Redox:
Definitions
Oxidation
:
loss of e- in an atom
increase in oxidation number
(ex: -1
0 or +1
+2)
Reduction
:
gain of e- by an atom
Trick to remember this
:
OIL RIG
O
xidation
i
s
l
oss
Another trick
….
LEO
(the lion says…)
GER
LEO
-
L
ose
E
lectrons
O
xidation
Assigning Oxidation Numbers
(What are they?)
Definition:
The apparent charge assigned to an atom of an element.
Numbers may
be
+, -, or 0
• These are not ionic charges, but maybe the same as the ionic charge.
To distinguish oxidation numbers from charges on ions, the sign of the oxidation number precedes the number.
Example: Ionic Charge = Mg2+
Assigning Oxidation Numbers
(Why do we use them?)
During an Oxidation-Reduction reaction electrons are exchanged or transferred.
Its not always possible to determine whether atoms have exchanged electrons by simply reading an equation.
Chemists have devised a system that makes it easier to keep track of the number of electrons lost or gained during a reaction.
Positive, negative or neutral values known as Oxidation Numbers or States
can be assigned to atoms.
Oxidation numbers help show which atoms and how many electrons are either gained or lost by an atom.
Assigning Oxidation Numbers
(How do we assign them?)
You must!!! Learn the rules for assigning oxidation states to atoms in an equation.
These numbers will be used to identify what is being oxidized and reduced. Rules for assigning Oxidation #’s
1. Each uncombined or lone element has an oxidation number of “0”. Ex. Na + Cl2 >>> NaCl – both Na and Cl2 have oxidation numbers of “0”
2. Monatomic ions have an oxidation number equal to its ionic charge.
Ex. Na + Cl2 >>> NaCl – the Na ion in NaCl will have an oxidation # of +1 the Cl ion in NaCl will have an oxidation # of -1.
3. The metals in group 1 always have oxidation numbers of +1 in compounds, and the metals in group 2 always have an oxidation number of +2 in
compound.
4. Fluorine always has an oxidation number -1 in compound. All the halogens (group 17) will be -1 if they are the most electronegative element in the
Assigning Oxidation Numbers (How
do we assign them? continued)
5. Hydrogen is +1 in compounds unless it is combined with a metal, in which it is -1.
Example: Hydrogen is +1 in HCl but -1 in LiH
6. Oxygen is usually -2 in compounds. Except when its is combined with fluorine, which is more electronegative, its +2.
Example: Oxygen is -2 in H2O, and +2 in OF2. In the peroxide polyatomic ion (O22-)
oxygen is -1.
These 6 rules can be used along with the following 2 additional rules to calculate oxidation numbers for other elements in compounds or polyatomic ions in an
equation.
7. The sum of the oxidation numbers in all compounds must be zero.
Practice Assigning Oxidation
Numbers
Example # 1 HNO3
What are the oxidation numbers of the atoms?
Use the first six rules to assign oxidation numbers first then if you still have some to assign make sure the sum of the oxidation numbers is 0 (rule 7)
Rule 5 states H must be +1
Rule 6 states each oxygen must be -2, and since there are three oxygen atoms they must total (3 x -2) -6
So the total of 1 hydrogen and 3 oxygen atoms is (+1 + -6) -5 which means using rule 7 each nitrogen must be +5 to make the total charge of the compound “0”.
Final answer H = +1 N = +5
More Practice Assigning Oxidation
Numbers
What is the oxidation number of chromium in the dichromate ion? Cr2O7
2-Use as many of the first six rules as possible
O has an oxidation number -2 (rule6), producing a total of (7 x -2) -14 for the seven O atoms.
Using rule 8 the sum of the oxidation numbers must be equal to the charge of the polyatomic ions charge
2 x (the oxidation of Cr) + (-14) = -2 2 x (the oxidation of Cr) must = +12 oxidation of Cr = +12/2
oxidation of Cr = +6
Oxidation
Losing electrons
Any chemical change in which the oxidation number
increases.
Example
2Mg(s) + O2(g) → 2MgO(s)
(Magnesium loses, oxygen gains)
Rusting
Occurs when metals react with oxygen.
(accelerated by water and salts).
Reduction
Gaining electrons.
Any chemical change in which an oxidation
number of an element
decreases
.
Example:
Metal Ores
Occur naturally in
combination or in oxidized state.
Practice
Given a redox reaction… Which is being oxidized?
Which is being reduced?
Magnesium + Oxygen yields Magnesium oxide 2Mg0 + O
20 2Mg+2O-2
Zinc + Copper (II) Sulfate Yields Zinc Sulfate + Copper Zn0 + Cu+2S+6O-2
Redox Reactions
Oxidation cannot occur without reduction.
When electrons from one atom are lost, they must be gained by
another.
We call reactions that involve oxidation and reduction REDOX
reactions.
Not all reactions are redox reactions. To determine whether or
not a reaction is redox, assign oxidation numbers to each atom,
both on the reactant and product side.
If there is a change in oxidation number for any atom, the
Redox Reactions (Continued)
If an element exists alone on one side and in a
compound on the other, then it is a Redox Reaction.
Most single replacement, decomposition and synthesis reactions are Redox
Reducing Agents: Metals
– The
reducing agent
is the species being oxidized.
Oxidizing Agents: Non-metals
The
oxidizing agent
is the species being
reduced.
Oxidation number decreases.
Half-Reactions
A half-reaction shows either the oxidation or reduction portion of a redox reaction, including the electrons
gained or lost.
A reduction half-reaction shows an atom or an ion
gaining one or more electrons while its oxidation number decreases
Ex. Na+1 + 1e- Na0
An oxidation half-reaction shows an atom or an ion
losing one or more electrons while its oxidation number increases.
1e-Half-Reactions (continued)
Half-reactions must show conservation of mass and charge when written properly.
In a half-reaction there will be only one type of atom or ion shown on both reactant and product sides.
In a half-reaction the net charge on both sides of the equation must be equal, but it does not have to be “0”. Ex. Sn+4 + 2e- Sn+2 Net Charge/side = +2
Mass is conserved with 1
Writing Half-Reactions
To write the half-reactions for an equation such as MgCl2 + 2Na 2NaCl + Mg Make sure the equation is balanced
Next assign oxidation numbers to each atom, then write a partial half-reaction to show the change in oxidation number including the coefficients from the balanced equation.
Oxidation: 2Na0 2Na+1
Reduction Mg+2 Mg0
Then place the correct number of electrons on one side of the equation to make the net charge equal on both sides. (remember electrons will always be a product in oxidation and a reactant in reduction)
Oxidation: 2Na0 2Na+1 +
2e-Reduction: 2e- + Mg+2 Mg0
When you have written a correct oxidation and reduction half-reaction the electrons lost (oxidation) should be equal to the electrons gained (reduction)
ELECTROCHEMICAL CELL
The exchange of electrons during a redox reaction can be useful
to us.
One practical use of a redox reaction is in an electrochemical
cell. An electrochemical cell involves a chemical reaction and
the flow of electrons.
There are two common electrochemical cells.
Voltaic cells, which use a spontaneous reaction to produce a flow of electrons or an electric current.
ELECTROCHEMICAL CELL
(continued)
Electrochemical cells have two surfaces called electrodes that
can conduct electricity.
An electrode is were oxidation or reduction will occur.
The electrode at which oxidation occurs is called the anode.
The electrode at which reduction occurs is called the cathode.
Voltaic Cells
Voltaic cells take advantage of electron transfers
during
spontaneous reactions
usually single
replacement.
When you have two metals that differ in
reactivity one will
lose electrons
easily
(oxidation) and the other will
gain the electrons
(reduction) the other has lost.
Voltaic Cells
The metals must be placed in separate ionic solutions that are connected by a salt bridge.
The metals act as electrodes the one that is more
reactive, loses electrons (oxidation), will be the anode, the less reactive metal, gains electrons (reduction), is the cathode.
A voltmeter can be attached to the wire connecting the metals to detect electric current.
Drawing an Voltaic Cell
Solution A Solution
B Salt Bridge
Voltmete r
Voltaic Cells
Each vessel is called a half cell because a half reaction takes place here.
The salt bridge connects the two half cell solutions allowing ions to flow from one solution to the other to balance out the charge.
The electrodes (metals) are connected by a wire which acts as an electrical conductor. The flow of electrons through the wire creates an electric current.
Voltaic Cells & Table J
Oxidation and reduction occur at the electrodes.
To identify which electrode is your anode or the site of oxidation and which is your cathode the sight of
reduction you must compare the reactivity of the metals on Table J in your reference table.
The more reactive metal will always lose electrons (oxidation) which makes it your Anode.
The less reactive metal will always gain electrons (reduction) it is always your cathode.
Table J
So in summary the metal higher up on Table J is your Anode (oxidation) the metal lower on Table J is your cathode (reduction).
Once you identify your anode and cathode you can identify the flow of electrons they flow from the Anode to the Cathode.
Electrolytic Cells
• In a
voltaic cell
electrons flow spontaneously
from anode to cathode.
• The opposite occurs in an
electrolytic cell
,
since the reaction is
non-spontaneous
there
must be an electrical
power source
placed in
the circuit to force the electrons from anode to
cathode.
Electroplating
• The most common use of electrolytic cells is an
electroplating cell.
• Electroplating involves plating a small layer of usually a precious metal on another metal.
• The material to be plated for example a spoon or a
piece of jewelry, would be the cathode were reduction occurs.
Electrolytic Cell For Electroplating
Anode
Cathode Power Source
•Anode-Oxidation •Ag0 Ag+1
Electrolysis
• Electrolysis can be used to break up compounds to reform the elements it is made up of.
• For instance if you wanted to break up water molecules into hydrogen and oxygen gas you could use an
Electrolytic Cell.
• Since the reaction will be non-spontaneous a power source must be used to force the reaction to occur.
• The power source is usually a battery.
Electrolytic Cell for Electrolysis
Power Source
•Write the half-reactions:
•Reduction: 4H+1 + 4e- 2H
2 (Cathode)
•Oxidation: 2O-2 O
Comparing Voltaic and Electrolytic
Comparing Voltaic and Electrolytic
Cells
Cells
Voltaic and Electrolytic cells have several
similarities and differences.
Similarities
1. Both use redox reactions 2. The anode is the site of
oxidation
3. The cathode is the site of reduction
4. The electron flow through the wire is from anode to cathode
Differences
1. Voltaic cells use spontaneous reactions to produce energy (voltmeter)
2. Electrolytic use non-spontaneous reactions that requires energy (power source)
3. Voltaic cells the anode is negative and the cathode is positive
4. Electrolytic cells the anode is
