Electrode Potentials Flashcards
Writing redox reaction
Calculate oxidation states of each element present in equation
Use this to determine what is ox/reduced
Balance O with H2O and then H+
Add electrons to side to balance charges/ox states
Manganate redox titrations
End point - permanent pink colour shows excess of MnO4- ions
Oxidising agent
Commonly to analyse Fe2+ ions and ethanedioic acid
Electrochemical cells
Converts chemical to electrical energy
Require reaction when one species transfers electrons to another species
Half cells
Chemicals present in redox half equation
Contains an electrode and an aqueous solution
Metal metal ion half cells
Metal electrode in a solution of its aqueous metal ions
Half cell equation is in equilibrium
Forward is reduction vice versa
Ion ion half cells
Contains same elements in different oxidation states
Eg aq Fe2+ & Fe3+
Inert metal electrode present
Electrode potentials
More reactive metal releases electrons more readily and is oxidised -ve electrode
Less reactive gains electrons and is reduced positive
Standard electrode potential
The emf of a half cell connected to a standard hydrogen half cell in standard conditions
Standard electrode (H) is exactly 0V
More negative the electrode potential value
Greater tendency to lose electrons and undergo oxidation (reverse equation)
Greater reactivity of a metal
More positive the Electrode pot value
Greater tendency to gain electrons - reduction (forward reaction)
Greater reactivity of non metal
Electron flow
From most negative to most positive
Calculating cell potential
Positive electrode - negative electrode
Large +ve electrode potential
Half equation has strong tendency to go from left to right
Predicting feasibility of equations
Write out half equations starting w most negative circle species reacting
EP left to right - EP right to left( don’t reverse sign) = EP reaction
If EP value is positive reaction will take place
Limitations predicting reaction using EP
Reaction has a high activation energy therefore the rate is very slow
The reaction may not happen under standard conditions
EP values apply to aqueous equilibrium - many reactions don’t take place in aqueous conditions
Primary cells
Non rechargeable - one use
Electrical energy is produced by oxidation and reduction at electrodes
Chemicals will be all used up
Alkaline based Zn/MnO2 and KOH electrolyte
Secondary cells
Rechargeable - cell reaction can be reversed
Chemicals in cells regenerated
Car batteries lithium ion batteries - laptops
Fuel cells
Energy from reaction of a fuel with O to create a voltage
Fuel and oxygen into the cell and products flow out electrolyte remains
Do not have to be recharged
Alkali hydrogen fuel cell
Hydrogen and oxygen in
H2 + 2OH- -> 2H2O+ 2e-
1/2 O2 + H2O + 2e- -> 2OH-
H2 + 1/2O2-> H2O
1.23v
Acid hydrogen fuel cell
H2 dissociates
O2 reacts with H+ forming water
1.23v
Manganate equation
MnO4- + 8H+ => Mn2+ + 4 H2O
