Transcript Redox reactions

REDOX REACTIONS AND
ELECTROCHEMISTRY
Oxidation and reduction are
important chemical reactions.
They work both for our benefit, in
forms such as batteries, and
against us, in ways such as the
corrosion of important metals.
OXIDATION AND REDUCTION
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In many chemical reactions, electrons are
removed from some particles (atoms or ions)
and are transferred to other particles. Such
reactions are called oxidation-reduction or
redox reactions.
REDOX: one atom loses an electron, and
another atom gains the electron
Oxidation Number
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Oxidation numbers, or oxidation state, is convenient
for keeping track of the number , as well as the
transfer of electrons in a chemical reaction.
The oxidation number of an atom is a positive,
negative or zero charge assigned to the atom in a
chemical formula in accordance with certain arbitrary
rules.
Rules
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1. Uncombined elements
2. Group 1 metals (1+ions)
Group 2 metals (2+ions)
4. Flourine
5. Hydrogen
except in metal hydrides (LiH)
(group 1,2 metal)
6. Oxygen
except in peroxides (H2O2)
except with flourine (OF2)
7. Monoatomic ions
0
+1
+2
-1
+1
-1
-2
-1
+2
Sample Problem
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What are the oxidation numbers of the atoms in H2SO4?
Solution: Identify the known and unknown values.
Using the six rules assign oxidation numbers to each element.
Rule 5: hydrogen (H) +1,
Rule 6: oxygen (O) -2,
Sulfur = ? (multiple oxidation numbers)
H2SO4 :
2(+1) + S +4(-2) = 0
S + (-6) =O
S=6
Practice problems
Sample problem: Oxidation Numbers in
Ions
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What is the oxidation number of chromium in the dichromate
ion (Cr2O7)2Solution: Identify the known and unknown values.
Using as many of the rules as possible.
O has an oxidation number of -2,
The sum of the atoms in a polyatomic ion must equal the
charge on the ion.
(Cr2O7)22Cr + 7(-2) = -2
2Cr = 12
Cr = 6
Redox reactions
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Redox reactions are a family of reactions that
are concerned with the transfer of electrons
between species. Redox reactions are a
matched set -- you don't have an oxidation
reaction without a reduction reaction
happening at the same time.
Each reaction, oxidation or reduction, is
called a half-reaction.
Oxidation:
LEO: lose electrons, oxidized
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Loss of electrons by an atom or ion.
Loss of electron results in a gain/increase in
oxidation number
Increase in oxidation number, ex. -2 to 0, 1
to 2 (becomes more positive)
Reduction:
GER: gain electrons, reduced
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Gain of electrons by an atom or ion.
Gain of electron results in a decrease in
oxidation number.
Decrease in oxidation number, ex. 0 to -1, 3
to 1 (becomes less positive or more
negative)
Recognizing Redox Reactions
MnO2 + 4 HCl
 MnCl2
+
Mn(+4), O(-2) H(+1),Cl(-1)  Mn(+2), Cl(-1)
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Cl2 +
Cl(0)
2H2O
H(+1), O(-2)
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 an
atom/ion, the reaction is a redox reaction.
Increase in oxidation number – atom/ion is oxidized
Decrease in oxidation number – atom/ion is reduced
Double Replacement reactions are NOT redox
reactions.
MnO2 + 4 HCl
 MnCl2
+ Cl2 +
2H2O
Mn(+4), O(-2) H(+1),Cl(-1) Mn(+2), Cl(-1) Cl(0)
H(+1), O(-2)
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Mn changes from a +4 to a +2: decrease in
oxidation number
***Mn+4 is reduced
Cl changed from a -1 to 0: increase in oxidation
number
***Cl-1 is oxidized
*** ATOM/IONS OXIDIZED OR REDUCED ARE
ALWAYS REACTANTS
Redox Reactions
Oxidizing Agents and Reducing Agents
 The substance oxidized is the reducing
agent
 The substance reduced is the oxidizing
agent
Half-Reactions
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A half-reaction shows either the oxidation or
reduction portion of a redox reaction,
including the electrons gained or lost.
 Reduction half-reaction
Shows an atom or ion gaining one or more
electrons. (electrons are reactants)
Shows the decrease in oxidation number
Fe 3+ (aq) + 3e-  Feo(s)
Charge = 0
+3 + -3 = 0
Charge = 0
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Shows the decrease in oxidation number
Fe 3+ (aq) + 3e-  Feo(s)
Oxidation half-reaction
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Shows an atom or ion losing one or more
electrons. (electrons are products)
Shows an increase in oxidation number
Feo(s)  Fe 3+ (aq) + 3e-
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Charge = 0
Charge = 0
+3 + -3 = 0
Half-reactions
***Half reactions follow the law of conservation of
matter. There must be equal number of atoms on
each side of the equation. (equal atoms between
reactants and products)
 ***Half-reactions follow the law of conservation of
charge. The net charge must be the same on both
sides of the equation.
The number of electrons lost must be equal to the
number of electrons gained.
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Writing half-reactions
Ex. Cu + 2AgNO3  Cu(NO3)2 + 2Ag
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1. assign oxidation number to each element
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2. write a partial half-reaction to show the change in oxidation state
Oxidation: Cu  Cu 2+
Reduction: Ag+  Ag
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3. show the number of electrons needed to explain the change in oxidation number
Oxidation: Cu  Cu 2+ + 2enet charge/side = 0
+
Reduction: 2Ag + 2e-  2Ag
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net charge/side = 0
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4. balance number of electrons lost and gained
Oxidation: Cu  Cu 2+ + 2eReduction: 2Ag+ + 2e-  2Ag
Balanced electrons: 2 lost by Cu, 2 gained by Ag
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Cu + 2AgNO3  Cu(NO3)2 + 2Ag
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Predicting Spontaneous reactions
(Activity SeriesTable)
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A redox reaction is spontaneous when the
pure element is higher up on the activity
series.
A redox reaction is nonspontaneous when
the pure element is lower down on the
activity series.
Hydrogen is the arbitrary standard used to
rank the activity of the elements
Table J
Electrode Potential
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Electrode potentials are relative values based on
assigning the voltage of half-reaction
2 H+ + 2 e-  H2 a value of zero.
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If a half-reaction is reversed from what it is on the
table the sign of the voltage is reversed.
If a redox reaction is at equilibrium, then the
voltage is zero!
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Electrode Potential
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Positive electrode potential – reaction is
SPONTANEOUS.
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Negative electrode potential – reaction is
NONSPONTANEOUS.
Balancing half reactions
A redox reaction can be balanced by separating it
into its two half reactions and then by balancing the
half reaction against each other by balance the
electrons!
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Electrochemical Cell
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Two types
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Voltaic cell: spontaneous chemical reaction
Electrolytic cell: non spontaneous reaction that
requires an electric current
Voltaic Cell-Spontaneous reaction
Chemical energy converted to electrical energy
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Two electrodes: conduct
electricity
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anode- site of oxidation
(AN OX)
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Half cell
negative electrode
cathode – site of reduction
(RED CAT)
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Half cell
Electrodes – site at which
oxidation and reduction occurs.
positive electrode
Electrons move through the
wire from the anode to the
cathode.
Site of
reduction
↑Mass of Cu(s) electrode
-----
↓[Cu2+] in solution
Site of
oxidation
↑[Zn2+] in solution
↓mass of Zn(s) electrode
Why? Cu+2 ions in solution
are gaining electrons,
forming Cu(s). The Cu(s)
is attaching to the electrode
Why? Zn(s) is changing to
ions and going into solution.
+ ions flow into the
half cell from the salt
bridge to replace +
ions lost through
reduction.
- ions flow into the half
cell from the salt
bridge to neutralize
the + ions added by
the oxidation of the
zinc.
Voltaic Cell
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Electrons travel through the
wire from the anode to the
cathode.
Salt bridge connects the two
cells and provides a path for
the movement of ions
between the two cells.
(maintains an ionic balance)
Chemical energy is
converted to electrical
energy.
Table J
Can be used to determine the anode and
cathode.
 Can be used to determine if a reaction will be
spontaneous.
Metal higher on the table will be oxidized and is
the anode.
Ion of the metal lower on the table will be
reduced and is the cathode
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Table J
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Fe/Fe2+ // Pb/Pb2+
Higher Metal is oxidized
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O: Fe Fe2+ + 2e-1
Lower Metal ion is reduced
–
R: Pb2+ + 2e-1  Pb
Electrolytic cells: Non spontaneous Reactions
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2 major differences:
1. Anode is positive, Cathode is negative
2. Electrical energy is converted into chemical energy.
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1. Electrolysis
2. Electroplating
Electrolysis: when electricity is used to force
a chemical reaction.
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Practical uses:
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1. Used to obtain active
elements by the electrolysis of
their molten salts.
2 NaCl (l)→ 2Na(s)+ Cl2(g)
(anode): 2 Cl-  Cl2 + 2 e(cathode): Na+ + e– → Na(s)
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The salt has been heated until it melts,
the Na+ ions flow toward the negative
electrode and the Cl- ions flow toward
the positive electrode.
Electrolysis of NaCl(aq)-brine
2 NaCl(aq) + 2 H2O→2 NaOH +Cl2(g) + H2(g)
(anode): 2 Cl-  Cl2 + 2 e(cathode): Na+ + e– → Na(s)
H2O + 2 e– → H2(g) + 2 OH–
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Electrolysis of water
Practical use
2H2O(l)+electricity2H2 + O2
4H2O + 4e-  2H2 + 4OH-
Reduction: (cathode )
H: (+1  0)
Oxidation:
O: (-2  0)
p. 680 text book
Electroplating
Applying a metal onto a surface
 Object to be plated (spoon)
is the cathode
R: Ag+ (aq) + 1e-  Ago (s)
 Ag electrode is the anode
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O: Ag
(s)
anode
Ag+1 (aq) + 1e-
**anode is positive
**cathode is negative
cathode
Electrochemical Cells
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Similarities between voltaic
cells and electrolytic cells
1. Both use redox reactions.
2. The anode is the site of
oxidation.
3. The cathode is the site of
reduction.
4. The electrons flow
through the wire from anode
to cathode.
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Differences between voltaic
cells and electrolytic cells
1. The redox reaction in a
voltaic cell is spontaneous,
but it is nonspontaneous in
an electrolytic cell
2. In a voltaic cell the anode
is negative and the cathode
is positive. In an electrolytic
cell, the anode is positive
and the cathode is negative.
3. Voltaic cell converts
chemical energy to electrical
energy. Electrolytic cell
converts electrical energy to
chemical energy.
Dry Cell vs Alkaline Battery