Transcript Redox Reactions

Objectives
• Assign oxidation numbers to reactant and
product species.
• Define oxidation and reduction.
• Explain what an oxidation-reduction reaction
(redox reaction) is.
19.1 Oxidation & Reduction
• Also called REDOX reactions
• Always involve a TRANSFER of
ELECTRONS
Redox Reactions:
• 2 half-reactions occur at the same time
2Na(s) + Cl2(g) → 2NaCl(s)
LEO – loss of electrons is oxidation
Alkali metals: Na → Na+ + 1e• electron(s) are product
• Sodium is oxidized
GER – gain of electrons is reduction
Halogens: Cl + 1e- → Cl-
• electron(s) are reactant
• Chlorine is reduced
Remember:
“LEO the Lion
says GER!”
Reducing agent
A
Oxidation:
A loses 2
electrons
A
A is oxidized
B is reduced
B
Oxidizing agent
Reduction:
B gains 2
electrons
B
Reducing & Oxidizing Agents
Reducing agent – causes another to be
reduced by donating electrons
• is oxidized (loses electrons)
Oxidizing agent – causes another to be
oxidized by accepting electrons
• is reduced (gains electrons)
Oxidation of Zinc
• Zn loses electrons.
Is zinc oxidized or reduced? Is it the
oxididizing or reducing agent?
• Zinc is oxidized. It is a reducing agent
Half-reaction: Zn → Zn2+ + 2e• Copper gains electrons.
Is copper oxidized or reduced? Is it the
oxididizing or reducing agent?
• Copper is reduced. It is an oxidizing agent.
Half-reaction: Cu2+ + 2e- → Cu
Oxidation States
• Redox reactions can be identified and understood
by assigning oxidation numbers (aka oxidation
states) to reactants and products
Oxidation numbers/states =
• artificial “bookkeeping” devices used to keep track
of overall electron distribution
• not physical characteristics of atoms
• hypothetical charge that an atom would have if all
bonds to atoms of different elements were 100%
ionic.
• Transition metals often have more than one
oxidation state
Ex. Oxidation States of Mn
• In Potassium Permanganate, Mn is in oxidation state +7.
This gives a purple salt.
• In alkaline conditions, and in the presence of a sugar, Mn
is slowly reduced to its +6 and +4 oxidation state. (+6
being green, and +4 being yellow)
Rules for Assigning Oxidation Numbers
1. Uncombined elements = 0
2. Monoatomic ions = charge
3. More electronegative element in binary
compound is assigned number equal to its
charge if it were an ion
4. Fluorine in compound is always -1
5. Oxygen is -2 unless combined with F when its
+1 or +2, or in a peroxide when its -1
6. Hydrogen is usually +1 unless combined with
a metal when it is -1
7. Group 1 & 2 and Aluminum have oxidation
numbers of +1, +2, and +3 respectively
8. The sum of the oxidation numbers of all
atoms in a neutral compound is zero.
9. The sum of the oxidation numbers of all
atoms in a polyatomic ion is equal to the
charge on the ion.
Sample Problem E p. 233
Assign oxidation numbers to each atom in the
following compounds:
UF6
Start with known:
F is always -1
Compounds must add up to zero
U is +6
Sample Problem E p. 233
H2SO4
Oxygen & sulfur more electronegative
Hydrogen assigned +1.
Oxygen not combined with fluorine or in
peroxide.
Oxygen assigned -2.
Sum must be zero.
2(+1) + 4(-2) + 1(x) = 0
x = +6 Sulfur atom assigned +6
ClO-3
Oxygen is more electronegative and is assigned -2
Sum must be equal to charge of polyatomic ion = -1
3(-2) + 1(x) = -1
x = +5
Chlorine is assigned an oxidation number of +5
Homework:
“Assigning Oxidation Numbers” worksheet