Redox

Question 1

Electrochemistry

Answer 1

The study of the exchange between electrical and chemical energy

Question 2

Battery

Answer 2

A galvanic cell or cells connected in a series with a constant amount of reagents. A battery stores energy in the form of electrical potential energy.

Question 3

Electrorefining

Answer 3

Process by which materials, usually metals, are purified by means of an electrolytic cell. The anode is the impure metal and the cathode is the pure sample of the metal.

Question 4

Oxidation

Answer 4

The loss of an electron from a species (an increase in its oxidation number).

Question 5

Oxidation Number

Answer 5

A conceptual bookkeeping numbering system that allows us to track the number electrons transferred during a redox reaction.

Question 6

Reduction

Answer 6

The gain of an electron by a species (and a decrease in the oxidation number)

Question 7

Redox

Answer 7

A reaction involving the transfer of one or more electrons from the reducing agent to the oxidizing agent.

Question 8

Reducing Agent

Answer 8

A reactant in a redox equation that donates an electron to the reduced species. The reducing agent is oxidized.

Question 9

Oxidizing Agent

Answer 9

A reactant in a redox reaction that accepts an electron from the oxidized species. The oxidizing Agent is reduced.

Question 10

Galvanic Cell

Answer 10

An electrochemical cell with a positive cell potential that allows chemical energy to be converted into electrical energy.

Question 11

Cell Potential

Answer 11

The overall electrical potential of an electrochemical cell. It is the sum of the reduction potential of the cathode and the oxidation potential of the anode.

Question 12

Oxidation Potential

Answer 12

The potential of a half-reaction written as an oxidation reaction, it is the opposite sign of the same reaction written as a reduction.

Question 13

Reduction Potential

Answer 13

Arbitrarily setting the potential of the standard hydrogen electrode, SHE, to zero, all other half reactions are measured by their power to reduce hydrogen. The voltage given by the construction of a galvanic cell between the SHE and the reduction of interest gives the standard reduction potential of that reduction.

Question 14

Work

Answer 14

Force over a distance

Question 15

Electrolytic Cell

Answer 15

A cell that consumes electrical energy to drive a non-spontaneous redox reaction.

Question 16

Anode

Answer 16

The electrode that is the source of the negative charge, designated by a minus sign (-); this electrode is the site of oxidation.

Question 17

Cathode

Answer 17

The electrode that is the source of positive charge, designated by a plus sign (+); this electrode is the site of reduction.

Question 18

Concentration Cell

Answer 18

A galvanic cell that has two compositionally equivalent half-cells of differing concentrations. One can calculate the potential developed by such a cell using the Nernst Equation.

Question 19

Electrode

Answer 19

A conducting material placed in physical contact with a half-reaction on which the electron transfers in the redox reaction take place.

Question 20

Fuel Cell

Answer 20

A galvanic cell with a constant flow of reagents in and products out used for the prodction of a constant supply of energy. Whereas batteries have a finite lifetime of useful energy production, fuel cells are only limited in their duration of energy production by the ammount of available fuel reagents.

Question 21

Half-Cell

Answer 21

A half-reaction and its electrode; it is half of a galvanic cell.

Question 22

Half-Reaction

Answer 22

Either an oxidation or a reduction reaction, it represents half of the redox reaction.

Question 23

Half-Reaction Method

Answer 23

The set of rules that have been developed to aid in balancing redox reactions.

Question 24

Line Notation

Answer 24

A shorthand way of describing an electrochemical cell without drawing a picture.

Question 25

pH Meter

Answer 25

A concentration cell that places a known concentration of acid in the meter’s electrode. When that probe is dipped into a solution of unknown acid concentration, a potential develops due to the differences in concentration which can be calculated by the Nernst Equation. From that potential, the meter calculates the concentration of the unknown acid and, therefore, its pH.

Question 26

Porous Disk

Answer 26

A disk placed between two half-reactions allowing ion flow between half-reactions without mixing the reactions. The mixing of the two cells would cause a short in the circuit.

Question 27

Potentiometer

Answer 27

A device that measures electrical potential.

Question 28

Reaction Quotient

Answer 28

Similar to the form of an equilibrium constant, the reaction quotient is the ratio of the product of the each product in a reaction raised to its stoichiometric power divided by the product of each reactant raised to its stoichiometric power from the balanced equation.

Question 29

Salt Bridge

Answer 29

A tube plugged with pourous material at either end (usually cotton) filled with a gel allowing ion flow between half-reactions without mixing the reactions. The actual mixing of the two cells would cause a short in the circuit.

Question 30

Standard State

Answer 30

An arbitrarily defined set of conditions–273K, 1atm for gasses, or 1M for solutions.

Question 31

Atoms in elemental form have oxidation states of zero.

Answer 31

N/A

Question 32

The charge on a monoatomic ion is equivalent to its charge.

Answer 32

N/A

Question 33

Hydrogen has an oxidation state of +1 when bonded to non-metals and -1 when bonded to metals.

Answer 33

N/A

Question 34

F, because it only forms one bond and is the most electronegative element, has an oxidation state of -1.

Answer 34

N/A

Question 35

O, unless bonded to F or itself, has an oxidation state of -2.

Answer 35

N/A

Question 36

The sum of all oxidation states in a compound must equal the total charge on the species

Answer 36

N/A

Question 37

1. Separate oxidation and reduction half-reactions.
2. Balance all atoms except for hydrogen and oxygen in each half-reaction.
3. Balance oxygen by adding H2O as needed.
4. To balance hydrogen, add H+ as needed.
5. Balance the charge of each reaction by adding electrons to side with the greater charge.
6. Multiply each half-reaction by the least integer factor that equalizes the number of electrons in each half-reaction. Then, add the half-reactions to obtain the overall balanced reaction in acidic solution.

If your redox reaction is in acidic solution, the above reaction is properly balanced. However, if the reaction you wish to balance is in basic solution, you need to add these three steps:

7. If the redox reaction is one in basic solution, then add OH- to both sides of the equation to “neutralize” each H+.
8. “React” H+ and OH- to form H2O and eliminate water molecules on both sides of the equation.
9. Make sure that all atoms and charges are indeed balanced in your overall balanced equation for the redox reaction in basic solution.

Answer 37

How to solve redox reactions