Chemical Reactions 2: Equilibrium & Oxidation-Reduction

Redox Reactions

Neutral atoms do not have charge since number of electrons equals number of protons (charge equals zero).

Charge is acquired when an atom gains (- charge) or loses (+ charge) electrons (last shell)

Redox Reactions

Oxidation

Process of losing electrons (usually in last shell) Most likely to occur to metals Element “gains” charge (e.g. O2- oxidizes to O, so charge changes from -2 to 0)(e.g. Zn oxidizes to Zn2+, so charge changes from 0 to +2)

Sodium lost one electron. It oxidized, so

from Na to Na+

Redox Reactions

Reduction

Process of gaining electrons (usually in last shell) Most likely to occur to non-metals Element “lose” charge (e.g. O oxidizes to O2-, so charge changes from 0 to -2)(e.g. Cu2+ oxidizes to Cu+, so charge changes from +2 to +1)

Chlorine gained one electron. It

reduced, so from Cl to Cl-

Redox Reactions

Oxidation half reaction produces electrons (M→M+ + e-)

Reduction half reaction consumes electrons (N + e- →N-)

Redox Reactions

Identify which reaction involves a reduction, and which an oxidation:

_Zn → Zn2+ + 2e-

_S + 2e- → S2-

_Fe2+ → Fe3+ + e-

_Al + 3e- → Al3-

Oxidation

Reduction

Oxidation

Reduction

Redox Reactions

Oxidizing agent: The one reactant that reduces in a redox reaction

(N + e- →N-)

N reduces, so it is the oxidizing agent (makes M undergo oxidation)

Reducing agent: The one reactant that oxidizes in a redox reaction

(M→M+ + e-)

M oxidizes, so it is the reducing agent (makes N undergo reduction)

Redox Reactions

Copper. Cu2+(aq) + 2e- → Cu(s)

Zinc. Zn(s) → Zn2+(aq) + 2e-

Copper reduces. Zinc oxidizes

Copper, oxidizing agent. Zinc, reducing agent

Redox Reactions

Oxidation and Reduction occur simultaneously There cannot be one without the other Both can be described by half-reactions Total redox reactions needs to have same amount of

electrons in both half reactions

Redox Reactions

Redox Reactions

Redox Reactions

Redox Reactions

Spontaneous Redox Reactions (Exothermic reactions)

_Half-redox reactions are ranked according to their standard reduction potential, which is a measure of the tendency of an element to gain electrons

_For a redox reaction to be spontaneous, the species acting as oxidizing agent (the one who reduces) must have a higher standard reduction potential than the species acting as reducing agent (the one who oxidizes)

Redox Reactions

E° = -1.18VE° = -2.37V

E° = 1.99V

E° = -0.13V

E° = -0.23V

E° = -1.66V

Redox Reactions

E° = -0.14V

E° = -2.37V

E° = 0.00V

E° = -0.73V

E° = 1.50V

E° = 0.34V

Redox ReactionsVolta’s cell was the first attempt to produce electricity.

***Even though Volta had little understanding of the way its cell worked, his discovery contributed to:

_Development of electrochemistry

_Discovery of new chemical elements

Redox ReactionsDaniel’s cell

_First truly usable cell

_Very heavy and big equipment needed

_Composed of:

Anode (-) (electrode where oxidation takes place)

Cathode (+) (electrode where reduction takes place)

*Electrons flow from anode to cathode

Redox Reactions

Redox Reactions

Cell Potential Difference

ΔE° = E°cathode - E°

anode (Reduction) (Oxidation)

*each E° is measured against the reduction potential of hydrogen electrode (zero)

Redox Reactions

ΔE° = E°cathode - E°

anode

ΔE° = (0.34V) – (-0.76V)ΔE° = 1.10 V

Calculate ΔE° if you replace Zn by Mg:

ΔE° = E°cathode - E°

anode

ΔE° = (0.34V) –(-2.37V)ΔE° = 2.71 V

Redox ReactionsCell notation

Ag(s)/Ag+(aq)||H+

(aq)/H2(g)

Anode (Oxidation):Ag(s)→Ag+

(aq) + 1e-

Cathode (Reduction):2H+

(aq) + 2e- →H2(g)

Cell reaction:Ag(s) + 2H+

(aq) →Ag+(aq) + H2(g)

|| (Salt bridge): maintains electrical neutrality of solutions in half cells

Anode Cathode

Electrons move from anode to cathode

Redox Reactions

Redox Reactions

Redox Reactions