Transcript Oxidation Numbers

Oxidation Numbers and Redox
Reactions
Section 7.2
Oxidation Numbers
• Oxidation number: the number of
electrons that must be added to or
removed from an atom in a combined
state to convert the atom into the
elemental form
• AKA oxidation state
• Do not have an exact physical meaning
Usage
• Useful for naming compounds, writing
formulas, and balancing chemical
equations
• There are specific rules for assigning
oxidation numbers to a species
• Shared electrons are assumed to belong
to the more electronegative atom in each
bond
Rules
1. Atoms in a pure element have an
oxidation number of 0
2. F has an oxidation number of -1
3. O has an oxidation number of -2, except
in peroxides (-1) or bonded with F (+2)
4. H is +1 when in a compound where the
other element is more electronegative
than it is
Rules Continued
5. H is -1 when it is bonded with metals
6. The more-electronegative element in a
binary molecular compound is
assigned the number equal to the
negative charge it would have as an
anion, the other gets the positive
charge it would get as a cation
Rules Continued
7. The algebraic sum of the oxidation
numbers of all atoms in a neutral
compound is 0
8. The algebraic sum of the oxidation
numbers of all atoms in a polyatomic ion
is equal to the charge of the ion
9. A monoatomic ion has an oxidation
number equal to its charge
Practice
• Assign oxidation numbers to each atom:
• UF6
• We know F has an oxidation number of 1, so the oxidation number for F6 is -6
• Since the compound is neutral, the U
must be +6
• See the board for how to write this
Practice
• Assign oxidation numbers to each atom:
• H2SO4
• H is bonded with more electronegative
atoms, so it is +1 for each (total +2)
• O is -2 for each for a total of -8
• S must be +6 to balance the neutral
formula
Practice
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Assign oxidation numbers to each atom:
ClO3O is -2 for a total of -6
The charge on the ion is -1, so Cl must be
+5
• Practice problems on page 234 in the
book at school.
Redox Reaction
• Reduction and oxidation occurs
• This is a reaction in which electrons are
transferred from one atom to another
• Oxidation: loss of electrons
• Also the increase in oxidation number
• Reduction: gain of electrons
• Also the decrease in oxidation number
• LEO goes GER or OILRIG
Example of Redox Reaction
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•
•
•
Zn(s) + I2(aq) → Zn2+(aq) + 2 I-(aq)
Half reactions are used to show redox
Zn(s) → Zn2+(aq) + 2e- (oxidized)
I2 (aq) + 2e- → 2 I-(aq) (reduced)
Oxidation Number in Redox
Reactions
• 2KBr(aq) + Cl2(aq)  2KCl(aq) + Br2(aq)
• Assign oxidation numbers to all elements
in the balanced equation
• K in KBr = +1, K in KCl = +1 (no change)
• Br in KBr = -1, Br in Br2 = 0 (increase)
• Cl in Cl2 = 0, Cl in KCl = -1 (decrease)
• Bromine is oxidized, chlorine is reduced
Reducing and Oxidizing Agents
• The species that is reduced is called the
oxidizing agent
• The species that is oxidized is called the
reducing agent.
• 2Zn + O2  2ZnO
• Zn in 2Zn = 0, Zn in ZnO = +2, Zn is
oxidized and is the reducing agent
Continued
• O in O2 = 0, O in ZnO = -2, O is reduced
and is the oxidizing agent
• Practice: identified which species is
reduced, oxidized, the oxidizing agent,
and the reducing agent
• N2(g) + 3H2(g)  2NH3(g)
• N is reduced, oxidizing agent, H is
oxidized, reducing agent
Disproportionation
• Disproportionation: a process by which a
substance acts as both the reducing
agent and the oxidizing agent
• 2 Cu+(aq) → Cu2+(aq) + Cu(s)
• +1
+2
0