2. Oxidation Reduction Chemisty: Redox Chemistry
Oxidation and Reduction reactions always take
place simultaneously.
Loss of electrons – oxidation (Increase in
Oxidation Number)
Ex:Na ------> Na+1 + e-1
Gain of electrons - reduction ( Decrease in
Oxidation Number)
Cl2 + 2 e-1 ------> 2 Cl-1
3.
4. Oxidation occurs when a molecule
does any of the following:
Loses electrons Loses hydrogen Gains oxygen
If a molecule undergoes oxidation, it has
been oxidized and it is the reducing agent
(aka reductant).
5. Reduction occurs when a molecule does any of the
following:
Gains electrons
Gains hydrogen
Loses oxygen
If a molecule undergoes reduction,
it has been reduced and it is the
oxidizing agent (aka oxidant).
6.
7. zinc is being oxidized while the copper is
being reduced. Why?
8. Redox reactions involve electron transfer:
Lose e - =Oxidation
Cu (s) + 2 Ag + (aq) Cu 2+ (aq) + 2 Ag(s)
Gain e - =Reduction
9. Oxidation Numbers
• Rules for Assigning Oxidation States
• The oxidation state of an atom in an uncombined element is 0.
• The oxidation state of a monatomic ion is the same as its charge.
• Oxygen is assigned an oxidation state of –2 in most of its covalent
compounds. Important exception: peroxides (compounds
containing the O2 2- group), in which each oxygen is assigned an
oxidation state of –1)
• In its covalent compounds with nonmetals, hydrogen is assigned
an oxidation state of +1
• For a compound, sum total of ON s is zero.
• For an ionic species (like a polyatomic ion), the sum of the
oxidation states must equal the overall charge on that ion.
10. Redox:
Reduction occurs when an atom gains one or more
electrons.
Ex:
Oxidation occurs when an atom or ion loses one or
more electrons.
Ex:
LEO goes GER
Copper metal reacts with silver nitrate to form silver
metal and copper nitrate:
Cu + 2 Ag(NO3) 2 Ag + Cu(NO3)2.
11. Identifying OX, RD, SI
Species
• Ca0 + 2 H+1Cl-1 ® Ca+2Cl-1
2 + H2
0
• Oxidation = loss of electrons. The species becomes more
positive in charge. For example, Ca0 ® Ca+2, so Ca0 is the
species that is oxidized.
• Reduction = gain of electrons. The species becomes more
negative in charge. For example, H+1 ® H0, so the H+1 is
the species that is reduced.
• Spectator Ion = no change in charge. The species does not
gain or lose any electrons. For example, Cl-1 ® Cl-1, so the
Cl-1 is the spectator ion.
12. Oxidizing Agent and Reducing Agent:
Oxidizing agent gets reduced itself and reducing
agent gets oxidized itself, so a strong oxidizing agent
should have a great tendency to accept e and a strong
reducing agent should be willing to lose e easily. What
are strong oxidizing agents- metals or non metals?
Why?
Which is the strongest oxidizing agent and which is the
strongest reducing agent?
13. Agents
• Ca0 + 2 H+1Cl-1 ® Ca+2Cl-1
2 + H2
0
• Since Ca0 is being oxidized and H+1 is being reduced, the
electrons must be going from the Ca0 to the H+1.
• Since Ca0 would not lose electrons (be oxidized) if H+1
weren’t there to gain them, H+1 is the cause, or agent, of
Ca0’s oxidation. H+1 is the oxidizing agent.
• Since H+1 would not gain electrons (be reduced) if Ca0
weren’t there to lose them, Ca0 is the cause, or agent, of
H+1’s reduction. Ca0 is the reducing agent.
14. Steps for Balancing a Redox Reaction:
Half Reaction Method
In half reaction method, oxidation and reduction half-reactions
are written and balanced separately before
combining them into a balanced redox reaction. It is
a good method for balancing redox reactions
because this method can be used both for reactions
carried out in acidic and basic medium .
15. Steps for Balancing Redox Reaction Using
Half Reaction Method IN ACIDIC MEDIUM:
Step 1: Write unbalanced equation in ionic form.
Step 2: Write separate half reactions for the oxidation and
reduction processes. (Use Oxidation Numbers for identifying
oxidation and reduction reactions)
Step 3: Balance atoms in the half reactions
•First, balance all atoms except H and O
•Balance O by adding H2O
•Balance H by adding H+
Step 4: Balance Charges on each half reaction, by adding
electrons.
Step 5: Multiply each half reaction by an appropriate number to
make the number of electrons equal in both half reactions.
Step 6: Add two half reactions and simplify where possible by
canceling species appearing in both sides.
Step 7: Check equation for same number of atoms and charges on
both sides.
16. Writing Half-Reactions
• Ca0 + 2 H+1Cl-1 ® Ca+2Cl-1
2 + H2
0
• Oxidation: Ca0 ® Ca+2 + 2e-
• Reduction: 2H+1 + 2e- ® H0
2
The two electrons lost
by Ca0 are gained by
the two H+1 (each H+1
picks up an electron).
PRACTICE SOME!
17. Practice Half-Reactions
• Don’t forget to determine the charge of each species first!
• 4 Li + O2 ® 2 Li2O
• Oxidation Half-Reaction:
• Reduction Half-Reaction:
• Zn + Na2SO4 ® ZnSO4 + 2 Na
• Oxidation Half-Reaction:
• Reduction Half-Reaction:
18. Steps for Balancing Redox Reaction Using
Half Reaction Method IN BASIC MEDIUM:
For balancing redox reactions in basic solutions, all the
steps are the same as acidic medium balancing,
except you add one more step to it. The H+ ions can
then be “neutralized” by adding an equal number of
OH- ions to both sides of the equation. Ex.
19. Standard Cell Potential
Just as the water tends to flow from a higher level to a lower
level, electrons also move from a higher “potential” to a
lower potential. This potential difference is called the
electromotive force (EMF) of cell and is written as Ecell.
The standard for measuring the cell potentials is called a
SHE (Standard Hydrogen Electrode).
Description of SHE (Standard Hydrogen
Electrode)
Reaction 2H+
(aq, 1M)+ 2e - H2(g, 101kPa) E0= 0.00 V
20. Standard Reduction Potentials
Many different half cells can be paired with the
SHE and the standard reduction potentials for
each half cell is obtained. Check the table for
values of reduction potential for various
substances:
Would substances with high reduction potential
be strong oxidizing agents or strong reducing
agents? Why?
21.
22. Activity Series
• For metals, the higher up the chart the
element is, the more likely it is to be
oxidized. This is because metals like to
lose electrons, and the more active a
metallic element is, the more easily it can
lose them.
• For nonmetals, the higher up the chart the
element is, the more likely it is to be
reduced. This is because nonmetals like to
gain electrons, and the more active a
nonmetallic element is, the more easily it
can gain them.
23. Metal Activity
• Metallic elements start out with a charge
of ZERO, so they can only be oxidized to
form (+) ions.
• The higher of two metals MUST undergo
oxidation in the reaction, or no reaction
will happen.
• The reaction 3 K + FeCl3 ® 3 KCl + Fe
WILL happen, because K is being
oxidized, and that is what Table J says
should happen.
• The reaction Fe + 3 KCl ® FeCl3 + 3 K
will NOT happen.
3 K0 + Fe+3Cl-1
3
REACTION
Fe0 + 3 K+1Cl-1
NO REACTION
24. Voltaic Cells (Galvanic Cells)
A voltaic cell converts chemical energy from a
spontaneous redox reaction into electrical energy.
Ex: Cu and Zn voltaic cell (More positive reduction
potential is the cathode)
Key Words:
•Cathode
•Anode
•Salt Bridge
How a Voltaic Cell Works: An Ox, Red Cat
A reaction is spontaneous if the metal with higher reduction potential is
made cathode.
25. Voltaic Cells
• Produce electrical current using a spontaneous redox
reaction
• Used to make batteries!
• Materials needed: two beakers, piece of the metals (anode,
- electrode and cathode + electrode), solution of each
metal, porous material (salt bridge), solution of a salt that
does not contain either metal in the reaction, wire and a
load to make use of the generated current!
• Use Reference Table J to determine the metals to use
– Higher = (-) anode (lower reduction potential)
– Lower = (+) cathode (higher reduction potential)
27. Electrolytic Cells
• Use electricity to force a nonspontaneous redox reaction to
take place.
• Uses for Electrolytic Cells:
– Decomposition of Alkali Metal Compounds
– Decomposition of Water into Hydrogen and Oxygen
– Electroplating
• Differences between Voltaic and Electrolytic Cells:
– ANODE: Voltaic (-) Electrolytic (+)
– CATHODE: Voltaic (+) Electrolytic (-)
– Voltaic: 2 half-cells, a salt bridge and a load
– Electrolytic: 1 cell, no salt bridge, IS the load
28. Decomposing Alkali
Metal Compounds
2 NaCl ® 2 Na + Cl2
The Na+1 is reduced at
the (-) cathode,
picking up an e- from
the battery
The Cl-1 is oxidized at
the (+) anode, the e-being
pulled off by the
battery (DC)
29. Decomposing Water
2 H2O ® 2 H2 + O2
The H+ is reduced at
the (-) cathode,
yielding H2 (g), which
is trapped in the tube.
The O-2 is oxidized at
the (+) anode, yielding
O2 (g), which is
trapped in the tube.
30. Electroplating
The Ag0 is oxidized to Ag+1
when the (+) end of the
battery strips its electrons
off.
The Ag+1 migrates through
the solution towards the (-)
charged cathode (ring),
where it picks up an electron
from the battery and forms
Ag0, which coats on to the
ring.
