Study Outline Flashcards
Represents what is happening to only one reactant in an overall equation.
Half-reaction.
A process in which electrons are lost.
Oxidation; LEO; Fe -> Fe3+ + 3e-
A process in which electrons are gained.
Reduction; GER; Cu2+ + 2e- -> Cu
A chemical reaction where electrons are transferred between particles.
Redox reaction.
Electron transfer theory.
Total electrons gained = total electrons lost.
Reducing agent; RA
The oxidized substance ie. Fe
Oxidizing agent; OA
The reduced substance ie. Cu2+
The real or apparent charge an atom or ion posses when bonds are falsely assumed ionic.
Oxidation number.
Rules for assigning oxidation numbers.
- Monatomic ions: charge on ion
- H: 1+, except hydrides: 1-
- O: 2-, except peroxides: 1-
- Atoms in elements (Cl2, O2): 0
- All others add up to total charge
Writing half-reaction equations.
- Balance other than O, H
- Balance O, +H2O
- Balance H, + H+
Basic: a) Add OH- = H+
Basic: b) Make H2O, and dump some - Add e- to balance charges.
Half-reaction method of balancing.
- Obtain half reactions
- Balance e-
- Combine, cancel e-, H2O, H+
- Check charge balance
Oxidation number method of balancing.
- Assign oxidation numbers
- Simples oxidation and reduction equations; balance electrons
- Write redox equation
- Balance O, +H2O
- Balance H, + H+
a) OH-
b) Make water - Check charges
Spontaneous reaction.
When the OA is able to take electrons away from the RA.
Redox spontaneity rule.
A redox reaction will be spontaneous if the oxidizing agent is above the reducing agent in the redox table.
Electric cell (galvanic cell).
Device in which spontaneous redox reaction draws electrons through a conductor producing an electric current.
Salt bridge.
Allows ions to move between half-cels to prevent polarization.
Reduction potential, E`r.
Measured voltage of a half-cell connected to the standard.
Standard half-cell.
Hydrogen, with a reduction potential assumed to be 0.
H2(g)/H+(aq)/ With E`r = 0v
Standard cell potential, /\E`.
Can be determined by the summation of the half-cell potentials. /\E`= Ecathode - Eanode.
Hydrogen-oxygen fuel cell.
O2 reduced/H2 oxidized
Aluminum-air fuel cell.
O2 reduced/Al oxidized
Occurance of a non-spontaneous redox reaction driven by external source of electrical energy.
Electrolysis.
Metal spontaneously oxidized by substances in the environment.
Corrosion.
Protective coatings.
Barrier to environment, ex: paint, oil, metallic coatings
Impressed current or sacrificial anode.
Cathodic protection.
Molten salt electrolysis.
Producing sodium metal.
[Na+ + e- -> Na / 2Cl- -> Cl2 + 2e-]
Production of aluminum.
High voltage
[Al3+ + 3e- -> Al / 2O2- -> O2 + 4e-]
Electrorefining.
Production of a pure metal at the cathode using an impure metal at the anode.
Electroplating.
Depositing a layer of metal on to another object which is the cathode of an electrolytic cell.
Faraday’s Law.
The mass of a substance formed at an electrode is directly related to the charge transferred.
9.65x10^4C of charge per every mole of electrons that flows.
Molar charge of electrons, F.
9.65x10^4 C/mol
Moles of electrons transferred, Ne-.
It/F.
