Download Electrochemistry power point

Electrochemistry
Oxidation – Reduction (Redox)
Reactions


Involves the transfer of electrons from one
species to another
Both oxidation and reduction must occur
Burning Magnesium Demo












The balance of the overall reaction seen above can be written as follows.
2 Mg (s) + O2 (g) -------> 2 MgO
1) What is happening to magnesium in this reaction?
Answer: Magnesium is losing electrons to become a positive ion . This is the process of
oxidation.
The reaction we can write for this process is called the oxidation half - reaction.
Mg0 (s) --------> Mg 2+ + 2 e - ( electrons )
Since two magnesium atoms are reacting a total of 4 electrons are lost by the
magnesium
2) What is happening to the Oxygen in this reaction?
Answer: Oxygen is gaining electrons to become a negative ion. This is the process of
reduction.
The reaction we write for this process is called the reduction half- reaction.
O0 (g) ----------> O 2- + 2e Since oxygen is a diatomic molecule two oxygen atoms are undergoing reduction
and a total of 4 electrons are gained
Note: The total number of electrons lost in oxidation = the total number gained in reduction.
This is true of all redox reactions
Oxidation – Reduction (Redox)
Reactions





Oxidation number: Charge on a single atom or ion
or the apparent charge on atoms in multiatom
combinations
Oxidation: the lose of electrons, oxidation number
increases
Reduction: the gain of electrons, oxidation number
decreases
Oxidizing Agent (oxidant): substance oxidizing
another (being reduced)
Reducing Agent (reductant): substance reducing
another (being oxidized)
Oxidation – Reduction (Redox)
Reactions








Oxidation state of a free element is zero
Oxidation state of group I, II, III metals: +1, +2, +3
Oxidation state of F is always –1
Oxidation state of H is +1 except in metallic hydrides
where it is –1
Oxidation state of O is –2 except
with F where it is +2,
in peroxides where it is –1 or
in superoxides where it is –1/2
Oxidation state of monatomic ion = ion’s charge
Sum of oxidation states of all atoms in a polyatomic
ion = ion’s charge
Sum of oxidation states of all atoms or ions in the
formula of a compound = zero
Example
In each of the following equations, is hydrogen
peroxide acting as a reducing or oxidizing
agent?
 2MnO4-(aq) + 5H2O2(aq) + 6H+(aq) 
2Mn2+(aq) + 5O2(g) + 8H2O(l)
 PbS(s) + 4H2O2(aq)  PbSO4(aq) + 4 H2O(l)
Balancing Redox Equations
Half Reaction Method (Acidic Solution)








Write the oxidation half reaction ( electrons are
products)
Write the reduction half reaction ( electrons are
reactants)
Balance elements other than H and O in each
reaction
Balance O by adding H2O
Balance H by adding H+
Balance charges by adding electrons
Add the 2 half reactions and simplify where possible
Check
Balancing Redox Equations Half
Reaction Method (Basis Solution)


Follow steps 1 thru 5 as if it were in acidic
solution
Balance H+ by adding OH-
Example:



Complete and balance the following
equations by the method if half reactions.
Cr2O72-(aq) + Cl-(aq)  Cr3+ (aq) + Cl2(g)
(acidic)
CN-(aq) + MnO4-(aq)  CNO-(aq) + MnO2(s)
(basic)
Voltaic (galvanic) Cells








Energy released in a chemical reaction can be used to do
electrical work
Device used to perform electrical work by the transfer of
electrons through an external pathway
Electrodes: 2 solid metals connected by the external circuit
Cathode (+): electrode where reduction occurs
Anode (-): electrode where oxidation occurs
Salt Bridge: Maintains neutrality in a galvanic cell
Contains an electrolytic solution whose ions will not react with
other ions in the cell
Electron Flow: From anode to cathode
Voltaic (galvanic) Cells
Standard cell notation:
anode/anode solution // cathode solution/
cathode
Zn/Zn2+ (1.0 M) //Cu2+ (1.0 M)/Cu
 Voltmeter: Measures the cell potential
usually in volts

Voltaic Cell Animation


Voltaic cell animation
http://www.chem.iastate.edu/group/Greenbo
we/sections/projectfolder/flashfiles/electroCh
em/voltaicCellEMF.html
Example:
The following redox reaction is spontaneous
Cr2O72-(aq) + 14H+(aq) + 6I-(aq)  2Cr3+(aq) + 3I2(s) +
7H2O(l)
 A solution containing K2Cr2O7 and H2SO4 is poured into
one beaker and a solution of KI is poured into another
beaker. A salt bridge is used to join the beakers. A
metallic conductor that will not react with either solution
is suspended in each solution and the two conductors
are connected with wires through a voltmeter. Indicate
the reaction occurring at the anode, the reaction at the
cathode, the direction of electron and the ion
migrations, and the signs of the electrodes

Cell EMF (Electromotive Force)




The potential difference between the two
electrodes of a voltaic cell which provides the
driving force that pushes the electrons
through the external circuit
EMF of a cell (Ecell) is also called cell
potential
Measured in volts, so called cell voltage
Must be positive value if spontaneous
reaction
Standard Reduction Potential (Eored)


The tendency of a given half cell to occur as
a reduction
The most positive value on the reduction
potential table will be the reduction half
reaction
Standard Cell Potential (Eocell)



The standard cell potential is the measured
cell potential when the ion concentrations in
the half cells are 1 M, any gases are at a
pressure of 1 atm and the temperature is
25oC.
Ecell = Ered + Eox
Positive values of E mean spontaneous
Example:


A voltaic cell is based on the following two
standard half reactions:
Cd2+(aq) + 2 e-  Cd(s)
Sn2+(aq) + 2 e-  Sn(s)
By using the data in appendix E, determine
the half reactions that occur at the cathode
and the anode and the standard cell potential
Practice:


A voltaic cell is to be designed at standard state
conditions using the following two half reactions.
Mn2+(aq) + 2e-  Mn(s)
Eored = -1.18 V
Cr3+(aq) + 3e-  Cr(s)
Eored = -0.74 V
Write the chemical equation describing the
spontaneous cell reaction and calculate Eocell. Draw
a figure representing the voltaic cell. Label all
components.
Practice

Determine which one of the three metals Zn,
Fe, or Na is the most active metal given the
following data:
Na+(aq) + e- Na(s)
Eored = -2.71 V
Zn2+(aq) + 2e-  Zn(s)
Eored = -0.76 V
Fe2+(aq) + 2e-  Fe(s)
Eored = - 0.44 V
Oxidizing Agents:
Reduced in a redox reaction
Large positive E values on reduction potential table

MnO4-(aq)  Mn2+(aq)
(in acidic)

MnO4-(aq)  MnO2 (s)
(in basic)

Cr2O72-(aq)  Cr3+(aq)
(in acidic)

CrO42-(aq)  Cr(OH)3(s) (in basic)
All strong nonmetals are excellent oxidizing agents

O2(g)  O2- (combustion if fast, corrosion if
slow)

Cl2(g)  Cl- (free halogens reduce to halide
anions)
Reducing Agents:
Oxidized in a redox reaction
More negative E values on reduction potential table

HSO3-(aq)  SO42(in acidic)

SO32-(aq)  SO42(in basic)

S2O82-(aq)  SO42(in acidic with Cl2)

S2O82-(aq)  S4O62(in acidic with I2)

Fe2+(aq) or Sn2+(aq)  Fe3+(aq) or Sn4+(aq)
Also, all strong metals are excellent reducing agents
 Na(s)  Na+
 Ba(s)  Ba2+
(free metals oxidize to metal cations)
Spontaneity of Redox Reactions




Ecell = Ered + Eox
Positive value of Ecell mean spontaneous
∆G = - nFE
n = # electrons transferred, E
is emf (voltage)
F (Faraday’s constant) = 96,500 C/mol =
96,500 J/V-mol
Nernst Equation:




Relates cell emf to concentration
E = Eo – (RT/nF)lnQ
E =Eo – (2.303 RT/ nF)log Q
E =Eo – (0.0592 / n)log Q
Q is reaction quotient
F is Faraday’s constant
N is # electrons transferred
Cell EMF and Chemical Equilibrium




EMF of a voltaic cell drops as it discharges
As reactants convert to products the value of
Q increases so the value of E decreases until
E =0
If E = 0 then ∆G = 0 and the system is at
equilibrium and Q = Keq
Log Keq= nEo / 0.0592
Batteries
A battery is a portable, self-contained electrochemical power source
of one or more voltaic cells.
 Lead – Acid Battery
Automotive battery uses 6 voltaic cells in series
Cathode is PbO2 and anode is Pb
Both electrodes are immersed in sulfuric acid
 Alkaline Battery
Anode is powdered Zn in contact with KOH
Cathode is graphite and MnO2
 Nickel-cadmium, Nickel-metal halide and lithium-ion batteries
Light weight rechargeable
Cd is anode and NiO(OH)(s) is cathode
 Corrosion Reactions: Spontaneous redox reactions that lead to
the corrosion of metals
Lead Battery Recycling
Lead-Acid Battery
Lead-Acid Battery
Alkaline Battery
Nickel-Cadmium Battery
Cathodic Protection




Prevents corrosion of iron
Make the metal the cathode in an
electrochemical cell
Sacrificial anode is the metal that will oxidize
Choose an anode metal with a reduction
potential that is less than that of the metal to
be protected. These metals are more easily
oxidized
Electrolysis




It is possible to use electricity to cause nonspontaneous redox
reactions to occur
These are called electrolysis reactions and they take place in
electrolytic cells
Electrolysis of molten salts is an important process for the
production of active metals such as sodium.
w = -nFE
1 F = 96500 C/mol eCouloumbs = amperes X seconds
watt (W) is a unit of electrical power
1 W = 1 J/s
If given current and time  find quantity of charge (coulombs) 
find moles of electrons (faradays)  find moles of substance
oxidized or reduced  find grams of substance
Electrolytic cell with salt bridge
Electrolysis



http://www.cressex.bucks.sch.uk//department
s/Sc/revision%20movies/electrolysis.htm
electrolysis of copper (II) chloride
Examples:


What mass of zinc metal is produced at the
cathode in an electrolysis cell when a
constant current of 10.00 amp is passed for 1
hr?
Calculate the mass of Mg produced in an
electrolytic cell if 2.67 x 102 kwh of electricity
is passed through a solution of MgCl2 at 4.20
volts.