an extra electron from the metal

• oxidising power of halogens decreases going down the group

  • they become less reactive

metal reactivity series

• metals in higher reactivity can displace less reactive metals from their compounds in solutions / oxides

• the more reactive metal acts as a reducing agent

acids with reactive metals

• acid + metal —> salt + hydrogen gas

• extent of reaction depends on reactivity of metal and strength of acid

• reactions of acids & metals can be written as ionic equations showing only species that has changed in reaction

primary (voltaic) cells

• convert energy from spontaneous redox reactions to electrical energy

• generate a potential difference known as EMF (E), also called cell potential

• a half cell is a metal in contact with an aqueous solution of its own ions, a primary cell consists of 2 different half cells

• 2 half cells are connected together to enable transfer of electrons to produce energy as electricity

• cells are connected by an external wire and salt bridge

  • salt bridge allows free movement of ions

primary cell example

• zinc metal strip (electrode) dipped into zinc sulphate connected to a copper electrode dipped into copper sulphate

• zinc is more reactive, so electrons flow from the zinc half cell towards the copper half cell

• to keep half cells electrically neutral, ions flow through the salt bridge

  • negative ions flow to the negative half cell (Zn)

  • positive ions flow to the positive half cell (Cu)

• voltage produced by a voltaic cell depends on the relative difference between the 2 metals in the reactivity series

  • bigger diff = higher voltage produced

fuel cells: primary

• electrochemical cell where a fuel donates electrons at one electrode and oxygen gains electrons at the other electrode

• as the fuel enters cell, it becomes oxidised, setting up a voltage within the cell

• e.g hydrogen-oxygen fuel cell

• advantages:

  • water is the only reaction product, so better for environment than other cells

  • no harmful oxides of nitrogen produced

• disadvantages:

  • hydrogen is highly flammable and its storage carries safety hazard

  • very thick pipes needed to store hydrogen, has economic impacts

convention for writing cells

• the half cell undergoing oxidation is placed on the left and half cell undergoing reduction is on right

• shorthand notation:

  • half cell is denoted by metal/metal ions and II to denote the salt bridge

  • e.g Zn/Zn²⁺II Cu²⁺/Cu

secondary (rechargeable) cells

• employ chemical reactions which can be reversed by applying a voltage greater than the cell voltage, causing electrons to push in the opposite direction

• examples:

  • Lead-acid battery

  • nickel-cadmium cell

  • lithium cell

lead-acid battery: secondary

• 6 cells joined together in series, produce total voltage 12 V, used in cars

• lead metal as negative electrode, lead (IV) oxide as positive electrode, sulphuric acid as electrolyte

• oxidation at anode: Pb + SO₄^2- → PbSO₄ + 2e^-

• reduction at cathode: PbO₂ + 4H+ + SO₄^2- + 2e^- —> PbSO₄ + 2H2O

• reverse reaction occurs during charging

• advantages

  • can deliver large amounts of energy fast

• disadvantages:

  • heavy

  • lead and sulphuric acid are polluting

NiCad cells: secondary

• produce 1.2V

• nickel as positive electrode, cadmium hydroxide as negative, potassium hydroxide as electrolyte

• reaction during discharge:

  • 2NiO(OH) + Cd + 2H2O → 2Ni(OH)₂ + Cd(OH)₂

• process is reversed during charging

• advantages:

  • long life

• disadvantages:

  • cadmium is a toxic heavy metal

  • low voltage

electrolytic cells

• ionic compounds conduct electricity when molten/in solution, the current causes ionic compound to split up and form new substances, this is electrolysis

• electrolyte contains positive and negative ions

  • negative ions move to anode(+) and lose electrons by oxidation

  • positive ions move to cathode (-) and gain electrons by reduction

• e.g electrolysis of molten lead bromide

  • Pb2+ ions move to cathode (-) where they are reduced

    • Pb²⁺ + 2e- —> Pb (s)

  • Br- ions move to anode (+) where they are oxidised

    • 2Br^- —> Br₂ + 2e-

redox in cells

• electrochemical cells are either voltaic or electrolytic

  • voltaic cells generate electricity from chemical reactions

    • spontaneous

  • electrolytic cells drive chemical reactions using electrical energy

    • non-spontaneous

• reduction always happens at cathode, and oxidation always at anode, but polarity changes for diff types of cells

oxidation of alcohols

• primary alcohols can be oxidised to form aldehydes which can be further oxidised to form carboxylic acids, using oxidising agents [O]

  • if an aldehyde is not distilled off, further oxidation with excess oxidising agent will oxidise it to carboxylic acid

• secondary alcohols can be oxidised to form ketones only

  • requires sustained heating

• tertiary alcohols do not undergo oxidation

  • bc there has to be hydrogen on the functional group carbon which breaks off to make water

  • only C-C bonds on functional group carbon in tertiary alcohol

distillation

• to make aldehyde from primary alcohol, mixture must be heated

• aldehyde product has lower bpt than alcohol (since it lost the H-bonding) so can be distilled off as soon as it forms

• distillation carried out using a side arm arrangement which acts as an an air condenser

heating under reflux

• for reactions requiring sustained heating

• to prevent loss of volatile reactants, apparatus includes condenser positioned vertically, which returns components back into reaction flask

reduction of carboxylic acids, aldehydes and ketones

• oxidation reactions an be reversed in presence of a suitable reducing agent [H]

  • carboxylic acid —> aldehyde —> primary alcohol

  • ketone —> secondary alcohol

• most common reducing agents:

  • lithium aluminium hydride, in anhydrous acid (stronger, so can reduce carboxylic acids straight to primary alcohols)

  • sodium borohydride, in aqueous solutions

  • both of these produce the nucleophilic hydride ion H^- which reacts with electron-deficient carbon atom of polar carbonyl group

reduction of unsaturated compounds

• reduction of alkene to alkane

  • addition of hydrogen

  • requires temp of 200 C, nickel catalyst, and 1000kpa pressure

• reduction of alkyne to alkene and alkane

  • addition of hydrogen

  • more hydrogen required to make alkane straight from alkyne

• degree of unsaturation can be deduced from the structural formula of a molecule

  • alkenes containing one double bond can be made saturated by adding one mol of H₂, so degree of unsaturation is 1

  • alkynes containing one triple bond can be made saturated by added 2 mol of H₂, so degree of unsaturation is 2

the hydrogen electrode

• the absolute value of a half cell potential cannot be measured, only differences in potential between pairs of half cells can

• hence, the half-cell used as a reference is the standard hydrogen electrode

  • under standard conditions and 1.0 mol dm hydrogen ion conc

  • platinum elecrode used

  • H ⁺ (aq) + e⁻ ⇌ 1/2H₂(g)

  • given the value of E^θ = 0.00 volts

• the more negative the E^θ value for a half cell, the better it can act as a reducing agent

standard cell potential

• also known as the standard emf, is the difference between the standard E^θ (reducton potential) values of 2 half cells

• E^θ_cell = E_right - E_left

  • or E^θ_cell = E^θ_red - E^θ_ox

  • E^θ_cell = E^θ_cathode - E^θ_anode

    • the half cell with the more negative electrode potential will be anode, more positive is cathode

• E^θ_cell is positive for a spontaneous reaction

gibbs energy and standard cell potential

• in electrochemical cells, a spontaneous reaction occurs when the combined half cells produce a positive voltage through the voltmeter, thus:

  • E^θ_cell is positive, reaction is spontaneous

  • E^θ_cell is negative, forward reaction non-spontaneous but reverse is

• therefore: ΔG^θ = -nFE^θ

  • n = number of electrons transferred

  • F = Faraday constant

  • expressed in kJ mol^-1 

• if both ΔG^θ and E^θ are 0, reaction is at equilibrium

electrolysis of aqueous solutions

• when aqueous of solutions of ionic compounds are electrolysed, products are complicated to predict as there are additional ions present from the water

• at cathode, metal ion M+ or water can be reduced

  • 2H₂O + 2e^- —> H₂(g) + 2OH^- (aq)

• at anode, anion A- or water can be oxidised

  • 2H₂O (l) —> 4H⁺(aq) + O₂ (g)+ 4e-

• which species is discharged depends on:

  • relative values of E^θ

  • concentration of ions present

  • identity of electrode

influence of relative E^θ values

• the lower the metal ion in the reactivity series (i.e the more positive its standard reduction potential), the more readily it will be reduced to form a metal at the cathode

• e.g in the electrolysis of a solution of NaOH, hydrogen will be evolved (reduction of water) at the cathode instead of sodium, because it’s lower in the reactivity

• but, in a solution of copper (II) sulfate, copper will be deposited at the cathode in preference to hydrogen

influence of the concentration of ions

• if one ion is more concentrated than another, then it is preferentially discharged

• e.g when electrolysing aqueous NaCl, both oxygen (from oxidation of water) and chlorine evolved at the anode

  • for dilute solutions of NaCl, mainly oxygen evolved

  • for more concentrated solutions of NaCl, more chlorine is evolved

influence of nature of electrode

• if copper electrodes are used during electrolysis of copper sulfate, then the anode itself is oxidised to release electrons and form copper (II) ions

• since copper is simultaneously deposited at the cathode, the concentration of the solution remains constant throughout electrolysis

electroplating

• electroplating involves the electrolytic coating of an object with a very thin metallic layer

• anode is usually made from the same metal to replenish the loss of the metal during electrolysis and maintain a constant concentration

• the object to be electroplated is attached as the cathode

• e.g in copper plating

  • cathode is the object to be plated

  • electrolyte is copper (II) sulfate

  • anode is copper to replenish

• as electricity passes through the solution, the copper anode dissolves in the solution forming Cu²⁺ ions and the Cu²⁺ in the solution are deposited onto the cathode