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

What happens to oxidized atom?

2H2 + O2 -> 2H2O

Back 2

atom loses electrons in redox reaction

Hydrogen is oxidized from 0 to +1

Front 3

What happens to reduced atom?

2H2 + O2 -> 2H2O

Back 3

atom gains electrons in redox reaction

Oxygen is reduced from 0 to -2

Front 4

oxidation states

Back 4

possible charge values that an atom may hold within a molecule

the oxidation state must add up to charge on molecule or ion; so the oxidation states of the atoms in a neutral molecule must add to 0!

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 like NaH; then -1

Front 8

oxidation state of oxygen (O)?

Back 8

equals -2

except when it is in a peroxide like H2O2; then -1

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

CH4 + 2O2 -> CO2 + 2H2O

Back 15

reductant. gives electrons to an atom. loses electrons, is oxidized

CH4 is reducing agent, Carbon goes from -4 to +4, is oxidized

Front 16

Oxidizing agent

CH4 + 2O2 -> CO2 + 2H2O

Back 16

compound containing atom being reduced. gains electrons, is reduced. oxidizes other atom

O2 is oxidizing agent, Oxygen goes from 0 to -2, is reduced.

Front 17

Electric potential (E)

Back 17

associated with any redox reaction because electrons are transferred

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

Galvanic cell (voltaic cell)

Back 20

uses electric potential between 2 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, it is a battery

Front 21

Salt bridge

Back 21

a type of liquid junction that minimizes potential difference between different solutions

the ionic conducting phase. The salt bridge allows ionic conduction between solutions without creating a strong extra potential within the galvanic cell

electrolyte solution

phase impermeable to electrons

carries current in form of ions

Front 22

Terminals

Back 22

electronic conductors such as metal wires (T)

Front 23

Electrodes

Back 23

electronic conductors (E)

Front 24

Ionic conductor

Back 24

salt bridge (I)

Front 25

Simple galvanic cell

Back 25

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 26

cell potential (E) aka electromotive force (emf)

Back 26

the 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 through the load (resistance) from anode to cathode

Cell potential for a galvanic cell is always positive, always has chemical energy that can be converted to work

Front 27

RED CAT

Back 27

REDuction CAThode

Front 28

AN OX

Back 28

ANode OXidation

Front 29

Positive cell potential

Back 29

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

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

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

Front 30

Reactions that do not occur at standard state

Back 30

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

Front 31

reactions that are at equilibrium conditions

Back 31

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

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

Front 32

relationship between K and ∆G(not)

Back 32

if K = 1, then ∆G(not) = 0

if K > 1, then ∆G(not) < 0

if K < 1, then ∆G(not) > 0

Front 33

Concentration cell

Back 33

a 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 34

galvanic cell

Back 34

positive cell potential

spontaneous

Front 35

electrolytic cell

Back 35

negative cell potential

forced by outside power source to run backwards

cathode is negative, anode is positive

RED CAT & AN OX still the same