Electrochemistry

3 fundamental electrochemical cells

1. Galvanic (voltaic)

2. Electrolytic

3. Concentration

Anode

Electrode where oxidation takes place

negative electrode

cathode

electrode where reduction takes place

positive electrode

Electromotive force (emf)

voltage or electrical potential difference of a cell

emf is positive = release energy (G<0)

emf is negative = absorb energy (G>0)

Movement of electrons in a cell

opposite to the flow of the current

e^- = anode to cathode

I = cathode to anode

Galvanic Cells

Spontaneous reactions, G < 0

electrons flow from anode to cathode

anions and cations flow via the salt bridge

Cell diagram construction

anode I anode solution II cathode I cathode solution

Electrolytic cells

nonspontaneous reactions that require the input of energy

G > 0

anode is positive

cathode is negative

Faraday’s Constant

10⁵ (C/mol e^-)

1 faraday (F) is equivalent to the amount of change contained in one mole of electrons

Electrodeposition equation

mol M = It / nF

Concentration Cells

Spontaneous redox reactions

type of galvanic cell, electrodes are chemically identical

current is derived from a concentration gradient

Reduction potential

ability for a species to gain electrons and become reduced

the more positive the value, the more likely it is to be reduced

Free energy equation

G = -nFE_cell

1. Split equation into half reactions

2. determine cathode and anode

3. determine emf (E_cell)

4. use equation

Nerst equation for reaction quotients

E_cell = E^o_cell - (RT/nF) lnQ