The document discusses the electrochemical series and how it relates to voltage in galvanic cells. It explains that electrons flow from the metal higher in the electrochemical series to the metal lower down, creating a voltage. It provides an example of connecting magnesium, copper, and tin, and explains the sign and magnitude of the voltage based on their positions in the series.

1. The Electrochemical Series

2. V
ELECTRONS
IONS
The metal HIGHER in the
electrochemical series
LOSES electrons to the
metal lower down.
ELECTRONS flow through
the wires.
IONS flow through the
solution to complete the
circuit.
A cell

3. The electrochemical series (ECS)
Aluminium
Zinc
Iron
Tin
Lead
Electrons flow from one metal to another to
create a voltage
Imagine you connect different metals to
TIN...
Mg Cu Sn
1.2 V -0.5 V 0 V

4. Mg Cu Sn
1.2 V -0.5 V 0 V
+ve – the metal is
above the other
metal in the ECS
-ve – the metal is
below the other
metal in the ECS
0 – the metal is the
same as the other
metal
large number –
large gap
between the
metals in the ECS
small number –
small gap
between the
metals in the ECS

5. Magnesium
Aluminium
Zinc
Iron
Tin
Lead
Copper
Above tin
Large gap
Below tin
Small gap

7. Introduction to Electrochemistry
 An electric cell converts chemical energy into electrical energy
 AlessandroVolta invented the first electric cell but got his inspiration from Luigi
Galvani. Galvani’s crucial observation was that two different metals could make the
muscles of a frog’s legs twitch. Unfortunately, Galvani thought this was due to some
mysterious “animal electricity”. It wasVolta who recognized this experiment’s potential.
 An electric cell produces very little electricity, soVolta came up with a better design:
 A battery is defined as two or more electric cells connected in series to
produce a steady flow of current
 Volta’s first battery consisted of several bowls of brine (NaCl(aq))
connected by metals that dipped from one bowl to another
 His revised design, consisted of a sandwich of two metals
separated by paper soaked in salt water.

8. 8 © Alexis Kwasinski, 2012
Li-ion batteries
• Positive electrode: Lithiated form of a transition metal oxide (lithium cobalt
oxide-LiCoO2 or lithium manganese oxide LiMn2O4)
• Negative electrode: Carbon (C),
usually graphite (C6)
• Electrolyte: solid lithium-salt electrolytes
(LiPF6, LiBF4, or LiClO4)
and organic solvents (ether)
http://www.fer.hr/_download/repository/Li-ION.pdf
discharge

9. 9 © Alexis Kwasinski, 2012
Li-ion batteries
• Chemical reaction (discharge)
• Positive electrode
• Negative electrode
•Overall
• In the above reaction x can be 1 or 0
• With discharge the Co is oxidized from Co3+
to Co4+
. The reverse process
(reduction) occurs when the battery is being charged.
LiCoO2 Li1-xCoO2 + xLi+
+ xe-
xLi+
+ xe-
+ 6C LixC6
Through electrolyte
Through load
LiCoO2 + C6 Li1-xCoO2 + C6Lx

10. 10 © Alexis Kwasinski, 2012
Li-ion batteries
• Advantages with respect to lead-acid batteries:
• Less sensitive to high temperatures (specially with
solid electrolytes)
• Lighter (compare Li and C with Pb)
• They do not have deposits every charge/discharge
cycle (that’s why the efficiency is 99%)
• Less cells in series are need to achieve some given
voltage.
• Disadvantages:
• Cost

11. What is a hydrogen fuel cell?
•Hydrogen fuel cells (HFCs) are a type of
electrochemical cell.
•HFCs generate electricity by reduction
and oxidation reactions within the cell.
•They use three main components, a
fuel, an oxidant and an electrolyte.
•HFCs operate like batteries, although
they require external fuel.
•HFCs are a thermodynamically open
system.
•HFCs use hydrogen as a fuel, oxygen as
an oxidant, a proton exchange
membrane as an electrolyte, and emit
only water as waste.

12. How do they work?
•Fuel (H2) is first transported to
the anode of the cell
•Fuel undergoes the anode
reaction
•Anode reaction splits the fuel
into H+
(a proton) and e-
•Protons pass through the
electrolyte to the cathode
•Electrons can not pass through
the electrolyte, and must travel
through an external circuit which
creates a usable electric current
•Protons and electrons reach the
cathode, and undergo the
cathode reaction

13. Chemistry behind the technology
Oxidation
At the anode of the cell, a
catalyst (platinum powder)
is used to separate the
proton from the electron in
the hydrogen fuel.
Anode half-reaction:
2H2  4H+
+ 4e-
Eo
= 0.00V
Reduction
At the cathode of the cell, a
second catalyst (nickel) is used to
recombine the protons,
electrons, and oxygen atoms to
form water.
Cathode half- reaction:
4H+
+ O2 + 4e-
 2H2O
Eo
= 0.68V
In electrochemistry, the Eo
cell value (energy) of a fuel cell is equal to the Eo
of
the cathode half-reaction minus the Eo
of the anode half-reaction. For a
hydrogen fuel cell, the two half reactions are shown above. So to calculate the
energy of one fuel cell, we need to subtract the anode energy from the
cathode energy. For a HFC, the Eo
cell = 0.68V – 0.00V which equals 0.68V

14. Uses of hydrogen fuel cells
There are many different uses of fuel cells being utilized right now. Some of these
uses are…
•Power sources for vehicles such as cars, trucks, buses and even boats and submarines
•Power sources for spacecraft, remote weather stations and military technology
•Batteries for electronics such as laptops and smart phones
•Sources for uninterruptable power supplies.

15. Voltaic Cells (Galvanic Cell)
 A device that spontaneously produces electricity by redox
 Uses chemical substances that will participate in a spontaneous redox reaction.
 The reduction half-reaction (SOA) will be above the oxidation half-reaction (SRA) in the activity
series to ensure a spontaneous reaction.
 Composed of two half-cells; which each consist of a metal rod or strip immersed in a
solution of its own ions or an inert electrolyte.
 Electrodes: solid conductors connecting the cell to an external circuit
 Anode: electrode where oxidation occurs (-)
 Cathode: electrode where reduction occurs (+)
 The electrons flow from the anode to the cathode (“a before c”) through an electrical
circuit rather than passing directly from one substance to another
 A porous boundary separates the two electrolytes while still allowing ions to flow to
maintain cell neutrality
 Often the porous boundary is a salt bridge,
containing an inert aqueous electrolyte
(such as Na2SO4(aq) or KNO3(aq)),
 Or you can use a porous cup containing
one electrolyte which sits in a container of a
second electrolyte.

16.  Voltaic cells can be represented using cell notation:
 The SOA present in the cell always undergoes reduction at the cathode
 The SRA present in the cell always undergoes oxidation at the anode
Voltaic Cells (Galvanic Cells)
The single line represents a phase
boundary (electrode to electrolyte)
and the double line represents a
physical boundary (porous boundary)
The single line represents a phase
boundary (electrode to electrolyte)
and the double line represents a
physical boundary (porous boundary)