Bio-electrochemistry

Question 1

electron transfer chains

Answer 1

overal chemical reaction takes place by a series of electron transfers betwen molecules

in the mitochondria, the electron transfer occurs with an iron ion, Fe3+ + e- –>–< Fe2+

Question 2

reduction potential

Answer 2

a measure of the work dine/used by the reaction
- can be measured using electrode potentials

Question 3

half cell for reduction of oxygen

Answer 3

4e3- + o2 + 4h+ –> 2h2o

Question 4

half cell for oxidation of nadh

Answer 4

nadh –> nad+ +h+ + 2e-

Question 5

during photosynthesis, the electron transport chains pt 1

Answer 5

1. NADH dehydrogenase - nadh –> nad+, 2e- leaving to go to ETC, H + pumped across membrane
2. ubiquinase = UQ, transfers e- within membrane to cytochrome C - the transfer of e- from UQ –> cytochrome BCL releases energy to pump h+ ions across matrix to IMS

Question 6

during photosynthesis, the electron transport chains pt2

Answer 6

cytochrome C is ETC, with fe3+ ion, fe3+ + e- -> fe2+ cc

this transfers electron to cytochrome C oxidase, which is where oxygen accepts electrons from the protein to form h2o
- this releases energy to pump H+ across membrane

Question 7

during photosynthesis, the electron transport chains pt3

Answer 7

at this point, there is a high concentration of H+ ions in IMS, creating an electrochemical gradient

the H + moves down conc gradient via chemiosmosis through atp synthesis, releases energy for condensation reaction of ADP +PI –> ATP

Question 8

electrochemical cells

Answer 8

consists of 2 half cells and measuring the difference between PD via simple redox reaction

zn(s) + cu2+ (aq) –> cu(s) + zn2+(aq)
- the zinc is oxidised losing 2e- and so dissolves
- copper is reduced, gaining 2e- and precipitates

Question 9

drawing electrochemical cells

Answer 9

on the left
- have the beaker with reaction that is oxidising , with metal electrode and solution of metal ions - zinc loses 2e- when oxidising, so electrons flow to wire to copper, zn2+ joins solution

on right
- beaker with reduction reaction, as copper recieves electrons from zinc, cu2+ in solution joins to for cu on electrode

also got a salt bridge to ensure charge stays sameand ions flow to complete circuit

measuring using voltmeter

Question 10

salt bridge

Answer 10

glass tube filled with saturated kcl

Question 11

electrochemical cell notation

Answer 11

Zn(s)|Zn2+(aq)|| cu2+ (aq)|cu(s)

Question 12

types of half cell - 1

Answer 12

metal electrode in a solution of its ions

cu2+ +2e- –> –< cu

Question 13

types of half cell - 2

Answer 13

gas in contact with solution of its ions - hydrogen electrode - uses platinum electrode as inert
2h+ +2e- –>–< h2

Question 14

types of half cell - 3

Answer 14

two different oxidation states of same soluble species - found usually in biological systems

fe3+ e- –>–< fe2+
using pt electrode

Question 15

types of half cell - 4

Answer 15

metal in contact with an insoluble salt

agcl +e- –> ag(s) + cl-

Question 16

standard conditions with electrochemical cells

Answer 16

1M, 1 arm, pure substances, inert pt electode, 298K

Question 17

cell reactions

Answer 17

they are all written as reduction reactions
- the positive E cell voltage has a greater reductive power, so thats the one reducing, the other one generally negative is the oxidising one.

RH = reducing
LH = oxidising

Ecell = RH-LH

Question 18

standard hydrogen electrode

Answer 18

This can be used to work out one half cell
hydrogen electrode is the LH, oxidising one

Question 19

free energy relationship

Answer 19

delta G = -nFEcell

n = number of electrons transferred in half cell

F = faradays constant = 96485C mol-1

Question 20

E cell and free energy

Answer 20

when E cell is positive (reduction) delta G is negative so spontaneous

when E cell is negative (oxidation), delta G is positive so not soon

when E cell = 0, no overal voltage so delta G = 0, battery is dead

Question 21

Nernst equation

Answer 21

the relationship between electrode potential and concentration

allows us to convert between non standard and standard conditions

for a half cell: xOx + ne- –> yRed
E = E^o + RT/nF x ln[ox]x/[red]^y

E = new potential
E^o = standard reduction potential
R = gas constant = 8.3143jk-1mol-1
T = temp in K
n = number of electrons transfered

Question 22

if things are in its standard state

Answer 22

it can be ignored and given the value of 1/ any solid

Question 23

nernst equation for a cell

Answer 23

produced by substracting 2 half cell nernst equations

becomes E = E ^o + RT/nF x ln [RHaq] /[LHaq]

Question 24

for hydrogen ions standard state

Answer 24

[H+] = 10-7

Question 25

how to know when to use nernst or not

Answer 25

if H+ is not involved, then regular way is fine

if H+ is involved,need to use Nernst equation

Question 26

selective transport

Answer 26

across membranes is important as it power much of cellular machinery, using embedded membrane bound proteins to pump ions across membranes

Question 27

cells and membrane potential

Answer 27

cells are always electroneutral, so charge inside and outside cell need to be opposite

this is why concentration gradients work as ion travel across the membrane for this

Question 28

Membrane potential

Answer 28

calculated using H+ ion electrochemical cells
- no hydrogen gas in biological systems, so displaying the transport of protons across the membrane rather than transfer of electrons

H+out –> H+ in

Question 29

example - potential difference across biological membrane

Answer 29

RH E =E^o + RT/nF x ln [H+out}/pH2^1/2

LH E =E^o + RT/nF x ln [H+IN}/pH2^1/2

THEREFORE. E =RT/nF x ln [H+out]/[H+in]

Question 30

E cell and electrical potential

Answer 30

E cell describes a redox potential difference electrical potential
- a neg elec potential indicates that the inside of the cell is negatively charged compared to the outside

Question 31

electrical potential equations

Answer 31

E = RT/nF x ln[H+] out/[H+]in - electrical potential

delta G =RT x ln [H+]in/[H+]out + Fx electrical potential

Question 32

proton motive force

Answer 32

driving force for protons, the transfer of H+ across a membrane produces a free energy change.

promotes the movement of H+ ions across membranes down the electrochemical gradient

For protons, the electronic potential equation simplifies to:
delta G = 2.303 RT (pHout - pH in) + F electrical potential

Question 33

the proton motive force equation

Answer 33

delta G h+ = -F delta P

delta P = proton motive force

Question 34

glass pH electrode

Answer 34

allows to measure changes in cell voltage that are due to changes in membrane potential

the thin glass membrane holds H+ in, this is dipped into a solution with unknown H+ out

the voltage is measured across the thin glass membrane

this works as the membrane has SiO2 - protonatable sites

measures pH of test solutions
given symbol E*

Question 35

Nernst slope

Answer 35

Ecell = E* - kpH
y = c-mx
ph = x
m= k
E* = c

E=const = 2.303RT/nF pH

2.303rt/f TERM = 0.0591V at 298k therefore K = 0.0591

Question 36

generalising this equation to other ions

Answer 36

E = RT/zF x ln [Mz+]out/[Mz+]in - eletrical potential

delta G = RTln[Mz+] in/[Mz+]out + zF electrical potential

z =the charge of the ion
z is -ve for -ve ions, +ve for +ve ions

Question 37

spontaneous electron transfer positive

Answer 37

passed from E more negative to E more