Transcript Pettinato OXIDATION-REDUCTIONREACTIONS

OXIDATION-REDUCTION REACTIONS
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DEFINITIONS
• Oxidation – loss of electrons, usually done by metals.
• Reduction – gain of electrons, usually done by nonmetals.
• Oxidizing Agent – gains electrons causing oxidation.
• Reducing Agent – loses electrons causing reduction.
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RULES
BEFORE
BALANCING
• Check Oxidation Numbers.
• Oxygen is always -2 except in peroxides where the O- is -1 and as OF2
where O is +2.
• All compounds have a zero net charge.
• In ternary compounds, add the known oxidation numbers to determine
the transition metal’s oxidation number.
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PAY ATTENTION!
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BALANCING RULES
• Assign oxidation numbers to elements undergoing oxidation and
reduction.
• Balance electrons.
• Follow balancing rules from your chemical reactions’ notes.
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DISPROPORTIONATION
• A chemical reaction in which a single compound serves as both
oxidizing and reducing agent and is thereby converted into a more
oxidized and a more reduced derivative.
• Example: The decomposition of hydrogen peroxide.
2H2O2  2H2O + O2
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HALF-REACTION METHOD
• Write down the unbalanced equation.
• Assign oxidation numbers to all of the elements in the equation.
• Decide which atoms change in oxidation number from the
reactant to the product side.
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Half-Reactions Continued
• Write half-reactions for the oxidation (electron loss) and
reduction(electron gain).
This is ONLY done for the elements changing in oxidation
number from the reactant to the product side of the original
equation.
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Half-Reactions Continued
• If an element has a subscript in the unbalanced equation,
carry this into the half-reaction.
– When the element is free, carry the subscript as a subscript.
– When the element is part of a compound or polyatomic ion, carry
the subscript as a coefficient.
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Half-Reactions Continued
• Balance the electron gain and loss in the half-reactions (multiply through the
equations by the appropriate whole numbers).
• Add the oxidation and reduction half-reactions. Since the electron gain and loss
should be equal, the electrons won’t appear in the sum of the equations.
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Half-Reaction Continued
•Use the coefficients in the sum of the half-reactions to determine the coefficients of
the unbalanced equation.
•This procedure will only balance the elements changing in oxidation number.
INSPECTION BALANCING will always be the last step.
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Practice: oxidation numbers in HCO3–
_
H O C O
O
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Practice: oxidation numbers in HCO3–
6 – (4 + 4) = –2
4 – (0 + 0) = +4
6 – (4 + 4) = –2
_
1 – (0 + 0) = +1
H O C O
O
# valence e– in neutral atom
# nonbonding e–
# bonding e– assigned
oxidation number
6 – (6 + 2) = –2
They add up to charge of ion:
(–2) + 4 + (–2) + 1 + (–2) = –1
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Practice: oxidation numbers in K2Cr2O7
(+1)
( ?)
(-2 )
K2Cr2O7
(+2)
+
2(+6)
+
(-14) = 0
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Balancing Redox Equations
• To balance the charge, add electrons
to the side of the equation with the
higher charge.
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Balancing Redox Equations
• Multiply the half-reactions by appropriate
numbers to balance electron gain and loss.
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Balancing Redox Equations
• Add the half-reaction equations, subtracting
any duplications on the left and right sides.
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Half-Reaction in
Acidic Solution
• Again, write the two half-reactions.
• Balance each half-reaction, keeping in mind to
do hydrogen and then oxygen as the last two
steps, respectively. O’s must be balanced with
H2O. H atoms must be balanced with H+ ions.
• Balance the number of electrons in each halfreaction.
• Add the two half-reactions to form a net
balanced ionic equation.
• Check to make sure everything is balanced.
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What if the solution is Basic?
• If the reaction occurs in basic solution, the final equation must
not contain H+.
• If your equation has H+ ions, add an equal number of OH- ions
to both sides of the equation.
• Combine the H+ and OH- into H2O and cancel any replications
of H2O.
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EXAMPLES
1. Al˚ + CuCl2 → AlCl3 + Cu˚
• Each Al loses 3 e•
Each Cu gains 2e- ( 2 X 3 ) = 6
• 2Al˚ + 3CuCl2 → 2AlCl3 + 3Cu˚
2. FeCl3 + H2S → FeCl2 + S˚ + HCl
• Each Fe gains 1 e•
Each S loses 2 e- ( 2 X 1 ) = 2
• 2FeCl3 + H2S → 2FeCl2 + S˚ + HCl
• Balance non-redox species
+ 2HCl
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Gram-Equivalent Weights
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Gram-Equivalent Weight
• Also known as a gram-equivalent or equivalent.
• Used to indicate the amount of a chemical used in a
reaction.
• Can be used instead of moles for redox reactions.
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Gram-Equivalent
• A gram-equivalent weight of a substance is the
amount of this substance which releases or
acquires one mole of electrons during a redox
reaction.
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Gram-Equivalent
• The chief advantage of working with gram-
equivalents is that you don’t have to write a
balanced equation to solve a stoichiometry problem
involving a redox reaction.
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Gram-Equivalent
• A gram-equivalent of oxidizing agent is the
amount of the chemical which accepts one
mole of electrons during a redox reaction.
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Gram-Equivalent
• A gram-equivalent of reducing agent is the
amount of the chemical which releases one
mole of electrons during a redox reaction.
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Gram-Equivalent
• In a redox reaction the total electron gain by
the oxidizing agent must equal the total
electron loss by the reducing agent.
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Gram-Equivalent
• In a redox reaction the number of gram-
equivalents of oxidizing agent must equal the
number of gram-equivalents of reducing agent.
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Gram-Equivalent
• The change in the oxidation number of an element
during a redox reaction is used to determine the
number of gram-equivalents of that substance per
mole.
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NORMALITY OF SOLUTIONS
(NORMALITY AS APPLIED TO REDOX REACTIONS)
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NORMALITY
• The normality of a solution, like its
molarity, indicates the concentration of a
solution.
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NORMALITY

Normality is expressed in terms of the
number of gram-equivalents of solute
per liter of solution.
gram - equivalents of solute
Normality =
liter of solution
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Normality
• The number of gram-equivalents per mole
of a solute may change depending on the
redox reaction that it undergoes.
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Normality
• During a redox titration involving solutions of
oxidizing agent and reducing agent, the gramequivalents of oxidizing agent must equal the gramequivalents of reducing agent.
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Normality
Since,
Volume (in liters) x Normality = G-Equiv.
VOA x NOA = VRA x NRA
Where OA = Oxidizing Agent and
RA = Reducing Agent
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