• loss of electron of one species (oxidized) and the gain of electrons of another (reduced)
• responsible indirectly for all work done by living organisims
• electrons move from various metabolic intermediated to specialized electron carriers who in turn donate them to acceptors with a higher affinity
• this releases energy which can be used for work

• consider a circuit with two chemical species with different electron affinities
• because of this affinity difference, electrons flow spontaneously through the circuit driven by a force proportional to the difference in electron affinity called the emf
• glucose is the electron source in a biological circuit and as it is oxidized it releases its electrons to carriers with higher affinity and to O2

Oxidation Reduction reactions can be described as half reactions

• redox reactions occur together but can be separated as 1/2 reactions:

Fe(2+) + Cu (2+) <--> Fe(3+) + Cu(+)

Fe(2+) <-->Fe(3+) + e(-)

Cu(2+) + e(-) <-->Cu(+)

• electron donating agent is called the reductant, other is oxidant
• fe(2+) and 3+ are conjugate redox pair

• Oxidation states of carbons based on the affinities of H<C<S<N<O
• in bio systems, the loss of electons, oxidation, is coincident with loss of hydrogen (b/c carbon can only get electrons from hydrogen)
• oxidation is therfore synonymous with dehydrogenation

1. directly as electrons. in
   

   Fe(2+) + Cu (2+) <--> Fe(3+) + Cu(+) the Fe pair donates electrons to the Cu pair

2. as hydrogen atoms. hydrogen atom  is a proton and H(+) and one electron
3. as a hydride ion (:H-) which has 2 electrons
4. through direct combination with oxygen. Oxygen is covalently inducted into the product and is considered the electron acceptor

reducing equivalent is pair a of electrons as a unit

• When two conjugate redox pairs are in solution together, electron xfer may proceed spontaneously
• tendency for such a reaction to happen depends on the relative affinity of the electron acceptor of each redox pair
• the standard reduction potential, E^o, a measure (volts) can be determined experimentally for each half reaction
• H(+) + e(-) --> 1/2H2 is standard of reference and is assigned 0.0v
• when this hydrogen electrode is connected through a circuit to another half cell with another redox pair in it, electrons flow through the exernal circuit from the 1/2 cell with lower to higher  E^o.
• a 1/2 cell that takes electrons from hydrogen cell is positive and one that loses electrons is negative
• that with the larger E will be reduced
• theses are at standared concentrations at std conditions

• tells us that reduction potential is not affected by chemical species but by relative donor/acceptor concentrations
• nernst equation relates standared reduction potential with actual reduction potential at any redox pair concentrations:
• E=E^o +(RT/nFaraday)ln[electron acceptor]/[electron donor]
• at 298K this is now
• E=E^o + (0.026V/n)ln[electron acceptor]/[electron donor]

• energy made avialable by this electron flow is proportional to delta E:
• delta G=-n*faraday*delta E or standard delta G and Standard delta E
• n is number of electrons transferred in the reaction
• can calculate the actual free energy change for any redox reaction from the E values in a reduction table