Download Oxidation/Reduction

Oxidation/Reduction
- originally, oxidation described rxns in which oxygen was
added to a reactant
- reduction first meant the removal of oxygen from a
compound
- today’s definition:
- oxidation is the process by which a substance loses
one or more electrons
- reduction is the process by which a substance gains
one or more electrons
- mnemonic devices:
- OIL RIG: Oxidation Is Loss of electrons; Reduction Is
Gain of electrons
- LEO the lion goes GER: Loss of Electrons is
Oxidation and Gain of Electrons is Reduction
- Oxidation and Reduction ALWAYS occur together
- the two rxns together are called oxidation-reduction or
redox rxns
Oxidation States
- system of keeping track of e - in redox rxns by assigning
charges to various atoms in a compound
- also called oxidation numbers
- oxidation states are NOT actual charges
Rules for Assigning Oxidation States
1 – the oxidation number of an atom of an uncombined
element is zero
2 – the oxidation number of any monatomic ion equals its
ionic charge
3 – in compounds, the oxidation number of many elements
corresponds to the elements position on the periodic
table
a. Alkali Metals are ALWAYS +1
b. Alkaline Earth Metals are ALWAYS +2
c. Aluminum is ALWAYS +3
d. Fluorine is ALWAYS -1 (b/c it is the most
electronegative atom)
e. Hydrogen has an oxidation number of +1 when
combined with nonmetals
f. Oxygen has an oxidation number of -2, unless it is
in peroxide in which case oxygen is -1 or it is
paired with Fluorine
4 – the oxidation numbers of elements in compounds are
written per atom
5 – the algebraic sum of the individual oxidation numbers
of all the atoms in the formula for a compound is zero
6 – the algebraic sum of the individual oxidation numbers
of all the atoms in the formula for a polyatomic ion is
equal to the charge of the ion
7 – in binary compounds, the element with the greatest
electronegativity is assigned a negative oxidation
number equal to its charge as an anion in its ionic
compounds
Oxidizing & Reducing Agents
- in redox rxns, e- are lost and gained at the same time
- the # lost must equal the # gained (law of conservation
of charge)
- the substance that gives up e - is the reducing agent
- the reducing agent causes reduction by being oxidized
- the substance that gains e - is the oxidizing agent
- the oxidizing agent causes oxidation by being reduced
- many strong oxidizing agents contain an atom with a
high electronegativity
- also many strong oxidizing agents contain an atom with
a high positive oxidation # of +5, +6, or +7 b/c they
easily accept electrons
- common reducing agents include active free metals b/c
they have low electronegativities and tend to give up
their valence e- easily
- other reducing agents include substances that burn
readily such as hydrogen and organic substances
Balancing Redox Equations
Half-Reaction Method
- half-rxns are equations that have e - as reactants or
products
- two half rxns are always devised: one to describe
reduction and one to describe oxidation
- two rules are used to write and balance half-rxns:
- always write the formulas of molecules and ions as
they actually occur
- in aqueous sol’n, H +, H2O, and OH- are available to aid
in balancing
- in acid sol’n, H+ and H2O are available
- in basic sol’n, OH - and H2O are available
- each half-rxn and redox rxn can be balanced three ways:
by electrons, by total charge and by atoms
- any two of these methods are sufficient to give an
equation overall balance
- the two half-rxns must ultimately be added
- the two rxns must show that the # e - lost equals the # e gained, so some multiplication may be required
To balance using half rxns – acidic solution:
1 – Separate the rxn into an oxidation half rxn and a
reduction half rxn
2 – Assign oxidation # to all atoms in the equation, writing
them above or below the chemical symbol for each
element
3 – Balance oxidized element and the reduce element
numerically
4 – Balance the oxygens by adding water
5 – Balance the hydrogens by adding H +
6 – Add the two rxns together, canceling as needed
To balance using half rxns – basic solution:
1 – Approach just like balancing in acidic solution
2 – Add the same # of OH- to the side with the H + ions, as
H+ions, this will make H 2O
3 – Cancel H2O so that in any given half rxn, the water
only appears on one side of the equation
4 – Add the two rxns together, canceling as needed
Oxidation Number Method
- for more complex redox rxns:
- Remember: # e - lost in oxidation = # e - gained in
reduction
1 – Assign oxidation # to all atoms in the equation, writing
them above or below the chemical symbol for each
element
2 – Identify the element oxidized and the element reduced
and determine the change in oxidation # of each
3 – Connect the atoms that change oxidation # by using a
bracket and write the change in oxidation # at the
midpoint of each bracket
4 – Choose coefficients that make the total increase in
oxidation # equal to the total decrease in oxidation #
5 – Balance the remaining elements by inspection, and
then check the final equation
- a rxn occurring in acidic aqueous sol’n can be balanced
by adding H2O and H+ to either side of the equation
- add H2O to the side of the equation that needs oxygen
and H+ to the side that needs hydrogen atoms
Electrochemistry
- study of the interchange of chemical and electrical
energy
- deals with electricity-related applications of redox rxns
- Remember: Redox rxns involve a transfer of e -
- can be used to produce an electric current
- this transfer also involves a release of heat
- separating the oxidizing agent from the reducing agent is
key in harnessing the energy from redox rxns
- if not separated e - exchange occurs directly with no
useful work obtained
Electrochemical Cells
- devices that use redox rxns to either produce or use
electricity
- voltaic cells – electrochemical cells that produces
electricity as a result of a spontaneous redox rxns
- also called galvanic cells
- the transfer of e - takes place through an external
pathway rather than directly between reactants
- electrolytic cells – electrochemical cells that uses
electrical energy to drive nonspontaneous redox rxns
Voltaic Cells
- Luigi Galvani first credited with discovering electricity
in the late 1700’s when he noticed twitching in the limbs
of animals he was dissecting when he touched them with
tools made of different metals
- In 1800, Alessandro Volta built the first batter by
alternating disks of dissimilar metals separated by pieces
of leather that had been soaked in salt water
- voltaic cells can be visualized as consisting of two half
cells
-
-
-
- one corresponding to the oxidation half-rxn and one
corresponding to the reduction half-rxn
electrodes – two solid metals are connected by an
external circuit
anode – the electrode at which oxidation occurs
cathode – the electrode at which reduction occurs
batteries are marked with (+) and (–) signs to indicate
positive and negative electrodes
- (–) indicates the anode b/c e - are released from the
electrode
- (+) indicates the cathode b/c e - are attracted to the
electrode
e- move spontaneously from anode to cathode
a salt-bridge is used to keep the # of positive and
negative ions equal in each half cell and is essential to
complete the electrical circuit
- without flow of ions within the cell, charge will begin
to accumulate in each half cell
- flow of e- between electrodes will stop almost
instantaneously if charge builds up
the salt-bridge is a tube or porous clay barrier that allows
ions to move from one half cell to another but prevents
the sol’n from mixing totally
Voltmeter
Anode
(-)
1.10V
Cathode
(+)
Salt-Bridge
Cell Potential
- e- move from the electrode with higher electrical
potential energy to the electrode with lower electrical
potential energy
- cell potential – represents the difference in the electrical
potential energy between two electrodes of the cell
- measured in volts
- also called cell voltage
- by convention, the potential associated with each
electrode is chosen to be the potential for reduction to
occur at that electrode
- standard electrode potentials are tabulated for reductions
rxns
- called standard reduction potentials
- standard reduction potentials of all electrodes are
measured against one standard electrode, the standard
hydrogen electrode (SHE)
-
- the potential of the SHE is defined as zero
the more positive the standard reduction potential, the
more readily that species is reduced
the more negative the standard reduction potential, the
harder the species is to reduce
o
the cell potential, Ecell
is calculated by
o
o
o
E

E

E
cathode
anode
- cell
to identify the anode and cathode:
- the more positive standard reduction potential 
reduction  cathode
- the more negative standard reduction potential 
oxidation  anode
Batteries
- different kinds of batteries are based on different redox
rxns
- those that cannot be recharged use redox rxns that cannot
be easily reversed (primary batteries)
- batteries that can be recharged use an outside power
source to reverse electrode rxns (secondary batteries)
Dry Cell Batteries
- a voltaic cell in which the electrolyte is a paste
- in common dry cell batteries, the paste consists of ZnCl 2,
MnO2, NH4Cl, and H2O
- a zinc container, which serves as the anode, is separated
from the other chemicals by a liner of porous paper that
acts as a salt-bridge
- a graphite rod is in the center of the paste to serve as an
inexpensive cathode
- the battery will continue to produce e - until the reduction
product ammonia comes out of its aqueous sol’n as a gas
- advantages: low cost
- disadvantages:
- not rechargeable
- if current drawn rapidly from cell, the voltage drops
from its normal 1.5V, as reactants close to the cathode
are depleted
- Zn reacts directly with the slightly acidic ammonium
ions, causing the cell to run down even when current is
not being drawn
- leads to poor shelf-life
- putting batteries in the fridge will extend life by
decreasing the rate of redox
Alkaline Dry Cell
- modification of the ordinary dry cell
- KOH replaces NH4Cl in the paste
- powdered zinc is used in the paste that surrounds the
anode instead of zinc container as the anode
- the zinc paste is separated from the alkaline paste by
porous paper or a fabric barrier
- pastes are contained in a steel jacket and have a brass
current collector
- advantages:
- longer shelf-life
- maintains a steady voltage under high current loads
- generates about 50% more total energy than dry cells
of the same size
- disadvantages: higher cost
Nickel-Cadmium Batteries
- rechargeable voltaic cells
- also called NiCad batteries
- anode is cadmium, cathode is NiO 2 paste
- products formed in this redox rxn are insoluble and
adhere to the electrode surfaces which allows for
recharging
- advantages: lightweight and produces a constant voltage
during discharge
- disadvantages:
- if discharged partially and then recharged, it develops
the tendency to need recharging after only a short use
- must be disposed of properly b/c of the toxicity of
cadmium and its compounds
Lead Acid Storage Batteries
- standard automotive battery
- most contain 6 cells that generate about 2V each for a
total output of 12V
- called lead storage battery b/c its electrodes are
composed of alternating sheet so lead and lead dioxide
- the electrolyte is sulfuric acid
- during discharge, the lead is the anode and the lead
dioxide is the cathode
- the lead and lead dioxide are separated from each other
by porous spacers
- the lead sulfate formed during discharge is insoluble and
adheres to the electrode surfaces, which makes it
possible to reverse the rxns and therefore recharge the
battery
- advantages: inexpensive, reliable at low temps, can be
recharged thousands of times, provides a large initial
supply of energy, long shelf-life
- disadvantages: mass and must be disposed of properly
Fuel Cells
- a voltaic cell in which a fuel is continuously supplied
form an external reservoir to the cell
- very efficient at converting chemical energy into
electrical energy (~90%)
- most common: hydrogen-oxygen fuel cell that powers
the space shuttle
- simply the oxidation of H 2 to produce H2O
- hydrogen flows into the anode and is oxidized
- oxygen flows into the cathode and is reduced
- cell contains KOH, which supplies OH - ions that
participate in the electrode rxns
- advantage: b/c the fuel for the cell is provided from an
outside source, fuel cells never “run down” they keep
producing energy as long as fuel is fed to them
Electrolytic Cells
- electrical energy is applied to cause nonspontaneous
redox rxns to occur
- consists of a source of direct electrical current that is
connected to electrodes, which are immersed in the
reactants
- inert electrodes are often used b/c they do not participate
in the redox rxn themselves but serve as surfaces on
which the redox rxns take place
- e- move from anode to cathode but are “pushed” through
the circuit by the external source of electrical energy
- electrolysis – the process by which electricity in an
electrolytic cell is used to bring about a nonspontaneous
chemical change
- uses: electrolysis rxns including isolation of active
metals, purifications of metals, and electroplating