Electrochemistry Flashcards
Define electrochemistry
the interconversion of chemical and electrical energy
- spontaneous reactions can produce electricity and electricity can cause non-spontaneous reactions to occur
redox reactions involve
the movement of electrons from one reagent to another
oxidation
loss of electrons
reduction
gain of electrons
oxidising agent
- reduced
- takes electrons from the substance being oxidised
reducing agent
- oxidised
- gives electrons away
how to balance redox reactions
- write down the two half-reactions
- balance the atoms and charges in each half-reaction
- first balance atoms other than O and H, then O, then H
- charge is b balanced by adding electrons (e-) to the reactant side of the reduction half-reaction and to the product side in the oxidation half-reaction
- electrons should cancel out in net reaction. if necessary, multiply one or both half-reactions by an integer so that number of e- gained in reduction = number of electrons lost in oxidation
- add the balanced half-reactions and include states of matter
two types of electrochemical cells
- voltaic cells
- electrolytic cells
voltaic/galvanic cells
spontaneous chemical reaction (ΔG<0) generates an electric current
- batteries contain one or more voltaic cells
- voltaic cell does work on the surroundings, converting higher energy reactants in the cell into lower energy products
electrolytic cells
non-spontaneous reactions (ΔG>0) are driven by electric current
- external power source supplies free energy to run electrolytic cells. the surroundings thus do work on the cell. lower energy reactants are converted to higher energy products in the cell
- used for electroplating, purification of metals, and more
electric current is
flow of electrons
electrons only flow if
the driving force (free energy change) is large enough
electrodes
usually metal strips/wires connected by an electrically conducting wire
anode
electrode where oxidation takes place
cathode
electrode where reduction takes place
describe salt bridge
U-shaped tube that contains a gel permeated with a solution of inert electrolyte (contains positive and negative spectator ions).
function of salt bridge
The salt bridge keeps half cells electrically neutral because ions flow in and out of the salt bridge, counteracting charge build-up due to electron flow.
conventions for notation for a voltaic cells
- anode components are written on the left
- cathode components are written on the right
- components of each half cell are written in the same order as in their half reactions
- single line shows a phase boundary between the components of a half cell
- double line shows that the half cells are physically separated
addition notation for voltaic cell that is more complex
any inactive (inert) electrode is specified
a comma is used to show components that are in the same phase
why does a voltaic cell work?
- differing abilities of metals to gain electrons gives rise to a voltage drop
- this is also known as electromotive force (EMF) or cell potential
multimeter
can measure voltage
cell potential (Ecell)
the difference in electrical potential between two electrodes
units for Ecell
measured in Volts (V), where
V = J/C (Colomb is SI unit of charge)
difference in Ecell for voltaic and electrolytic cells
Voltaic: Ecell>0 (spontaneous e- flow)
electrolytic: Ecell<0 (non-spontaneous e-flow)
when Ecell = 0,
the redox reaction has reached equilibrium so the cell can do no more work
define standard potential (E’cell)
cell potential under standard state conditions
how to calculate Ecell from Ehalf-cell
the cell potential of any electrochemical cell is the sum of the half cell potentials for the oxidation and reduction half cells
Ecell = E ox + E red
what does it mean to say that E is an intensive property?
Bigger cells with more moles of redox components will last longer, but will have the same standard output voltage
which half cell will form the oxidation half cell?
the lower Ered (ie more negative/less positive = worse at being reduced)
use of standard hydrogen electrode
- potentials are determined experimentally from the difference in potential between two electrodes
- the reference point is called the standard hydrogen electrode
Standard Hydrogen Electrode
consists of a platinum electrode in contact with H2 gas (1atm) and aqueous H+ ions (1M)
standard hydrogen electrode is assigned (arbitrarily) a value of
exactly 0.00 V
how is Ecelll related to ΔG?
ΔG = -nFE(cell) = -RTln(Keq)
- ΔG is in J/mol
- n is in mol and is the no of moles of electrons transferred per mole of the reaction
- F is faraday’s constant
at equilibrium, ΔG = and Ecell =
0
Nernst Equation
E = E’ - RT/nF (lnQ)
E = E’ - 0.0592V/n logQ
E - The cell potential (electromotive force, EMF) under non-standard conditions, measured in volts (V).
The standard cell potential, which is the voltage of the electrochemical cell under standard conditions
how does cell potential depend on the relative concentration of reactants and products?
when Q<1, lnQ<0 so Ecell>E’cell
when Q=1, lnQ=0 so Ecell = E’ cell
when Q>1, lnQ>1 so Ecell<E’cell
two consequences of cell potential depending on the relative concentration of reactants and products:
- as a cell is operated, concentration of reactants will decrease and products will increase. thus, cell potential will decrease over time
- concentration cells - capture the electrical energy from a concentration difference
how is a concentration cell created
- has the same half-reaction in both cell compartments, but with different concentrations of electrolyte
- there is a potential difference between cells, which drives current flow until both compartments have an equal concentration of ions
- Ecell>0 as long as the half-cell concentrations are different
- once the concentrations equalise, Ecell = 0 and current stops flowing
write equations for corrosion
how to protect against corrosion
Galvanisation: coating of iron with zinc.
Zinc is more easily oxidised than iron, giving up electrons to it. Iron is more easily reduced than zinc.
