Transcript Slide 1

ELECTROCHEMISTRY REVIEW
OXIDATION REDUCTION REACTIONS
Oxidation - A process in which chemical entities
lose e Cu was oxidized to Cu2+
OXIDATION REDUCTION REACTIONS
Reduction – A process in which chemical entities
gain e Each Ag+ was reduced to Ag
OXIDATION REDUCTION REACTIONS

reaction in which one reactant is oxidized and
the other is reduced is called a redox reaction.
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Therefore, in any redox reaction, the number of
e- lost must equal the number of e- gained. It
must balance.


If we look at the transfer of electrons in this
reaction we will see that Cu looses e- and Ag
gains eCu(s) + 2 Ag+ (aq)  2 Ag(s) + Cu2+ (aq)
Cu becomes Cu2+, a loss of 2 e Each Ag+ becomes Ag, a gain of 2 e

TIP:

To remember which is oxidation and which is
reaction you can use of these memory aids.
 LEO says GER
 Lose Electrons Oxidation
 Gain Electrons Reduction

OR
OIL RIG
 Oxidation Is Loss – Reduction Is Gain

OXIDIZING AND REDUCING AGENTS
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A substance that removes electrons from
another substance is known as a oxidizer or a
Oxidizing agent
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A substance that gives electrons to another
substance is known as an reducer or a
reducing agent.
STEPS TO IDENTIFYING OXIDATION AND
REDUCTION
Step 1) Identify repeating entities
 Write the total ionic equation
 Eliminate ions common on both sides
Step 2) Label charges
 Uncombined have a charge of 0
Step 3) Identify loss and gain of electrons
EXAMPLE
Zn(s) + CuSO4(aq)  ZnSO4(aq) + Cu(s)
Step 1)
Zn + Cu2+(aq) + SO42-(aq)  Zn2+(aq) + SO42-(aq) + Cu(s)
Step 2)
 Zn0(s) + Cu2+(aq)  Zn2+(aq) + Cu0(s)
Step 3)
 Zn becomes Zn2+, loss of 2 e- (oxidation)
 Cu2+ becomes Cu, gain of 2e- (reduction)
Oxidation Numbers
 The sum of all of the oxidation numbers in a
molecule must be zero because the molecule
does not have an overall charge.
 If it was a polyatomic ion it would have a
charge but the oxidation numbers can still be
calculated.
Rules for Assigning Oxidation
#’s
I) The oxidation number of a atom in an
uncombined element is always Zero.
 Example: Na, K, O2, H2, Cl2, S8, Li,
Rules for Assigning Oxidation
#’s
II) The oxidation number of a simple ion is the
charge of the ion.
 Example: Ca2+ = + 2
 Na+ = + 1
 Cl- = - 1
Rules for Assigning Oxidation
#’s
III) The oxidation number of Hydrogen in most
compounds is +1.
 Example: H2O, H2SO4, NH3, H = +1
 Except in metal hydrides: NaH, H = -1
Rules for Assigning Oxidation
#’s
IV) The oxidation number of oxygen in most
compounds is -2
 Example : MgO, HNO3, OH- = -2
 Except in peroxides: H2O2, = -1
Rules for Assigning Oxidation
#’s
V) The oxidation number of group 1 elements is
+1
 Example: Na, Li, = + 1
 The oxidation number of group 2 elements is
+2
 Example: Mg, Ca = + 2
Rules for Assigning Oxidation
#’s
VI) The sum of oxidation numbers in a
compound must = 0
 Example: H2O = 2 (+1) + (-2) = 0
Rules for Assigning Oxidation
#’s
 The sum of oxidation numbers in a
polyatomic ion must equal the charge of the
ion.
 Example: OH- = (-2) + (+1) = -1
Identifying Redox Reactions
Using Oxidation Numbers
 Not all reactions are redox
 By using oxidation numbers, you can tell if a
reaction is redox. A redox reaction will have a
change in oxidation numbers where a reaction
that is not redox will not.
 Example: AgNO3 + HCl  HNO3 + AgCl
Ag+ + NO3- + H+ + Cl-  H+ + NO3- + Ag+ + Cl-
Metal Activity Series
 The Activity series is an arrangement of
metals in order of their tendency to react
(become oxidized).
 The most reactive metals (most easily
oxidized) are at the top of the list.
 The least reactive are at the bottom.
Metal Activity Series
 This allows us to predict whether a single
displacement reaction will occur not. For it to
occur, the metal must be higher on the list
than the metal in the compound that it is
trying to displace.
Activity series
Metals
Lithium
Potassium
Calcium
Sodium
Aluminum
Zinc
Chromium
Iron
Nickel
Tin
Lead
Hydrogen *
Copper
Mercury
Silver
Platinum
Gold
Decreasing Activity
Halogens
Fluorine
Chlorine
Bromine
Iodine
 The Activity series is used to predict the
products of single displacement reactions.
 A + BC  AC + B
 In general, an element that is higher on the
activity series will displace an element that is
lower. The lower element is, thus left as a
pure metal.
Galvanic cell
 This apparatus is called a galvanic cell.
 A device that converts chemical energy from redox reactions into
electrical energy.
Galvanic cell
 This apparatus is called a galvanic cell.
 A device that converts chemical energy from redox reactions into
electrical energy.
Galvanic cell
 A Galvanic cell is a spontaneous reaction
 A reaction that proceeds on its own without outside assistance
(energy).
 The oxidation of zinc and reduction of copper occur in
separate beakers called half cells.
 One of the two compartments in a galvanic cell
 Composed of an electrode and a electrolytic solution.
Galvanic cell
 The metal in each beaker is called a electrode.
 A solid electrical conductor where the electron transfer
occurs.
Galvanic cell
 Each electrode has a special name
 Anode – The electrode where oxidation occurs (think of
anion, a negative ion)
 Cathode – The electrode where reduction occurs (think
cation, a positive ion)
 REDCAT (Reduction Cathode)
 Metals and non-reactive conductors such as
graphite are often used as conductor electrodes.
How does a salt bridge work?
 How does a salt bridge work? (do not look at text, they
used the wrong metals)
 The purpose of the salt bridge is to provide ions to prevent
charge from building up.
 In a way it is much like simple diffusion.
How does a salt bridge work?
 Every time a zinc atom is oxidized to an ion it would make
the solution more positive, which would then stop the
reaction. The nitrate from the salt bridge moves in and
balances it.
 On the other side, every time a copper ion is reduced it
would make the solution more negative and stop the
reaction. A sodium on then moves in a negative nitrate ion
leaves. This allows the circuit to continue without a build up
of charge.
Cell Reactions
 The chemical equation for this reaction can be broken down into 2
parts, called half cell reactions.
 Anode half –reaction = Zn(s)  Zn2+(aq) + 2e- (oxidation)
 Cathode half –reaction = Cu2+(aq) + 2e-  Cu(s) (reduction)
Cell Reactions
 Therefore, as the cell operates the mass of the zinc electrode decreases
and the mass of the copper electrode increases.
 Anode half –reaction
 Cathode half –reaction
 Overall cell reaction =
Zn(s)  Zn2+(aq) + 2eCu2+(aq) + 2e-  Cu(s)
Cu2+(aq) + Zn(s)  Cu(s) + Zn2+ (aq)
Corrosion
 Corrosion – The deterioration of metals as a result of
oxidation.
 A few metals such as copper, zinc and aluminum form
protective coatings when they oxidize. This makes them more
corrosion resistant than other metals that are lower on the
activity series. This is why they are commonly used to coat
and protect other metals.
Rusting Iron
 Rust is produced when iron reacts with oxygen to form rust oxide.
This new compound does not stick well to the existing metal and
flakes off leaving the metal underneath it to rust. This process
continues until all the metal is gone.
The Redox Reaction of Rust
 A corroding metal is a galvanic cell in which the anode and
the cathode are found at different points on the same metal
surface.
 The metal itself is the conducting material that allows the
electrons to flow from the anode to the cathode. The anode is
normally starts due to a scratch dent or impurity.
The Redox Reaction of Rust
The Redox Reaction of Rust
Cathode
 O2(aq) + 2H2O(l) + 4e-  4OH-(aq)
Anode
 Fe(s)  Fe2+(aq) + 2e-
Factors that Affect the Rate of
Corrosion
Moisture
 Since water takes part in the reaction, it must be present for
the reaction to occur. A relative humidity of at least 40% is
needed for the reaction to take place.
Electrolytes
 When salts dissolve in water they become ions which
increase the conductivity of water. The chlorine ions also act
in a similar manner to a salt bridge in the way they offset the
increase of Fe2+ ions at the anode. The sodium ions play a
similar role at the cathode as they help to offset the negative
charge build up from the hydroxide ions.
Contact with Less Reactive Metals
 When two different metals come in contact with each other,
the more reactive metal becomes oxidized. This is why metal
fabricators must use the same type of metal when fabricating
materials to avoid corrosion.
Mechanical Stress
 Bending, shaping, or cutting metal, stresses the structure of
the metal which creates weak points. The weak points are
then prone to corrosion.
Preventing Corrosion
Protective Coatings
 The simplest method of corrosion resistance is to cover the
metal with a protective coating. Once the coating is scratched
or exposed the metal will rust, even though the rust may
appear to only be at the surface it can actually spread deep
into the metal.
Galvanizing
 The process of coating iron or steel with a thin layer of zinc.
This can be done by dipping the metal in a hot vat of molten
zinc or by electroplating. When the zinc oxidizes it forms a
tough, protective coating.
Cathodic Protection
 A form of metal corrosion prevention in which the metal
being protected is forced to be the cathode of a cell, using
either impressed current or a sacrificial anode.
Sacrificial Anode
 A form of protection where a metal that is more easily
oxidized is attached to another metal to protect it. The more
reactive metal acts as the anode, thus protecting the other
metal by making it the cathode.
Sacrificial Anode
 This method does not require complete covering of the
metal; all it needs is some sort of conductive connection that
allows it to pass electrons to the metal that needs protecting.
The sacrificial anode will need periodic replacement.
Impressed Current
 In this method, the metal needing
protection is attached to the negative
terminal of a power source making it
the cathode. Continually pumping
electrons into the cathode prevents
corrosion