SAI Redox Introduction

Definition of Redox Reaction §

• Redox reaction is a reaction involving the REDuction and OXidation (REDOX) of two chemical species that involves a transfer for electrons from the reductant to the oxidant AND thus involves a change in the OXIDATION state of the reactants.
• Reduction -> process of gaining electrons -> oxidation state reducing
• Oxidation -> process of losing electrons -> oxidation state increasing
  • OILRIG -> oxidation is losing and reduction is gaining
• reductant -> something that causes another species to be reduced (gain e-)
  • losing e- and being oxidised
• oxidant -> something that causes another species to be oxidised (lose e-)
  • gaining e+ and being reduced

oxidation states §

• number/charge of an atom if its bonds were purely ionic (it existed as a free ion)
• oxidation state also indicate the extent to which an element has been oxidised/reduced.
  • a larger positive oxidation number -> has a lot of electrons.
  • a larger negative oxidation number -> has gained a lot of electrons.
• calculating oxidation states
  • oxidation state of any free element is 0.
  • Monoatomic ions have an oxidation state equal to their charge.
  • oxidation state of Hydrogen in all compounds is +1 except
    • in metal hydrides (-1)
  • oxidation state of oxygen in a compound is always -2 except
    • peroxides (-1)
    • difluorine monoxide (+2)
  • the sum of oxidation states in a compound is equal to its ionic charge
  • the most electronegative element is given the negative OS whereas less EN are given positive
• this is because more EN elements attract electrons more
• e.g. in F2O, F is more EN than O so it gets the -1 OS where as O gets the +2

redox §

• for redox reactions to occur, there needs to be a transfer of electrons.
• when a redox reaction occurs, there is a change in oxidation states.
• oxidant is NOT the one that gets oxidised !!!
• combustion is an oxidisation reaction!

Exception case: no evident transfer !! so silly §

• In some cases, it is not evident if there has been any electron transfer between species. These reactions are quite rare and are almost always synthesis and degradation reactions.

1. Synthesis reactions
   • a reaction where two smaller products (sometimes more) are combined to form a single larger product.
2. Degradation reactions
   • a reaction where a LARGE product is degraded to produce two or more smaller products.

• sublimation of solid carbon $C_{(s)}+O_2(g)\to C{O}_2(g)$ (CO2 is oxidant and reductant)
• transfer of electrons is INTRA molecular and electrons are shared. this means electrons are passed between elements and not species and there is no complete transfer of the electrons: they remain shared between two elements.
• carbon is oxidised and oxygen is reduced.

Half equation and rules §

• A Half reaction is a theoretical (not actual) reaction that represents the process of oxidation and reduction separately. REMEMBER -> a half-reaction cannot occur in isolation. Oxidation MUST be coupled with reduction (simultaneously.)

types of reaction	definiton / example
single displacement	redox reaction where one species displaces another from a molecule
double displacement	a non redox reaction where two ions of an aqueous species displace each other
synthesis	a non-redox/redox reaction where 2 species combine to produce a larger compound
decomposition	a non-redox/redox reaction where a large
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corrosion	a redox reaction involving the oxidation of metal such as Fe in the presence of oxygen or some oxygen containing oxidtant
disproportionate

disproportionation §

• oxygen “returns” their electron and loses the electron they received from hydrogen (for peroxide reaction)

conjugate pairs §

• conjugate redox pair is a term that denotes an electron donor and its corresponding electron acceptor form. For example, Cu+ is a donor and Cu2+ is an acceptor
• an example is given below:
• pair must come from the same half equation.
• one gets oxidised is the reducing agent and the one getting reduced is the oxidising agent.
• pair is called conjugate reducing agent.

standard reduction potential §

• SRP represents the tendency of a substance to reduce (gain e-)
  • relative to the H2/H+ half cell
  • at 25degC, 100kPa or pressure and utilise electrolyte with a 1mol/L concentration
  • measured in votls.
  • standard oxidation potential represents the tendency of a substance to oxidise (lose e-)and is = -SRP
• fluorine has a tendency to reduce at a higher preference.
• srp is relative.
• srp ranks the different metals, ions and chemicals species according to the tendency to reduce.
• three major limitations MUST MEMORISE
  • only given at 25degC, 100kPa, and 1mol/L concentrations - changing the conditions will change the ranking.
  • only applies in aqueous solutions
    • a lot of organic reactions etc. occur in liquid and other mixtures while involving redox reactions so SRP is useless in comparing redox strength in these cases.
  • does not reflect the rate of a reaction
    • a reaction may theoretically occur but in reality it may have such a high activation energy that its progress is negligible.
• negative srp
  • a negative srp means that relative to H+, the reactant species will not reduce.
  • A larger and negative SRP means that the forwards reduction reaction is very unlikely relative to the reaction $2H^{+}+2\overline{e}\to H_2$. that the reverse oxidation reaction is very likely.
• classifications
  • thus elements on the top left hand side of the SRP table are easily reduced (strong oxidants)
  • thus elements on the bottom right hand side of the SRP table are easily oxidised (strong reductants)
• sacrificial anode: strong reducing agent (more reactive) metal that will be oxidised preferentially attached to iron (or other metals) to protect it (from oxidising)