Topic 9: Redox Processes Flashcards
oxidation
- loss of electrons
- increase in oxidation no
- gain of oxygen
- loss of hydrogen
reduction
- gain of electrons
- decrease in oxidation no
- loss of oxygen
- gain of hydrogen
oxidation state
apparent charge of an atom in a molecule or ion.
- used to measure electron control or possession relative to the atom in pure element
relation between reducing power and reactivity in metals
more reactive metals are stronger reductants
relation between oxidizing power and reactivity in non-metals
more reactive non-metals are stronger oxidizing agents
redox titration
redox reaction between an oxidizing agent and reducing agent
chemically speaking, what happens in a redox titration?
electrons are transferred from reductant to oxidant
redox titrations: analysis of iron with manganate (VII)
5Fe2+ + MnO4- + 8H+ -> 5Fe3+ + Mn2+ + 4H2O
- uses KMnO4 in acidic soln. as oxidant
- oxidises Fe (II) ions to Fe (III)
- MnO4- is reduced to Mn2+
- colour change: deep purple to colourless
redox titrations: iodine-thiosulfate reaction
2I- (aq) + oxidant -> I2 (aq) + reduced product
- oxidant reacts with excess iodides to form iodine diatomic molecules
- oxidants can be KMnO4, KIO3, K2Cr2O7, NaOCl, etc.
2S2O3 2- (aq) + I2 (aq) -> 2I- (aq) + S4O6 2- (aq)
- the liberated iodine is then titrated with sodium thiosulfate (Na2S2O3)
- using starch as an indicator (NOT added at the start but during titration
- initially forms deep blue colour due to starch, but as I2 is reduced to I-, the blue colour disappears
redox titrations: winkler method (what it is, its function, etc)
- calculates dissolved oxygen content of water
- used to measure degree of pollution
- as oxygen is used by bacteria in decomposition reactions
biological oxygen demand (BOD)
- amount of oxygen used to decompose organic matter in a sample of water over a specified time period
- usually 5 days at a specified temp
- high BOD = high quantity of degradable organic waste = low level of dissolved oxygen
redox titrations: winkler method (process)
- Dissolved oxygen O2 (g) is fixed by the addition of a manganese (II) salt, such as MnSO4. This causes oxidation of Mn (II) to higher oxidation states:
2Mn2+ (aq) + O2 (g) + 4OH- (aq) -> 2MnO2 (s) + 2H2O (l) - Acidified iodide ions (I-) are added and are oxidised by Mn (IV) to I2:
MnO2 (s) + 2I- (aq) + 4H+ (aq) -> Mn2+ (aq) + I2 (aq) + 2H2O (l) - Iodine produced is titrated with sodium thiosulfate:
2S2O3 2- (aq) + I2 (aq) -> 2I- (aq) + S4O6 2- (aq)
redox titrations: winkler method (ratio of O2 : S2O3 2-)
1:4
voltaic cells
- generates electricity from spontaneous redox reactions
- separates two half-reactions into half-cells, allowing e-s to flow between them through an external circuit
2 connected half-cells = 1 voltaic cell
electrode potential
charge separation between the metal and its ions in solution
half cell
- where a half-reaction occurs
- simplest one is made by putting a strip of metal into a solution of its ions
relationship between cell equilibrium and reactivity (of metal)
- more reactive metals are stronger reducing agents
- they have a higher tendency to lose electrons
- the less reactive the metal, the more to the right the equilibrium position will be
voltaic cell connections between half-cells
- external electronic circuit
- salt bridge
external electronic circuit
- connected to the metal electrode in each half-cell
- can have a voltmeter attached to record generated voltage
- electrons flow from anode to cathode through wire
salt bridge
- glass tube/strip of absorptive paper containing an aq soln of ions
- movement of these ions neutralise build-up of charge and maintains potential difference
- anions move in salt bridge from cathode to anode, which opposes the flow of e-s in external circuit
- cations move in salt bridge from anode to cathode
- solution chosen is typically aq NaNO3 or KNO3
what is affected by the difference in reducing strength between electrodes in a voltaic cell?
- direction of electron flow (e-s always flow to the electrode with less reducing power)
- voltage generated (the greater the reducing power difference, the greater the voltage generated)
electromotive force
- potential difference
- generated due to e- flow between half-cells
- depends on difference of reducing power
- also called cell potential/electrode potential
standard hydrogen electrode
reference standard for measuring reducing power of half-cells
standard hydrogen electrode
- modified form of pH electrode
- platinum is used as the conducting metal in the electrode
- as platinum is inert and doesn’t ionise
- it also acts as a catalyst for proton reduction reaction
- the form of platinum used is platinized platinum (the surface of the metal is covered with very finely divided platinum, or platinum black)
- this causes the electrode reaction to occur rapidly
- due to the large surface area helping in the adsorption of hydrogen gas
standard conditions for measurement of standard electrode potential
- concentration: 1 mol/dm3
- pressure: 100 kPa
- only pure substances
- temp: 298 K
- platinum is used as electrode if the half-cell doesn’t include a solid metal
standard electrode potential
- denoted by Eϴ
- potential of a redox system to lose/gain electrons when connected to the SHE (assigned a value of 0 volts) via a salt bridge and high-resistance voltmeter
standard reduction potential
- half-cells exist as equilibria
- they can occur as either oxidation or reduction
- std. electrode potential is always given for the reduction reaction
- so standard electrode potential values are sometimes also called standard reduction potentials
- they don’t depend on total no. of electrons so they’re not affected by stoichiometry
- the more positive the Eϴ value, the more readily it is reduced
determining spontaneity using Eϴ (cell)
- if Eϴ (cell) = positive, the rxn is spontaneous and the reverse is non-spontaneous
- a voltaic cell always runs in the direction that gives a positive Eϴ (cell) value
electrolytic cell
- uses external electrical energy source to bring about a redox rxn that would otherwise be non-spontaneous
- redox rxns occur at electrodes
- oxidation at anode and reduction at cathode
- e- flow from anode to cathode
determining products in electrolytic cells
- Identify all ions in electrolyte and determine which will migrate to which electrode
- Write the half-equation for the rxn at each electrode
- Balance e-s lost and gained and add the two half-equations
- Consider what changes would be observed as a result of the redox processes (e.g. colour change, precipitation of solid, gas discharge, pH change…)
electrolysis of molten salt
- only 1 ion migrates to each electrode
- as ionic compounds have v high m.pts, another compound might be added to lower m.pt
electrolysis of aq. soln.s
- cathode: H2O can be reduced to H2
- anode: H2O can be oxidised to O2
- difficult to predict as there is more than 1 redox rxn possible for each electrode
selective discharge
the discharge of an ion at the electrode in cases where there is more than 1 possible redox rxn
how to predict the products of electrolysis of aq. soln.s
- Eϴ values of the ions
- relative conc.s of ions in electrolyte
- nature of electrode
