Transcript Galvanic Cells – Converting chemical energy to electrical

Lesson 2
Galvanic Cells
 In the reaction between Zn and CuSO4, the zinc is
oxidized by copper (II) ions.
 Zn0(s) + Cu2+(aq) + SO42-  Cu0(s) + Zn2+ + SO42-(aq)
 Zn becomes Zn2+, a loss of 2e- (oxidation)
 Cu2+ becomes Cu, a gain of 2e- (reduction)
 Separating the zinc metal from the solution containing
copper ions, and placing a metal conductor between
the metals creates electricity. Electrons that are lost by
zinc are forced to travel through the metal conductor
to reach the copper ions. The movement of electrons is
known as electric current.
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.
Galvanic cell
 Each electrode is immersed in an electrolytic solution
that contains the same metal ions as the electrode.
 Zinc is in a zinc nitrate solution
 Copper is in a copper (II) nitrate solution
Galvanic cell
 Half cells are connected by a salt bridge.
 Concentrated solution of electrolyte, it should not
react with the other chemical in the reaction.
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.
Questions
 Page 400 # 1-7
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)
Writing Cell Reactions
 Step 1) Establish elements oxidized and reduced
 Step 2) Balance charges
 You may need to multiply the half reactions by a
common multiple to balance out the number of
electrons.
Example:
 A strip of silver in a solution of silver nitrate , and a
copper strip in a solution of copper (II) nitrate
Step 1) Copper is higher on the activity series so it will
be oxidized.
 Anode half-reaction
Cu(s)  Cu2+(aq) + 2e Cathode half-reaction
Ag+(aq) + e-  Ag(s)
 -note: the number of e- lost by copper is not equal to
the number of e- gained by silver.
Step 2) Multiply both sides of the cathode reaction by 2 so that
the number of e- is equal to the e- gained.
Cu(s)  Cu2+(aq) + 2e Cathode half-reaction 2 Ag+(aq) + 2e-  2Ag(s)
_
 Overall cell reaction Cu(s) + 2 Ag+(aq)  Cu2+(aq) + 2Ag(s)
 Anode half-reaction
 Practice – page 397
 Activity 5.8 – Designer Cells
Corrosion
Corrosion

Corrosion – The deterioration of metals as
a result of oxidation.

With the exception of a few un-reactive
metals, most metals are found as minerals
and not in elemental form. Without protection
against corrosion, most metals are oxidized
by their environment.

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.
Example

Aluminum reacts with oxygen to form
aluminum oxide. It is one of the hardest
compounds known; it is chemically stable
and heat also resistant. It is a great
protectant because it forms as quickly as it is
scraped off
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-
The Redox Reaction of
Rust
In the presence of oxygen and moisture, iron
oxidizes releasing 2 e The electrons then travel through the metal
to the cathode, where they are used to
reduce oxygen molecules. Water has two
purposes in this reaction; it acts as a salt
bridge for ions to flow and it also takes part
in the reaction with hydroxide ions and the
iron (III) ions to form rust.

The Redox Reaction of
Rust
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.

Page 416 # 1-6
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.
Protective Coatings
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

Page 425 # 1-7