Transcript 10~chapter_8_electro..

Electrochemical Engineering
& Alternate Energy
How to store
some electricity
for later use
Layout of this lecture
• Green energy
– Load leveling
• Electrochemical principles
– Anodes and cathodes
– Half cells and simple electrochemical cells
• Fuels Cells
• Ragone plot and battery capacities
Why electrochemical engineering?
• Batteries and fuel cells are deeply embedded in
“Green Energy” because solar and wind energy
systems need to store electrical energy
• It’s also a hot topic because some vehicles use
batteries for propulsion such as in hybrid cars
and trucks
• Electrochemical engineering is the study of what
happens inside batteries, fuel cells and
‘ultracapacitors’
– So be prepared for a little more chemistry!
Green Energy
• “Green energy” refers to renewable energy
supplies that do not spew forth greenhouse
gases nor toxic impurities
– Wind and solar energy are two favored sources
• Big disadvantage #1: Only works when the wind blows
or the sun shines
• Big disadvantage #2: May make too much electricity
exactly when you don’t need it.
• Solution: Store the electrical energy until you
do need it.
Load leveling
– These may be very
large batteries!
Power demand
Evening power
demand
Time
Power produced
• Electricity not used
when generated, nor
available when needed
if the sun or wind go
down
• Solution: Battery
storage
Morning power
demand
Noon peak
generation
Time
Figure
2: Mismatch
daily power
production
Mismatch
of greenofelectric
power
and use and
Solar or wind need a lot of equipment
Batteries
Why do batteries work?
• Matter is inherently electrically charged.
–Simplest case is ionic bonding (i.e., attraction) in
compounds such as common table salt: Na+Cl- in which
Coulombic forces hold together positively charged sodium
ions and negatively charged chlorine ions. The force between
these ions is:
2
e
F 2
kr
where e is the charge on an electron and r is the interionic
distance. k is the dielectric constant, which is ~80 for water.
Electrochemistry
• When Na+Cl- dissolves in water with k ~ 80,
the forces between ions lessen allowing free
ions as Na+ and as Cl– Ion-containing solutions are called “Electrolytes”
– Overall ion-containing solutions are electrically
neutral
– Locally ion-containing solutions have both charged
species at short distances to each other
– Batteries use these ions when they can be
separated
Electrolytes, Anodes and Cathodes
• Electrodes are classified
as anodes and cathodes
• Anodes are “sources” of
electrons, and cathodes
are “sinks” for electrons
–
–
–
–
e- = electrons
E = Electrolyte
C = Cathode
A = Anode
e- flow
-
+
C
A
E
Discharge through a load
Electrolytes, Anodes and Cathodes
• Anodes are “sources”
of electrons, and
cathodes are “sinks”
for electrons
–
–
–
–
e- = electrons
E = Electrolyte
C = Cathode
A = Anode
e- flow
+
C
A
E
Charge through an source
Electrolytes, Anodes and Cathodes
• Beware Franklin’s
error!
–
–
–
–
e- = electrons
E = Electrolyte
C = Cathode
A = Anode
Conventional Current
flow
-
+
C
A
E
Discharge through a load
Lead-Acid Batteries
• These are the batteries you find in a car
– Both electrodes are based on lead, Pb one with a
PbO2 coating
– The electrolyte is sulfuric acid written H2SO4,
which dissolved in water is 2H+ /SO=4 (the sulfate
ion has two electrons/molecule)
– The principle anodic reaction is: Pb  Pb++ + 2e– The two electrons flow through the external circuit
to the cathode on which:

PbO2  4H  SO

2e
 PbSO4  2H2 0
4
Lead-Acid Batteries
• Product of reaction is PbSO4 which precipitates
during discharge and dissolves during charging.
• The anodic voltage at the anode is 0.36V above a
reference cell and the cathodic is 1.69 V below.
• Overall cell voltage = ~2.0 V
C
A
Anodic
0.36 V
Reference
cell
E
Cathodic
1.69V
Can you power a car using batteries alone?
Property
Mass
Energy
Mass energy storage density
Volumetric energy storage
density
Power
Lead-acid
battery
25. kg
3,000 kJ
120 kJ/kg
250 kJ/liter
Gasoline
25. kg
1.2  105 kJ
46,500 kJ/kg
34,400 kJ/liter
5 kW
Typically >100 kW
• This battery is too heavy, contains too little
energy, and delivers too little power – that’s why
hybrids are a popular substitute.
The Electrochemical Series
• Cell voltage set by the tendency to transfer
electrons
Half Cell Chemistry
Li+ + e-  Li(s)
Na+ + e-  Na(s)
Mg++ + 2e- Mg(s)
Zn++ + 2 e-  Zn(s)
Fe++ + 2 e-  Fe(s)
Ni++ + 2e-  Ni(s)
2H+ + 2 e-  H2(g)
Cu++ + 2 e-  Cu(s)
Cu+ + e-  Cu(s)
Ag+ + e-  Ag(s)
Pd++ + 2e-  Pd(s)
Potential in volts
–3.05 V
–2.71 V
–2.37 V
–0.76 V
–0.44 V
–0.25 V
0.00V (Hydrogen ½ cell is defined as zero)
0.34 V
0.52 V
0.80 V
0.95 V
The Electrochemical Series
• These half-cell reactions in principle are
reversible. The more negative the more they
want to lose electrons; the more positive the
more to gain them
– This determines how cells will behave:
Daniell cell
• The electrolytes are
ZnSO4(aq) with a Zn
anode and CuSO4(aq)
with a Cu cathode.
Write down the
reactions in each half
cell and explain what
happens in the salt
bridge. What is the
voltage?
Daniell cell
• In bulk aqueous solution , Zn++ and SO4= must
be in balance with each other.
ZnSO4  Aq   Zn  SO


4
• But the anode is also dissolving and thus yields
some locally extra Zn++ ions according to:

Zn(s)  Zn  2e

(V  0.76V)
Daniell cell
• Electrons from the anode flow through the external
circuit precipitating copper at the cathode:
Cu  2e  Cu  s   0.34V


• We have removed copper ions from solution;
therefore there must be a corresponding
reduction in SO4= ions from the electrolyte. They
must move into the salt bridge to exactly counteract
the Zn++ ions from the anodic side.
The cell potential is equal to:
(+0.76 V) + (+0.34 V) = +1.10 V.
Electroplating
• An electroplater wants to coat a 10.0 cm by
10.0 cm copper plate with 12.5 micrometers of
silver. How many electrons must pass in the
external circuit? How many coulombs are
passed? If the plating takes 1,200. s what’s the
electrical current in amperes in the external
circuit?
Electroplating
• Know: Atomic mass of
Ag is 108 kg/kmol. Its
density is 10,500 kg/m3.
Avogadro’s number
(NAv) is 6.02  1026
atoms/kmol. What we
call current is nothing
but the rate of flow of
electrons, so 1.00 A =
1.00 coulomb/s and one
electron carries –1.60 
10-19 coulombs.
Conventional
current flow
+
-
Electron
flow
Ag
Cu
Ag+, NO3- , H2O
Electroplating
• The reaction at the anode is Ag(s) → Ag+ + eand the reaction at the cathode is Ag+ + e- →
Ag(s); hence one atom of silver dissolves at
the anode and one atom of silver is deposited
at the cathode. For each atom of silver
dissolving at the anode and depositing at the
cathode, one electron must circulate in the
external circuit.
Electroplating
2
12.5×10-6 ×100.×10,500
2
3




cm
m




Mass =
μm
cm
kg/m
/


2

 
 μm  
m 
1.00×104
= 1.31×10-3 kg
• Next convert to kmols:
 kmols 
1.31  10 3
5
kg 
kmols 

1
.
22

10
kmols

108
 kg 
Electroplating
• Next convert to atoms of Ag(s):
Ag atoms deposited 
5
1.22 10  6.02 10
26
 7.32 10 21[kg][ atoms / kg]
 7.32 10 atoms
21
Electroplating
• Next convert to mass of Ag(s):
7.32 10 21 1.6010 19[e  ][coulombs/e - ]
 1.17 10 coulombs
3
1.17 10
Hence current 
[C/s ]  0.975A
1200.
3
Fuel Cells
• Fuel cells are just continuously refueled
batteries.
– They will not discharge while
electrochemical fuel is being fed to them.
• Most fuel cells depend on “Proton Exchange
Membrane” or “PEM” to catalyze electrode
reactions
2 H 2  4 H   4e  Anodic Rxn




4 H  4e  O2  2 H  2OH  2 H 2 0
Cathodic Rxn
Fuel Cells
http://aq48.dnraq.state.ia.us/prairie/images
Fuels cells
• Note: Only H2 and O2(i.e.,air) in and only H2O
out.
• Cell voltage is 1.23 V for overall rxn H2 + O2
= H2O
– Apparently no green house gas pollution!
– Unfortunately to make H2 needs copious CO2
The Ragone Chart
• Batteries must supply both energy and power
– Typically batteries supply current measured at
mA/cm2 of electrode area at a few volts
– The more electrode area, the greater the current;
this may be internal area packing or simply more
cells placed in parallel
– The more cells in series the higher the voltage
• But it’s the application that demands whether
the cell can deliver both enough energy (say
mileage between charges on an electric car)
and power (say to pass another car)
The Ragone Chart
• The best measures of energy and power
efficiency are their mass densities: e = E/Wt
(Whr/kg) and p = P/Wt (W/kg)
• The energy density delivered by a power
source for a time t is simply e = p  t.
– Take log base 10 of this equation:
• log10 e = log10 p + log10 t
– Plot the log10 energy density of a battery vs. log10
power density for the same battery and you get the
Ragone Plot
The Ragone Chart
Modified from a graphic of Maxwell Technologies:
http://www.maxwell.com.
The Ragone Chart
• Very convenient way to compare different
electrochemical sources
– Ideally you want to have your cake and eat it by
being in the upper right corner
– Reality shows what can be achieved by competing
electrochemical sources
• “Ultracapacitors” are storage devices that can
store thousands of times the energy capability of
an electrical capacitor.
Ragone Chart
1,000
rs
u
o
10h
our
h
1
100
Energy, Whr/kg
• That that the form
of log10 e = log10 p
+ log10 t is y = mx
+ c and has a slope
of m =1 given the
log10 scaling in
chart.
• The battery’s
discharge time is
given by e/p
s
0
6
3
10
s
36
1
0.1
s
3.6
s
m
6
6s
3
.
0
3
0.01
10
100
Power, W/kg
1,000
10,00
Summary:
• Green energy and load storage and leveling
• Electrochemical series
– Simple electrical cells
– Simple electrochemistry
• Principles of fuel cells
• Ragone chart to characterize