electron transfer reactions ch 19 oxidation reduction or redox reactions
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Electron Transfer Reactions: CH 19: Oxidation-reduction or redox reactions. Results in generation of an electric current (electricity) or caused by imposing an electric current. In previous chapter energy was in form of heat, here it is electricity.

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electron transfer reactions ch 19 oxidation reduction or redox reactions
Electron Transfer Reactions: CH 19: Oxidation-reduction or redox reactions.
  • Results in generation of an electric current (electricity) or caused by imposing an electric current.
  • In previous chapter energy was in form of heat, here it is electricity.
  • Therefore, this field of chemistry is often called ELECTROCHEMISTRY.
terminology for redox reactions
Terminology for Redox Reactions
  • OXIDATION—loss of electron(s) by a species; increase in oxidation number; increase in oxygen.
  • REDUCTION—gain of electron(s); decrease in oxidation number; decrease in oxygen; increase in hydrogen.
  • OXIDIZING AGENT—electron acceptor; species is reduced.
  • REDUCING AGENT—electron donor; species is oxidized.
you can t have one without the other
You can’t have one… without the other!
  • Reduction (gaining electrons) can’t happen without an oxidation to provide the electrons.
  • You can’t have 2 oxidations or 2 reductions in the same equation. Reduction has to occur at the cost of oxidation

LEO the lion says GER! Another way to remember.



activity series
Activity Series
  • For metals, the higher up the chart the element is, the more likely it is to be oxidized. This is because metals like to lose electrons, and the more active a metallic element is, the more easily it can lose them.
  • For nonmetals, the higher up the chart the element is, the more likely it is to be reduced. This is because nonmetals like to gain electrons, and the more active a nonmetallic element is, the more easily it can gain them.
metal activity
Metal Activity

3 K0 + Fe+3Cl-13


  • Metallic elements start out with a charge of ZERO, so they can only be oxidized to form (+) ions.
  • The higher of two metals MUST undergo oxidation in the reaction, or no reaction will happen.
  • The reaction 3 K + FeCl3 3 KCl + Fe WILL happen, because K is being oxidized, and that is what the activity series says should happen.
  • The reaction Fe + 3 KCl  FeCl3 + 3 K will NOT happen.

Fe0 + 3 K+1Cl-1


electrochemical cells
Electrochemical Cells
  • Allows a redox reaction to occur by transferring e- through an external connector.
  • Product favored reaction --> voltaic or galvanic cell

Converts chemical energy from a spontaneous redoxrxtn into electrical energy.

  • Reactant favored reaction --> electrolytic cell.

Converts chem. energy to electrical energy using a nonspontaneousredoxrxtn & an electrical energy source.

Batteries are voltaic cells

chemical change electric current

With time, Cu plates out onto Zn metal strip, and Zn strip “disappears.”

  • Zn is oxidized and is the reducing agent Zn(s) ---> Zn2+(aq) + 2e-
  • Cu2+ is reduced and is the oxidizing agentCu2+(aq) + 2e- ---> Cu(s)
chemical change electric current1
  • To obtain a useful current, we separate the oxidizing and reducing agents so that electron transfer occurs thru an external wire.

This is accomplished in a GALVANIC or VOLTAIC cell.

A group of such cells is called a battery.


Basic Concepts of Electrochemical Cells

Electrons travel thru external wire.

Salt bridge allows anions and cations to move between electrode compartments, without mixing of metal atoms.

Anode: half cell where oxidation occurs

Cathode: half cell where reduction occurs

Zn --> Zn2+ + 2e-

Cu2+ + 2e- --> Cu







Notation: Anode I Anode product I Salt I Cathode ICathode product

Zn (s) I Zn2+ (aq, 1 M) I NaNO3I Cu2+ (aq, 1 M) I Cu (s)





terms used for voltaic cells
Terms Used for Voltaic Cells

Electrode: metal connected by an external circuit

cell potential e cell emf electromotive force
CELL POTENTIAL,EcellEMF:electromotive force
  • For Zn/Cu cell, potential is +1.10 V at 25 ˚C and when [Zn2+] and [Cu2+] = 1.0 M.
  • —a quantitative measure of the tendency of reactants to proceed to products when all are in their standard states at 25 ˚C.
calculating cell voltage
Calculating Cell Voltage
  • Balanced half-reactions can be added together to get overall, balanced equation.

Zn(s) ---> Zn2+(aq) + 2e-

Cu2+(aq) + 2e- ---> Cu(s)


Cu2+(aq) + Zn(s) ---> Zn2+(aq) + Cu(s)

If we know Eo for each half-reaction, we could get Eo for net reaction.



ability of ion





+ 2e- Cu



2 H

+ 2e- H




+ 2e- Zn


reducing ability

of element

TABLE OF STANDARD REDUCTION POTENTIALSthe standard for measuring is SHE: Standard Hydrogen Electrode


To determine an oxidation from a reduction table, just take the opposite sign of the reduction!

zn cu electrochemical cell
+Zn/Cu Electrochemical Cell

Ox: Zn(s) ---> Zn2+(aq) + 2e- Eo = +0.76 V

Red:Cu2+(aq) + 2e- ---> Cu(s) Eo = +0.34 V


Cu2+(aq) + Zn(s) ---> Zn2+(aq) + Cu(s)

Eo = +1.10 V

Anode, negative, source of electrons

Cathode, positive, sink for electrons

e o for a voltaic cell
Eo for a Voltaic Cell

Cd --> Cd2+ + 2e-


Cd2+ + 2e- --> Cd

Fe --> Fe2+ + 2e-


Fe2+ + 2e- --> Fe

All ingredients are present. Which way does reaction proceed?

e o for a voltaic cell1
Eo for a Voltaic Cell

From the table, you see

• Fe is a better reducing agent than Cd

• Cd2+ is a better oxidizing agent than Fe2+

  • Eo =
  • Eoreduction(reduction process-cathode )– Eoreduction (oxidation process- anode)
  • Positive value indicates spontaneous process. Negative value indicates nonspontaneous.
  • Steps for prediction spontaneity:
  • Write the reaction as oxidation and reduction half reactions.
  • Plug standard reduction potentials in the equation given above.
  • Check for the spontaneity by a positive or negative value.
more about calculating cell voltage
More About Calculating Cell Voltage

2 H2O + 2e- ---> H2 + 2 OH- Cathode

2 I----> I2 + 2e- Anode


2 I- + 2 H2O --> I2 + 2 OH- + H2

Assume I- ion can reduce water.

Assuming reaction occurs as written,

E˚ = E˚cat+ E˚an= (-0.828 V) + (- +0.535 V) = -1.363 V

Minus E˚ means rxn. occurs in opposite direction

(the connection is backwards or you are recharging the battery)

charging a battery
Charging a Battery

When you charge a battery, you are forcing the electrons backwards (from the + to the -). To do this, you will need a higher voltage backwards than forwards. This is why the ammeter in your car often goes slightly higher while your battery is charging, and then returns to normal.

In your car, the battery charger is called an alternator. If you have a dead battery, it could be the battery needs to be replaced OR the alternator is not charging the battery properly.

dry cell battery
Dry Cell Battery

Anode (-)

Zn ---> Zn2+ + 2e-

Cathode (+)

2 NH4+ + 2e- ---> 2 NH3 + H2

alkaline battery
Alkaline Battery

Nearly same reactions as in common dry cell, but under basic conditions.

Anode (-): Zn + 2 OH- ---> ZnO + H2O + 2e-

Cathode (+): 2 MnO2 + H2O + 2e- --->

Mn2O3 + 2 OH-

h 2 as a fuel
H2 as a Fuel

Cars can use electricity generated by H2/O2 fuel cells.

H2 carried in tanks or generated from hydrocarbons

Balance the following in acid solution—

VO2+ + Zn ---> VO2+ + Zn2+

Step 1: Write the half-reactions

Ox Zn ---> Zn2+

Red VO2+ ---> VO2+

Step 2: Balance each half-reaction for mass.

Ox Zn ---> Zn2+


Step 3: Balance half-reactions for charge.

Ox Zn ---> Zn2+ + 2e-

Red e- + 2 H+ + VO2+ ---> VO2+ + H2O

Step 4: Multiply by an appropriate factor.

Ox Zn ---> Zn2+ +2e-

Red 2e-+ 4 H+ + 2 VO2+ ---> 2 VO2+ + 2 H2O

Step 5: Add balanced half-reactions

Zn + 4 H+ + 2 VO2+ ---> Zn2+ + 2 VO2+ + 2 H2O

Add H2O on O-deficient side and add H+ on other side for H-balance.

2 H++

VO2+ ---> VO2+ + H2O