Chapter 20 - Electrochemistry

Oxidation

Loss of electrons, gain of oxygen, loss of hydrogen, increase in oxidation number

Reduction

Gain of electrons, loss of oxygen, gain of hydrogen, decrease in oxidation number

Oxidation-reduction reactions

Split into two half reactions, must be balanced by both mass and number of electrons (if an atom loses one electron, there must be an atom in the same reaction that gains one electron)

Oxidation numbers/states

A hypothetical number/charge assigned to atoms in a compound to indicate the number of electrons gained or lost by that atom

Oxidation rules

1) Oxidation number of pure elements is always zero

2) Oxidation number of ions is equal to the charge of the ion

3) The sum of oxidation numbers in a neutral molecule = 0, an ion = charge of the ion

F oxidation number

Always -1

O oxidation number

-2, except in Peroxides where it is -1

H oxidation number

+1 except with metals where it is -1

Group 1 oxidation numbers

+1

Group 2 oxidation numbers

+2

Group 3 oxidation numbers

+3

Transition metal/post transition metal oxidation numbers

Various charges

Halogen oxidation numbers

-1

Group 6 oxidation number

-2

Group 4 oxidation number

-4

Metals ____ electrons and non metals ____ electrons

Lose, gain

An atom that loses an electron

Is oxidized, therefore it is the reducing agent because it gave its electron to reduce another atom

An atom that gains an electron

Is reduced, therefore it is the oxidizing agent because it took an electron from another atom

Balancing redox reactions under acidic and basic conditions

1) identify who is oxidized and who is reduced

2) divide the rxn into two half reactions

3) for any number of oxygens, add H2O to the opposite side; for every H add H+ to its opposite side

4) balance mass, number of electrons, cancel similar ions on opposite sides or add similar ions on the same side, and write balanced rxn (you now have acidic)

5) for basic, just add OH- according to the number of H+ to the existing H+ and H2O, then cancel any similar molecules and rewrite equation

Electrical current

Flow of electrons/electrical energy, amount of electrical charge that passes a point in a given period of time

Voltaic/galvanic cell

A spontaneous rxn takes place to generate electricity, produces electricity, converts chem energy to electrical energy, oxidation occurs at anode (-), reduction occurs at cathode (+), electrons flow from anode to cathode

Electrolytic cell

Electricity is used to run a non spontaneous rxn, consumes electricity, oxidation occurs at anode, reduction occurs at cathode, anode is connected to (+) end of battery, cathode is connected to (-) end of battery, electrons flow from anode to cathode

Electrodes

Conductors for electricity to enter or leave

Anions are attracted to the _____ and cations are attracted to the _______

Anode, cathode

In electrolytic cells, the anode _____ weight and the cathode _____ weight

Loses, gains

Ecell > 0

Spontaneous (for galvanic cell, Ecell must be positive in order to generate electricity)

Ecell equation

Ecell = Ecathode - Eanode

The larger the E(V)/Ereduction

The stronger the oxidizing agent (tends to be cathode) and the weaker the reducing agent

The smaller the E(V)/Ereduction

The stronger the reducing agent (tends to be anode) and weaker the oxidizing agent

Ereduction values are intensive

DO NOT CHANGE THEM EVEN IF YOU MULTIPLY A RXN BY A FACTOR!!!!

Standard reduction potential

Reduction of H+ to H2 is the reference for measuring cell potential (H2 electrode is set to zero and other cell potentials are relative to standard hydrogen electrode (SHE))

DeltaGrxn < 0, DeltaSuniv > 0, Ecell > 0

Spontaneous reaction

Salt bridge

Maintains neutrality and enables a continuous flow of current in voltaic cell

Dissolving metals in acids

Metals that are stronger reducing agents than H2 will dissolve in acids (look at Ered table provided), metals that have more negative Ered than SHE will dissolve in acids, to decide if a metal will dissolve in an acid, find Ecell by having the metals as the anode

Relationship between Keq, ΔG, E_cell

ΔG°_rxn = -RTlnKeq, ΔG°_rxn = -nFE°_cell, E°_cell = 0.0592/n(logKeq) (n = electrons transferred)

Nernst Equation

E_cell = E°_cell - 0.0592/n(logQ)

Concentration cell

Electricity can be generated when we have two different concentrations of the same electrode. In this case, the anode is the less concentrated solution.

Electrolysis

The process of using electrical current to drive nonspontaneous reaction, can be used to separate compounds into their elements, cations get reduced and anions get oxidized

What happens if there is more than one cation present to be reduced? (Mixture of ions)

The cation with the most positive E°_reduction will be reduced; the anion with the least positive E°_reduction will be oxidized.

Overvoltage

When an electrolysis reaction requires more voltage than E_cell predicts

Electrolysis of pure compounds

The compound must be in molten (liquid) state, electrodes are normally graphite, cations are reduced and anions are oxidized

Electrolysis of aqueous solutions

Possible cathode reactions: reduction of cation to metal, reduction of water to H2; possible anode reactions: oxidation of anion to element, oxidation of H2O to O2; oxidation of electrode: particularly Cu, graphite doesn’t oxidize; Half-reactions that lead to least negative E_cell will occur (unless overvoltage changes the condition)

Stoichiometry of electrolysis

The amount of product made is related to the number of electrons transferred (essentially electrons are a reactant), # of moles of electrons that flow through the cell depends on the current and length of time (1 amp = 1 coulomb of charge/second, 1 mole of e = 96485 coulombs of charge)

Corrosion

The spontaneous gradual oxidation of a metal by oxidizing agents in the environment. A metal must usually be reduced to extract from its ore, but in corrosion, the metal is oxidized back to its more natural state.