Transcript Redox electrochemistry PPT

Redox- Electrochemistry
“Oxidation-Reduction Reactions”
LEO SAYS GER
The Meaning of Oxidation and
Reduction (called “redox”)
OBJECTIVES
Define oxidation and
reduction in terms of the loss
or gain of oxygen, and the
loss or gain of electrons.
The Meaning of Oxidation and
Reduction (Redox)
OBJECTIVES
State the characteristics of
a redox reaction and identify
the oxidizing agent and
reducing agent.
The Meaning of Oxidation and
Reduction (Redox)
OBJECTIVES
Describe what happens to
iron when it corrodes.
Oxidation and Reduction (Redox)
Early chemists saw “oxidation”
reactions only as the combination of
a material with oxygen to produce
an oxide.
• For example, when methane
burns in air, it oxidizes and forms
oxides of carbon and hydrogen.
Oxidation and Reduction (Redox)
But, not all oxidation processes
that use oxygen involve burning:
•Elemental iron slowly oxidizes to
compounds such as iron (III)
oxide, commonly called “rust”
•Bleaching stains in fabrics
•Hydrogen peroxide also releases
oxygen when it decomposes
Oxidation and Reduction (Redox)
A process called “reduction” is the
opposite of oxidation, and originally
meant the loss of oxygen from a
compound
Oxidation and reduction always occur
simultaneously
The substance gaining oxygen (or
losing electrons) is oxidized, while the
substance losing oxygen (or gaining
electrons) is reduced.
Oxidation and Reduction (Redox)
Today, many of these reactions may
not even involve oxygen
Redox currently says that electrons
are transferred between reactants
Mg
+
S→
Mg2+
+
S2-
(MgS)
•The magnesium atom (which has zero charge) changes to a
magnesium ion by losing 2 electrons, and is oxidized to Mg2+
•The sulfur atom (which has no charge) is changed to a
sulfide ion by gaining 2 electrons, and is reduced to S2-
Oxidation and Reduction (Redox)
0
1
0
1
2 Na  Cl 2  2 Na Cl
Each sodium atom loses one electron:
1
0
Na  Na  e

Each chlorine atom gains one electron:
0

1
Cl  e  Cl
LEO says GER :
Lose Electrons = Oxidation
1
0
Na  Na  e

Sodium is oxidized
Gain Electrons = Reduction
0

1
Cl  e  Cl
Chlorine is reduced
LEO says GER :
- Losing electrons is oxidation, and the
substance that loses the electrons is
called the reducing agent.
- Gaining electrons is reduction, and
the substance that gains the electrons is
called the oxidizing agent.
Mg is the
reducing
agent
Mg is oxidized: loses e-, becomes a Mg2+ ion
Mg(s) + S(s) → MgS(s)
S is the oxidizing agent
S is reduced: gains e- = S2- ion
Oxidation and Reduction (Redox)
It is easy to see the loss and gain of
electrons in ionic compounds, but what
about covalent compounds?
In water, we learned that oxygen is
highly electronegative, so:
the oxygen gains electrons (is
reduced and is the oxidizing agent),
and the hydrogen loses electrons (is
oxidized and is the reducing agent)
Not All Reactions are Redox Reactions
- Reactions in which there has been no
change in oxidation number are NOT
redox reactions.
Examples:
1 5 2
1
1
1
1
1 5 2
Ag N O 3 ( aq )  Na Cl ( aq )  Ag Cl ( s )  Na N O 3 ( aq )
1  2  1
1
6 2
1
6 2
1
2
2 Na O H ( aq )  H 2 S O 4 ( aq )   Na 2 S O 4 ( aq )  H 2 O (l )
Corrosion
•Damage done to metal is costly to
prevent and repair
•Iron, a common construction metal often
used in forming steel alloys, corrodes by
being oxidized to ions of iron by oxygen.
•This corrosion is even faster in the
presence of salts and acids, because
these materials make electrically
conductive solutions that make
electron transfer easy
Corrosion
•Luckily, not all metals corrode easily
•Gold and platinum are called noble
metals because they are resistant to
losing their electrons by corrosion
•Other metals may lose their electrons
easily, but are protected from corrosion by
the oxide coating on their surface, such as
aluminum
•Iron has an oxide coating, but it is not
tightly packed, so water and air can
penetrate it easily
Corrosion
•Serious problems can result if bridges,
storage tanks, or hulls of ships corrode
•Can be prevented by a coating of oil,
paint, plastic, or another metal
•If this surface is scratched or worn away,
the protection is lost
•Other methods of prevention involve the
“sacrifice” of one metal to save the second
•Magnesium, chromium, or even zinc
(called galvanized) coatings can be applied
Oxidation Numbers
OBJECTIVES
Determine the oxidation
number of an atom of any
element in a pure substance.
Oxidation Numbers
OBJECTIVES
Define oxidation and
reduction in terms of a
change in oxidation number,
and identify atoms being
oxidized or reduced in redox
reactions.
Assigning Oxidation Numbers
• An “oxidation number” is a positive or
negative number assigned to an atom
to indicate its degree of oxidation or
reduction.
• Generally, a bonded atom’s oxidation
number is the charge it would have if
the electrons in the bond were
assigned to the atom of the more
electronegative element
Rules for Assigning Oxidation Numbers
1) The oxidation number of any
uncombined element is zero.
2) The oxidation number of a
monatomic ion equals its charge.
0
0
1
1
2 Na  Cl 2  2 Na Cl
Rules for Assigning Oxidation Numbers
3) The oxidation number of oxygen in
compounds is -2, except in
peroxides, such as H2O2 where it is -1.
4) The oxidation number of hydrogen in
compounds is +1, except in metal
hydrides, like NaH, where it is -1.
1
2
H2O
Rules for Assigning Oxidation Numbers
5) The sum of the oxidation numbers of the
atoms in the compound must equal 0.
1
2
H2O
2(+1) + (-2) = 0
H
O
2
2  1
Ca(O H ) 2
(+2) + 2(-2) + 2(+1) = 0
Ca
O
H
Rules for Assigning Oxidation Numbers
6) The sum of the oxidation numbers in
the formula of a polyatomic ion is equal
to its ionic charge.
? 2
N O3

X + 3(-2) = -1
N
O
thus X = +5
? 2
S O4
2
X + 4(-2) = -2
S
O
thus X = +6
Reducing Agents and Oxidizing Agents
• An increase in oxidation number = oxidation
• A decrease in oxidation number = reduction
1
0
Na  Na  e

Sodium is oxidized – it is the reducing agent
0

1
Cl  e  Cl
Chlorine is reduced – it is the oxidizing agent
Trends in Oxidation and Reduction
Active metals:
Lose electrons easily
Are easily oxidized
Are strong reducing agents
Active nonmetals:
Gain electrons easily
Are easily reduced
Are strong oxidizing agents
Balancing Redox Equations
OBJECTIVES
Describe how oxidation
numbers are used to identify
redox reactions.
Balancing Redox Equations
OBJECTIVES
Balance a redox equation
using the oxidation-numberchange method.
Balancing Redox Equations
OBJECTIVES
Balance a redox equation
by breaking the equation into
oxidation and reduction halfreactions, and then using the
half-reaction method.
Identifying Redox Equations
In general, all chemical reactions can
be assigned to one of two classes:
1) oxidation-reduction, in which
electrons are transferred:
• Single-replacement, combination,
decomposition, and combustion
2) this second class has no electron
transfer, and includes all others:
• Double-replacement and acidbase reactions
Identifying Redox Equations
In an electrical storm, nitrogen and
oxygen react to form nitrogen monoxide:
N2(g) + O2(g) → 2NO(g)
•Is this a redox reaction?
YES!
•If the oxidation number of an element
in a reacting species changes, then
that element has undergone either
oxidation or reduction; therefore, the
reaction as a whole must be a redox.
•Conceptual Problem 20.4, page 647
Balancing Redox Equations
It is essential to write a correctly
balanced equation that represents
what happens in a chemical reaction
• Fortunately, two systematic methods
are available, and are based on the
fact that the total electrons gained in
reduction equals the total lost in
oxidation. The two methods:
1) Use oxidation-number changes
2) Use half-reactions
Using Oxidation-Number Changes
Sort of like chemical bookkeeping, you
compare the increases and decreases in
oxidation numbers.
•start with the skeleton equation
•Step 1: assign oxidation numbers to all
atoms; write above their symbols
•Step 2: identify which are oxidized/reduced
•Step 3: use bracket lines to connect them
•Step 4: use coefficients to equalize
•Step 5: make sure they are balanced for
both atoms and charge – Problem 20.5, 649
Using half-reactions
A half-reaction is an equation showing
just the oxidation or just the reduction that
takes place
they are then balanced separately, and
finally combined
Step 1: write unbalanced equation in ionic
form
Step 2: write separate half-reaction
equations for oxidation and reduction
Step 3: balance the atoms in the halfreactions (More steps on the next screen.)
Using half-reactions
continued
•Step 4: add enough electrons to one side
of each half-reaction to balance the charges
•Step 5: multiply each half-reaction by a
number to make the electrons equal in both
•Step 6: add the balanced half-reactions to
show an overall equation
•Step 7: add the spectator ions and balance
the equation
Choosing a Balancing Method
1) The oxidation number change
method works well if the oxidized
and reduced species appear only
once on each side of the equation,
and there are no acids or bases.
2) The half-reaction method works
best for reactions taking place in
acidic or alkaline solution.
Electrochemistry
Applications of Redox
Review
• Oxidation reduction reactions involve a
transfer of electrons.
• LEO-GER
• Lose Electrons Oxidation
• Gain Electrons Reduction
Applications
• Batteries
– Moving electrons is electric current.
– 8H++MnO4-+ 5Fe+2 +5e Mn+2 + 5Fe+3 +4H2O
• Helps to break the reactions into half
reactions.
• 8H++MnO4-+5e-  Mn+2 +4H2O
• 5(Fe+2  Fe+3 + e- )
• In the same mixture it happens
without doing useful work, but if
separate
• Connected this way the reaction
starts
• Stops immediately because charge
builds up.
H+
MnO4-
Fe+2
Electrochemical Cell
H+
MnO4-
Fe+2
Salt
Bridge
allows
current
to flow
because
it allows
for
migration
of ions.
• Electricity travels in a complete
circuit
• Instead of a salt bridge
H+
MnO4-
Fe+2
External
wire
allows the
movement
of
electrons
Porous
Disk
H+
MnO4-
Fe+2
e-
e-
e-
e-
Anode
e-
Reducing
Agent
Cathode
e-
Oxidizing
Agent
Cell Potential
• Oxidizing agent pushes the electron.
• Reducing agent pulls the electron.
• The push or pull (“driving force”) is
called the cell potential Ecell
• Unit is the volt(V) = 1 joule of work/coulomb of
charge
• Measured with a voltmeter
Electrochemical Cell
•
•
1)
2)
3)
4)
The reaction always runs
spontaneously in the direction that
produced a positive cell potential.
Four things for a complete
description.
Cell Potential
Direction of flow
Designation of anode and cathode
Nature of all the components-
Practice
• Completely describe the galvanic cell
based on the following half-reactions
under standard conditions.
• MnO4- + 8 H+ +5e-  Mn+2 + 4H2O
Eº=1.51
• Fe+3 +3e-  Fe(s)
Eº=0.036V
Old Regent’s
reference table N.
Batteries are Electrochemical Cells
• Car batteries are lead storage batteries.
• Pb +PbO2 +H2SO4 PbSO4(s) +H2O
• Dry Cell
Zn + NH4+ +MnO2  Zn+2 + NH3 + H2O
• Alkaline
Zn +MnO2  ZnO+ Mn2O3 (in base)
• NiCad
• NiO2 + Cd + 2H2O  Cd(OH)2 +Ni(OH)2
Corrosion
• Rusting - spontaneous oxidation.
• Most structural metals have reduction
potentials that are less positive than
O2 .
• Fe  Fe+2 +2eEº=
0.44 V
• O2 + 2H2O + 4e- 4OHEº=
0.40 V
• Fe+2 + O2 + H2O Fe2 O3 + H+
Salt speeds up process by increasing
conductivity
Wate
r
Rust
e-
Iron Dissolves- Fe  Fe+2
Preventing Corrosion
• Coating to keep out air and water.
• Galvanizing - Putting on a zinc coat
• Has a lower reduction potential, so it
is more. easily oxidized.
• Alloying with metals that form oxide
coats.
• Cathodic Protection - Attaching large
pieces of an active metal like
magnesium that get oxidized instead.
Electrolysis
• Running a electrochemical cell
backwards.
• Put a voltage bigger than the
potential and reverse the direction of
the redox reaction.
• Used for electroplating.
1.1
0
e-
Zn
1.0 M
Zn+2
Anode
e-
1.0 M
Cu+2
Cathode
Cu
e-
Zn
e-
A
battery
>1.10V
1.0 M
Zn+2
Cathode
1.0 M
Cu+2
Anode
Cu
Other uses
• Electroysis of water.
• Seperating mixtures of ions.
• More positive reduction potential
means the reaction proceeds forward.
• We want the reverse.
• Most negative reduction potential is
easiest to plate out of solution.