Download Redox Reactions

RedOx Reactions
&
Electrochemistry
Outline and Section
Review 8.3
Usefulness of RedOx
• Making Batteries
• Corrosion Control
• Isolating Elements from
Compounds
• Used for dyes and pigments
Reduction + Oxidation =
RedOx
RedOx
Reduction – Gaining of
electrons
Oxidation – Loss of electrons
OIL RIG
Oxidation is Losing Electrons
Reduction is Gaining Electrons
Reduction and oxidation
always happens in pairs
thereby
Reduction / Oxidation
2Al + 3CuCl2  2AlCl3 + 3Cu
0
2Al

2+
3Cu
+
3+
2Al
6e
+

6e
0
3Cu
Lab – Single Replacement
Section Review 19.1
Questions 2, 3 (b-d), 5
Electrochemistry
OPTIONAL READING:
Section 20.1, pg 658 and
pg 662-665
WS – Voltaic Cell
We will be discussing how to
produce electrical energy from a
RedOx Reaction just as is done in
a common battery.
Batteries are a bank of voltaic cells.
A battery’s usefulness is measured thru
Voltage (or Electric Potential)
& Wattage / Current.
We will discuss how to calculate the
expected voltage and how to improve
on the voltage
Electric Potential (or Voltage) is the
driving force behind a transfer of
electrons.
Voltage can
be thought of
as Electrical
Pressure.
Voltage is created by a
RedOx Reaction
RedOx of Two Metals
• Redox Reactions half reactions

– Oxidation:
Zn0
– Reduction:
Cu2+ +
Zn2+
+
2e
2e-  Cu0
The repulsion of the electrons at the
site of oxidation creates a flow of
electrons towards the site of reduction.
If we separate the metals and
channel the electrons through a
circuit we can use the flow of
electrons (electricity)
Electrodes
• The Zinc and Copper Strips are our
Electrodes.
• Electrodes are conductors that make a
connection with the nonmetallic parts of the
solution.
• Electrodes are also the location of oxidation
and reduction
Electrodes
• ANODE is the site of OXIDATION
– In a battery the anode is negative because it
is the site where electrons are released.
• CATHODE is the site of REDUCTION
– In a battery the cathode is positive because
it is the site where electrons are absorbed.
Simple Battery
Activity
Cu and Zn
Do Not Dip the Alligator
Clips and Rinse them off
We have a table that is similar to the
activity series, but gives data that
can be used to calculate the possible
voltages.
It is called a
Standard Reduction Potential Table.
All the reactions on a Reduction
Potential Table are shown as
Reductions and Compared to
Hydrogen’s ability to be reduced.
How to calculate voltage
• We know that in a redox reaction one substance /
element has to gives up electrons and one absorb
electrons.
• The one that gives up electrons (oxidation) will
be lower on the reduction potential table.
• We reverse the reduction half-reaction for the
element that has a lower reduction potential
because it is actually oxidized instead of reduced.
Find the Reduction Half-Rxns and
Reverse the One That’s Oxidized
Cu2+ + 2e-  Cu0
Zn2+ + 2e-  Zn0
+0.34
-0.76
Volts
Volts
Zn is lower so…………
0
Zn

+2
Zn
+
2e
+0.76
Volts
Calculating Cell Potential
Zn0  Zn+2 + 2eCu+2 + 2e-  Cu0
+0.76 V
+0.34 V
+1.10 V
Combined force with which the
reaction takes place 1.10V
E cell = E cathode + E anode
E cell = E reduction + E oxidation
This equation assumes that the sign of the
voltage for the oxidized substance has
already been changed!
Substances will have different
electron releasing or absorbing
abilities as evidenced by the
activity series which is really an
oxidation table and the
reduction potential table.
Larger distances between metals
on reduction potential table the
larger the voltage.
Better means more voltage
(electrical potential).
Lab – Making a Battery
Lab – Building a Reduction
Potential Table
Lab – Powering an LED
Corrosion of Metals
25% of metal
production is to replace
corroded metals
RUSTING 4Fe(s) +3O2
0
4Fe
6

(g)
3+
4Fe
0
O + 12e

 2Fe2O3 (s)
+
12e
26O
Issues Corrosion Creates
•Safety
•$$$$$$
•Aesthetics / Cleanliness
Videos of Corrosion Impacts
• Aloha Airlines Flight 243
• I-35 Bridge
Galvanic
Corrosion
GALVANIC
CORROSION
GALVANIC Corrosion
• Insulating material
between Copper Statue
and Iron Scaffolding
broke down and needed
to be replaced at the
cost of $90,000,000
What is a half reaction & net equation?
• Redox Reactions half reactions
– Oxidation:
– Reduction:
Fe0  Fe2+ + 2eCu2+ + 2e-  Cu0
• Net Ionic Reaction
Fe0 + Cu+2  Fe+2 + Cu0
• Corrosion will occur between any
two substances (in the case of
galvanic corrosion – two metals)
which have a positive voltage
• Anytime two dissimilar metals are
in contact there will be corrosion.
How do you prevent corrosion ?
• Avoid contact – separation
• Paint or Plastic Insulation
• Coat with metal that makes a good
oxide layer like Cr (Stainless) or Zn
(Galvanized)
• Keep materials dry
– Salt water exacerbates the issue
Sacrificial Anode –
Prevents Corrosion
• Connect structurally important metal to a more
easily oxidized metal.
• Metal is not important to structure or function
• Must be replaced
• Ex: Zinc coated nails, zinc piece on boat motor,
bridges have zinc slabs.
Lab –
Corrosion of a Nail
What can speed or impede
corrosion of an Fe nail?
??????????
Cu Wrap
Zn Wrap
Electroplating
What is electroplating?
• Electroplating is coating one metal onto another
through precipitation of metal ions onto a
negatively charged part (anode)
• Electroplating helps with corrosion resistance
and appearance.
• Zn, Cr, Ni are coated onto Fe to prevent rust
(redox rxns) because they form an impervious
and durable oxide layer.
Chrome Plating
Chrome and Iron Electrodes in Nickel Ion
+2 + 2eo = +0.26 V
Ni

Ni
E
solution
Cu+2 + e-  Cu
Eo = +0.34 V
Fe3+ + 3e-  Fe0
Eo = -0.45 V
Etotal = 0.60 V
Cr0  Cr3+ + 3eEo = 0.75 V
Etotal = 0.30 V
Electroplating Cell
+
Power
Supply
-
Oxidation
Reduction
Cr
Cr3+
Electrodes
Cleaned
off
Cr3+
Fe
Plating
Plating Nickel is a Nonspontaneous plating reaction
Naturally Ni is reduced and Fe is oxidized
Fe3+ + 3e-  Fe0
Ni0  Ni+2 + 2e-
Eo = 0.26 V
Eo = -0.45 V
Etotal = -0.19 V
To reverse the process we would need to
add 0.19 Volts to the process.
Types of Redox Reactions
• Decomposition
• Synthesis
• Single-Replacement
– Follows activity series
(NH4)2Cr2O7 (s) 
Cr2O3 (s) + N2 (g) + 4H2O (g)
Decomp of Ammonium
Dichromate
(NH4)2Cr2O7(s)Cr2O3(s)+ N2(g) +4 H2O(g)
Heat – Exothermic Reaction
Decomposition Reaction
Self Sustaining
Electrolytic Cells
• Because reactions are merely the exchange of
electrons we can use electricity and LeChatelier’s
Principle to force reactions to proceed in
whatever direction we desire.
• Demonstration of electrolysis of H2O
– Indicators around both the cathode and anode to
detect the presence of OH- and H+
• Aluminum metal and Sodium metal
Bleaches
• Color caused by the
movement of electrons
up and down energy
levels
• Removing electrons –
OXIDATION – will
prevent this movement
The End
Confusion of terms
Reducing agents get oxidized and
Oxidizing agents get reduced
The exchange of electrons
can be real (ionization and
charges) or merely helpful
(oxidation numbers for
covalent compounds)
Oxidation number Na
Charge
Na
1+
+1
Charge -
1+
Na
Oxidation #s -
1Cl
+1
(H
)2
-2
O
Pairing of oxidation and
reduction
-3
• 2N
 N2 + 6
+6
• 2Cr
+6
e

e
+3
2Cr
Example charge conservation
Black and White
Photography
• Br  Br + e- OXIDATION
• Ag + e-  Ag REDUCTION
Reaction of Silver Halide in Black
and White Photos
Ag+
hv
Br-
Ag+ +
Br +
(s)
e
e
Ag
+ Ag+
Ag +
Ag
Ag
Ag
Ag
Ag Ag
+
Ag
Ag
Ag
Ag
Ag
Use the newly developed negative to
make a print through contact copying
Common
Dry Cell
Batteries
Zn (s)  Zn+2 (s) + 2e- Oxidation
2MnO2 (s) +2NH4+(aq) +2e- 
Mn2O3 (s) +2NH3 (aq) +H2O(l)
Reduction
Alkaline batteries also use
the reduction of
manganese and oxidation
of zinc to create an
electrical current
Dry Cell
Alkaline
NH3Cl reacts with
KOH replaces
Zn even when not NH3Cl and barriers
in use
installed
Cheap
More Expensive
Inconsistent voltage Consistent voltage
and short life
and longer life
Rechargeable Batteries
Reversed poles
extracts electrons
from reduced and
supplies electrons
to oxidized
How do Hybrid
Cars work?
Gelatin Battery Demo p 703
Each cell will become more and more positively charged moving from the right to
left if the flow of electrons does not increase. This will build-up current to the point
of internal resistance and then voltage is built up. When it is disconnected you are
creating a deficit of electrons in the left most cell.
e-
e-
H+
e-
Fuel Cells
eH2 Gas
Catalyst
converts
H2 into H+
H2O
Gas
Battery
and
Engine
2H+
e
Catalyst
Converts O2
into O- and
2H+ + O-  H2O
O2 Gas
Fuel Cell Chemistry
Chemistry of a Fuel Cell
• Anode side:
2H2 => 4H+ + 4e• Cathode side:
O2 + 4H+ + 4e- => 2H2O
• Net reaction:
2H2 + O2 => 2H2O
Character of Fuel Cell
• Low Temperature 176 degrees Celsius
• Each cell is 0.7 volts so many are needed
• Water is pure enough to drink
• Good efficiency (26% compared to 20%)
• Hydrogen is being taken from fossil fuels,
alcohol, and methane
Rules for finding
oxidation numbers
• Periodic Table Position Helps
–1A  +1 (except hydrogen w/ a metal)
–2A  +2
–3A  +3
–6A  -2 (most of the time)
–7A  -1
More Rules
• Uncombined elements always have a
charge of zero
• Charge on any monoatomic ion =
ionic charge
• Oxidation numbers for elements in a
compound are written on a per atom
basis
• The sum of all the individual atom’s
oxidation numbers in a polyatomic ion
equal to the charge on the polyatomic
ion
• The sum of all oxidation numbers in a
compound must equal zero
Practice Finding
Oxidation Numbers
for Elements within
a Compound
You can Mass Balance
a complex redox reaction
by Charge Balancing and
Oxidation Numbers
Charge is conserved
S + HNO3  SO2 + NO + H2O
Oxidation Numbers
Change
Left
Right
Side
Side
+4
S=0
S = +4
N = +5
N = +2
-3
3(+4) = +12
S + HNO3  SO2 + NO + H2O
4(-3) = -12
3(+4) = +12
3S + 4HNO3 SO2 + NO + H2O
4(-3) = -12
Finish Balancing
3S
3S ++ 4HNO
4HNO33
 3SO
3SO22 ++ 4NO
4NO +2H
+__H
2O
2O
Steps used to Balance Equations Using Ox #s
1. Find oxidation numbers
2. Identify what is being reduced and what is being
oxidized
3. Draw arrows overhead reduced and oxidized
substances with the degree of oxidation over top of the
arrow
4. Balance these oxidation numbers with coefficients
5. Use these coefficiencts for a first guess
6. Check the oxidized and reduced species to see if they
are balanced in the entire equation, make adjustments
as necessary
7. Finish balancing the equation by placing coefficients
on reactants and products which were not involved in
the oxidation or reduction.