Chemistry Ch. 7

Front 1

Oxidized

Back 1

-An atom that looses electrons

Front 2

Reduced

Back 2

-An atom that gains electrons

Front 3

Reducing agent (reductant)

Back 3

-A species that gives its own electrons to another atom (an atom in the species is oxidized in the reaction)

Front 4

Oxidizing agent (oxidant)

Back 4

-A species that takes electrons from another atom (an atom in the species is reduced in the reaction)

Front 5

Electric potential (E) in redox reactions

Back 5

-Since in a redox reaction electrons are transferred, and since electrons have charge, there is an electric potential associated with any redox reaction.

Front 6

Reduction potentials

Back 6

-The electric potential created when a species is reduced. Shown as a half reaction.
-Compared to an SHE (standard hydrogen electrode)
-A positive reduction potential indicates that the atom is easily reduced
-A negative reduction potential indicates that the atom is more easily oxidized.

Front 7

The standard hydrogen electrode half reaction

Back 7

2H+ + 2e- --> H2 E= 0.00V

-the standard to which all reduction potentials are compared.

Front 8

Finding the potential of an ionic reaction

Back 8

-Find the reduction potential of each half reaction
-If the direction of the half reaction is reversed, reverse the sign
-Add the potentials together (do NOT multiply by the number of times the reaction occurs, reduction potential is an intensive property)

Front 9

7 steps for balancing a reaction in acidic solution

Back 9

1. Divide the reaction into its corresponding half reactions.
2. Balance the elements other than H and O
3. Add H2) to one side until the O atoms are balanced
4. Add H+ to one side until the H atoms are balanced
5. Add e- to one side until the charge is balanced
6. multiply each half reaction by an integer so that an equal number of electrons are transferred to each reaction.
7 Add the two half reactions and simplify.

Front 10

Galvanic (voltaic) cell

Back 10

-uses the electric potential between multiple phases to generate a current of electrons from one phase to another in a conversion of chemical energy to electrical energy.
-all phases must conduct electricity
-since the electrons in the anode have a higher potential energy than those in the cathode, electrons flow through the load from the anode to the cathode.
-the cell potential of a galvanic cell is always positive; always has chemical energy that can be converted to work.

Front 11

Salt bridge

Back 11

-a type of liquid junction that minimizes the potential difference. (when there are two separate solutions at the anode and the cathode).
-cell will short circuit without this.
-Typically made form an aqueous solution of of KCl.
-Allows ionic conduction between solutions without creating a a strong extra potential within the galvanic cell.
-in a galvanic (voltaic) cell, a phase that is impermeable to electrons which is an ionic conductor carrying the current in the form of ions.
-often an electrolyte solution

Front 12

Terminals of a voltaic (galvanic) cell

Back 12

electron conductors such as metal whires

Front 13

Electrodes of a voltaic (galvanic) cell

Back 13

-the anode and cathode

Front 14

The components of a simple galvanic cell

Back 14

T-E-I-E'-T'

T-terminal
E-electrode
I-Ion conductor (salt bridge)
E'-second electrode
T'-second terminal

Front 15

Anode of galvanic cell

Back 15

-marked with a negative sign
-sight of the oxidation half reaction
-looses electrons, which then flow towards the

(AN OX)

Front 16

Cathode of a galvanic cell

Back 16

-the electrode in a galvanic cell at which the reduction half reaction occurs
-because it is positive, attracts electrons from the anode.

(RED CAT)

Front 17

Cell potential (E)

Back 17

-also called the electron motive force (EMF)
-the potential difference between the terminals when the are not connected.

Front 18

Electron motive force

Back 18

-the potential difference between the terminals of a battery when they are not connected
-also known as cell potential (E)

Front 19

Equation relating Gibbs free energy and cell potential

Back 19

DeltaG = -nFE(max)

n = number of moles of electrons that are transferred in the balanced redox reaction
F = Faraday's constant
E = cell potential

-a positive cell potential always indicates a spontaneous reaction (deltaG is negative)

Front 20

Equation relating reaction under nonstandard conditions and Gibbs free energy

Back 20

deltaG = G(standard) + RTln(Q)

R = Gas constant
T = temperature (Kelvin)
Q = reaction quotient

If use only 1 molar concentrations, Q = 1, RTln(Q) = 0, leaving deltaG = G. (standard conditions don't indicate temperature)

Front 21

Equation relating free energy of an equation at equilibrium

Back 21

deltaG(standard)= -RTln(K)

-At equilibrium, there is no available free energy to do work, deltaG = 0 by definition.
-In this equation, both K and deltaGstandard vary with temperature.

Front 22

Concentration cell

Back 22

-a limited form of a galvanic cell with a reduction half reaction taking place in one half of the cell and the exact reverse of that half reaction taking place in the other half cell.
-due to entropy, the more concentrated side will become less concentrated and electrons will flow accordingly.

Front 23

Electrolytic cell

Back 23

-If a power source is hooked up across the resistance of a galvanic cell, an electrolytic cell is created.
-Any electrolytic cell will have a negative EMF.
-the cathode is marked as negative, the anode is marked as positive. Reduction still takes place at the cathode and oxidation takes place at the anode.