Mcat Chemistry Lecture7

Front 1

Redox reaction

Back 1

oxidation-reduction reaction

electrons are transferred form one atom to another

atom that loses electrons is oxidized

atom that gains electrons is reduced

Front 2

Oxidized

Back 2

atom loses electrons in redox reaction

Front 3

reduced

Back 3

atom gains electrons in redox reaction

Front 4

oxidation states

Back 4

possible charge values that an atom may hold within a molecule

necessary for redox reactions

must add up to charge on molecule or ion

Front 5

oxidation state = 0

Back 5

atoms in their elemental form

Front 6

oxidation state of fluorine (F)

Back 6

equals -1

Front 7

oxidation state of hydrogen (H)

Back 7

equals +1

except when bonded to a metal; then -1

Front 8

oxidation state of oxygen (O)

Back 8

equals -2

except when it is in a peroxide like H2O2

Front 9

Oxidation state = +1

Back 9

group 1 elements

alkali metals

Front 10

oxidation state = +2

Back 10

group 2 elements

alkaline earth metals

Front 11

oxidation state = -3

Back 11

group 15 elements

nitrogen family

Front 12

oxidation state = -2

Back 12

group 16 elements

oxygen family

Front 13

oxidation state = -1

Back 13

group 17 elements

halogens

Front 14

LEO the lion says GER

Back 14

LEO: Lose Electrons Oxidation

GER: Gain Electrons Reduction

Front 15

Reducing agent

Back 15

reductant

giving electrons to an atom

losing electrons, is oxidized

Front 16

Oxidizing agent

Back 16

compound containing atom being reduced

gains electrons, is reduced

oxidizes other atom

Front 17

Electric potential (E)

Back 17

associated with any redox reaction

Front 18

Half-reaction

Back 18

Each component of a redox reaction


oxidation half-reaction potential is opposite reduction half-reaction potential

usually listed as reduction potentials (sign is reversed for oxidation potential)

Front 19

Half-reaction of standard hydrogen electrode

Back 19

2H+ + 2e- --> H2
Half-reaction potential = 0.00V

Front 20

Balancing redox reactions in acidic solutions

Back 20

1. divide reaction into its corresponding half reactions
2. balance the elements other than H and O
3. Add H20 to one side until O atoms are balanced
4. Add H+ to one side until H atoms are balanced
5. Add e- to one side until charge is balanced
6. multiply each half reaction by an integer so that an equal number of electrons are transferred in each reaction
7. add the 2 half reactions and simplify

Front 21

Balancing redox reactions in basic solutions

Back 21

same steps as acidic solutions

neutralize H+ ions by adding same number of OH- ions to both sides of reaction

Front 22

Galvanic cell (voltaic cell)

Back 22

uses electric potential between phases to generate a current of electrons from one phase to another in a conversion of chemical energy to electrical energy

turns chemical energy into electrical energy

Front 23

Salt bridge

Back 23

ionic conducting phase

electrolyte solution

phase impermeable to electrons

type of liquid junction that minimizes potential difference between different solutions

carries current in form of ions

Front 24

Terminals

Back 24

electronic conductors such as metal wires (T)

Front 25

Electrodes

Back 25

electronic conductors (E)

Front 26

Ionic conductor

Back 26

salt bridge (I)

Front 27

Simple galvanic cell

Back 27

T-E-I-E'-T'

has 2 electrodes: anode (-) and cathode (+)

oxidation half reaction takes place at anode

reduction half reaction takes place at cathode

2 terminals of cell is made from same material

Front 28

cell potential (E)

Back 28

electromotive force (emf)

potential difference between the terminals when they are not connected

connecting the terminals reduces the potential difference due to internal resistance within the cell

drop in emf increases, as current increases

current flows in direction opposite electron flow

electrons flow from anode to cathode

always positive, always has chemical energy than can be converted to work

Front 29

RED CAT

Back 29

REDuction CAThode

Front 30

AN OX

Back 30

ANode OXidation

Front 31

Positive cell potential

Back 31

equals a negative Gibbs free energy (deltaG), which equals a spontaneous reaction (work is being done by system and not on system)

deltaG = -nFEmax
deltaG: Gibbs free energy
n: number of moles of electrons that are transferred in balanced redox reaction
F: Faraday's constant
E: voltage

Free energy (deltaG) represents the product of total charge (nF) times voltage (E)

Front 32

Reactions that do not occur at standard state

Back 32

deltaG = deltaG(not) + RT[ln(Q)]
deltaG: Gibbs free energy
deltaG(not): Gibbs free energy (standard conditions)
T: temperature
Q: reaction quotient

Front 33

reactions that are at equilibrium conditions

Back 33

at equilibrium, there is no available free energy with which to do work; deltaG = 0

deltaG(not) = -RT[ln(K)]

Front 34

relationship between K and deltaG(not)

Back 34

if K = 1, then deltaG(not) = 0

if K > 1, then deltaG(not) < 0

if K < 1, then deltaG(not) > 0

Front 35

Concentration cell

Back 35

limited form of a galvanic cell with a reduction half reaction taking place in 1 half cell and the exact reverse of that half reaction taking place in the other half cell

type of galvanic cell

it is never at standard conditions, so Nerst equation is required to solve for cell potential

if concentrations were equal on both sides, the concentration cell potential would be zero

Front 36

galvanic cell

Back 36

positive cell potential

spontaneous

Front 37

electrolytic cell

Back 37

negative cell potential

forced by outside power source to run backwards

cathode is negative
anode is positive

RED CAT & AN OX still the same