Chapter 19:  Oxidation-Reduction Reactions

Oxidation and Reduction

I.  Types of reactions

A.     no apparent change in the electronic structure, such as double displacement and most acid-base

neutralization reaction

B.     changes in the electronic structure:  electrons are transferred from one atom to another or there is a rearrangement of the sharing; 

1.      the oxidation number of an atom - by itself, as part of a molecule or part of a polyatomic ion –

appears to change

2.      oxidation number:  the apparent charge on an atom based on it’s relative electronegativity

compared with other atoms in the molecule/ion

3.  these reactions are called oxidation-reduction reactions (or redox for short) 

a.      the atom that loses, or apparently loses, the electron(s) is said to be oxidized; since electrons

are negatively charged, the atom losing the electron(s) will become more positively charged;  the oxidation number increases

b.      the atom that gains, or apparently gains, the electron(s) is said to be reduced; since electrons

are negatively charged, the atom gaining the electron(s) will become more negatively charged;  the oxidation number is decreases (REDUCED)

c.      there is no reduction if there is no oxidation,  electrons that are lost must go someplace and

those that are gained must come from someplace

d.      a handy way to remember which is which is:  LEO says GER  (Lose Electrons Oxidized, Gain Electrons Reduced)        or   OIL RIG  (Oxidized Is Loss, Reduction Is Gain)

II.  You must be able to determine oxidation numbers of atoms in ions or compounds or of an element to

find out if an atom has been oxidized or reduced.  You must MEMORIZE the following rules:

      RULE 1:  The oxidation number of any free (uncombined) element is 0.  Cu is 0, H2 is 0, Na is 0, etc.

      RULE 2:  The oxidation number of a monatomic ion is equal to the charge on the ion.  Cu+2 is +2,

Ag+1 is +1, S–2       is –2, N–3 is –3, etc.

RULE 3:  In compounds, the element of Group IA and Group IIA and aluminum have positive

oxidation numbers numerically equal to their group number in the periodic table.  K is +1, 

Ca is +2, and Al is +3.

      RULE 4:  The oxidation number of each hydrogen atom in most common compounds is +1.  The

exception is a hydride, where it is –1.  LiH, NaH, CaH2, BaH2,  etc. the H is –1, these are the reactive metals on the left side of the periodic table.

      RULE 5:  The oxidation number of each oxygen atom in most common compounds is –2.  Two

exceptions are:  1)a peroxide where it is –1.  H2O2, Na2O2, etc.  the O is –1, these are the reactive metals on the left side of the periodic table, and 2)when combined with fluorine it is +2.

      RULE 6:  In a binary compound, the more electronegative element is assigned it’s oxidation number as

an ion.  Group VIIA would usually be assigned a –1.

      RULE 7:  The oxidation number of fluorine is –1.

      RULE 8:  The sum of the oxidation numbers of all the atoms in a particle must equal the apparent

charge of that particle.  H2O has a total charge of 0,  NH4+1 has a total charge of +1,  SO4–2 has a total charge of –2,  NO3–1 has a total charge of –1, etc.

 

Oxidizing and Reducing Agents

I.        An oxidizing agent is a compound, ion or element that causes oxidation; it contains the atom that is

reduced (gained the electron(s))

II.     A reducing agent is a compound, ion or element that causes reduction; it contains the atom that is oxidized (lost the electron(s))

III.   The relative strength of oxidizing and reducing agents is found in Table 19-3 on page 603 in the

textbook.

 

Balancing oxidation-reduction reactions

Redox reactions balance the type of elements, the numbers of those elements, and the charge.

      A.  Assign oxidation numbers to every atom.

      B.  Decide what is oxidized and what is reduced.

C.     Redox reactions are made of two distinct processes (oxidation and reduction).  Therefore these

reactions can be written as two half reactions:  the oxidation half reaction and the reduction half reaction.

            1.  Write the two half reactions with the number of electrons lost or gained per atom. 

            2.  Electrons lost are written on the right side of the oxidation half reaction as a product. 

            3.  Electrons gained are written on the left side of the reduction half reaction as a reactant.

            4.  Balance the number of atoms of the element oxidized or reduced on the reactant and product

side. 

            5.  Balance the number of electrons gained or lost with the number of atoms.

D.     Balance the number of electrons lost with the number of electrons gained by using the lowest

common multiple and multiplying the entire half reaction by the correct number.

      E.  Add the two half reactions.

F.      Check the TOTAL charge on the reactant side and the product side.  If they are not equal, do one of the following:

            1.  in an acidic solution:  add H+1 to balance the charge.

            2.  in a basic solution:  add OH-1 to balance the charge.

G.     Check the number of hydrogen atoms on the reactant side and the product side.  If they are not

equal, add the correct number of water molecules to balance - remembering that every water molecule has two hydrogens.

H.     Check the number of oxygen atoms on the reactant side and the product side.  They should now be

equal.  If not -       start at the beginning and recheck each step.

 

Electrochemistry

A.  Electrochemistry - the study of the interchange or interaction of chemical and electrical energy

1.   A redox reaction (chemical change or reaction) causes an electric current:  voltaic cell

2.   electrical current causes a redox reaction (chemical change or reaction):  electrolytic cell

  1. Conductors -  metals - electrons free to move in conduction band;  ionizable compounds in

solutions - electrolytes, dissociate to form ions in solution

C.  Work -  is done when electricity flows through a wire

D.  Galvanometer -  instrument used to measure electric current using two dissimilar metals

  1. Salt bridge -  a U-shaped tube used to connect two separate solutions without mixing them, it

contains an ionic compound (an electrolyte) in solution  (a porous disk maybe used in place of a salt bridge, it's function is the same)

F.  Electric current - the flow of electrons through a conductor

G.  Volts - measure of electric potential difference

H.  Voltage - potential difference

  1. Potential difference - measure of the relative tendency of two substances to gain electrons; therefore

it refers to reduction reactions

J.  Amperes - the rate of flow of electrons, current

K.  Insulators vs. conductors - insulators have large forbidden zones and it is difficult for electrons to

gain enough energy to jump this zone and enter the conduction band; but if the  P.E. is high enough, the insulator can break down and conduct:  conductors have no forbidden zone so is it very easy for electrons to move into the conduction band

  1. Redox in a solution:  the free movement of ions between electrodes in an aqueous solution, electrons are transferred directly – spontaneous
  2. Electrolyte - a substance that dissociates in an aqueous solution forming ions; the solutions

conduct electricity

N.  Construction of an electrolytic cell

1.      Redox in a solution:  because electrons are transferred directly when particles collide, no useful work can be done;  therefore the two half reactions must be separated from each other

            a.  each separated half reaction is called a half cell

            b.  the two half cells make up a cell

            c.  the electrodes are connected by an external wire 

d.      electrons flow from the reducing agent (containing the atom oxidized) to the oxidizing agent

(containing the atom reduced) through an external wire - work can be done 

e.      in this process there is a build up a negative charge in one half cell and a build up of positive

charge in the other half cell causing a polarization of the cell;  a salt bridge or porous disk is needed to allow positive and negative ions to migrate and balance the charge in the half cells

            f.  electron transfer occurs at the interface between the electrode and the solution 

      2.  Cathode – electrode where reduction occurs

      3.  Anode – electrode where oxidation occurs

4.      Inert electrode - place of oxidation or reduction but the electrode itself is not oxidized or

reduced, examples are carbon rods and platinum used as electrodes

5.      if two reactions can occur at an electrode, the reaction with the lower potential difference will

predominate

      6.  Cations - positively charged ions;  attracted to and move toward the cathode

      7.  Anions - negatively  charged ions;  attracted to and move toward the anode

  1. What happens at the anode?  anions (or an atom) have higher potential energy than the anode so

electrons move from the anions (or atom) to the anode:  anions (or atoms) are oxidized at the anode by losing electrons;  this is the oxidation half-      reaction of a redox reaction

  1. What happens at the cathode?  cations (or an atom) have lower potential energy than the cathode so

electrons move from the cathode to the cations (or atom):  cations (or atoms) are reduced at the cathode by gaining electrons;  this is the reduction half-reaction of a redox reaction

  1. A handy way to remember electrochemical cells:  EMFAT Cat – Electrons Move From Anode To

Cathode;  and Red Cat - Reduction at the Cathode

  1. Electrolytic cell:  electrolysis - electric current used to produce a chemical change, requires outside energy - not spontaneous

S.  Electronic conduction  or metallic conduction:  movement of electrons through a metal

  1. Voltaic cell or galvanic cell:  a battery - a chemical change (redox reaction) is used to produce or generate an electric current; oxidizing agent and reducing agent are separate and electrons travel from anode to cathode through a wire or an external circuit

U.  Representation of cell reactions -  standard line notation

      1.  with both electrodes being conducting metals

            X/X+2//Y-1/Y represents

a.      X being oxidized to X+2 ( X Υ X+2 + 2e- ) at the anode and

b.      Y-1 being reduced to Y ( Y-1 + 1e- Υ Y ) at the cathode

c.      the overall reaction is X  + 2Y-1 Υ 2Y + X+2

d.      the // represents the salt bridge

      2.  with the anode a conducting metal and the cathode an inert electrode

            X/X+2//(Pt) Z-3/Z;  platinum (Pt) is the inert electrode

V.  Standard reduction potentials

1.      Table 19-4 on page 615 in the textbook

2.      all half-reactions listed are REDUCTION half-reactions and are referenced to the hydrogen half-reaction which is zero

W.  Cell Potential

      1.  add the potential difference of the two half-reactions

      2.  the oxidation half-reaction must be switched and the sign listed in the chart must be reversed

      3.  a reaction will be spontaneous when the cell potential is positive

            4.  if the cell potential is negative you must add enough energy to overcome the negative potential

for the reaction to take place