Electrochemistry

1
Electrochemistry

• Review of important supporting concepts
• Types of chemical changes
• Oxidation numbers
• Recognizing redox reactions
• Balancing redox reactions
• Electricity and chemistry
• Chemical cells

2
Recognizing Chemical Changes

• Chemical change change in composition
• Gray area changes (textbook inconsistencies)
• Solution events
• Substance dissolving WITHOUT change in species or
  component species is a physical change
• Sodium chloride OR sodium fluoride dissolving
  (neutral salt)
• NaCl (s) ? Na (aq) Cl- (aq) NaF (s) ?
  Na (aq) F- (aq)
• On both sides of the reaction are ions but in
  different locations
• Substance precipitating when two different
  species are combined IS a change in composition
• Calcium carbonate precipitating out of solution
  when sodium carbonate and calcium chloride are
  mixed
• Na2SO4 (aq) CaCl2 (aq) ? CaCO3 (s) 2 NaCl
  (aq) full equation
• Ca2 (aq) CO32- (aq) ? CaCO3 (s)
  net ionic equation
• A new substance different from what we started
  with has been formed
• A molecular compound dissolving in solution and
  producing ions IS a change in composition
• Hydrogen chloride gas dissolving in water
• HCl (g) H2O (l) ? H3O (aq) Cl- (aq)
• Molecular species in reactants ions in products
  
• An ionic compound dissolving in solution and
  producing some molecular species (conjugate acid)
  IS a change in composition
• Fluoride ion (produced above) reacting with water
  in solution
• F- (g) H2O (l) ?? HF (aq) OH- (aq)

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Types of Chemical Changes

• Two types of chemical changes
• Redox
• Oxidation-reduction reaction
• Change in oxidation numbers of at least two
  species in the reaction
• Non redox
• No change in oxidation numbers
• Skill ? knowing when oxidation numbers change

4
Non redox chemical changes

• Acid base reactions
• Weak or strong or any combination
• Neutralizations
• Hydrolysis reactions
• Precipitations (dissolutions)
• Formation of complexes
• Including biological binding
• Hydrations/dehydrations
• Most polymerizations
• Synthetics
• Biological (formation of proteins, nucleic acids,
  carbohydrates, lipids)
• Combination reactions involving like atoms
• Organic chemistry

5
Redox chemical changes

• Every other chemical change!
• How to spot
• Elemental form ? Combined form (or vice versa)
• Zn (s) CuSO4 (aq) ? ZnSO4 (aq) Cu (s)
• MgI2 (aq) Br2 (aq) ? MgBr2 (aq) I2 (aq)
• More H in species ? Less H in species (or vice
  versa)
• XH2 C2H4 ? C2H6 X
• Organic or biochemical reduction (X FAD)
• More O in species ? Less O in species (or vice
  versa)
• HNO3 (aq) C2H6O (l) K2Cr2O7 (aq) ? KNO3 (aq)
  C2H4O (l) H2O (l) Cr(NO3)3 (aq)
  (unbalanced)
• 8 H (aq) Cr2O72- (aq) 3 SO32- (aq) ? 2 Cr3
  (aq) 3 SO42- (aq) 4 H2O (l)
  (unbalanced)
• Change in oxidation numbers

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Oxidation Number Concept

• A set of rules that indicates the arbitrary
  assignment of electrons in compounds to the atoms
  of the elements that comprise the compound
• Based primarily on electronegativity (en)
• The atom(s) of the more electronegative element
  get all the electrons that can be given up by the
  atom(s) of the more electropositive element
• Applies even if en difference if very small!
• Atoms of element in compound getting electrons
  will have negative oxidation number
• Atoms of element in compound giving up electrons
  will have positive oxidation number
• Highly artificial but useful concept
• Oxidation number Charge
• Oxidation numbers apply to all types of matter
  charges apply ONLY to ionic substances

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Rules for Oxidation Numbers

• The oxidation number of atoms in elements is
  always 0 (ZERO)
• If there is only one element present in a formula
  (no matter how many atoms there are in the
  formula), the atoms are in the elemental state
• Oxidation numbers of atoms of elements found in
  combined form have positive or negative
  oxidations numbers
• If there are two or more different kinds of
  elements in a formula, no matter what the formula
  looks like, the atoms are in the combined state
• The maximum positive oxidation number for an
  element is the column number or second digit of
  column number
• Obtained from oxide combinations
• Exceptions O and F
• The maximum negative oxidation number for an
  element is the number of hydrogens in the hydride
  of the element IF THE ELEMENT IS MORE
  ELECTRONEGATIVE THAN H

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Common Oxidation Numbers

• F in combined form will always have an oxidation
  number of -1
• it beats every other element in tug of war for
  electrons
• O in combined form will have an oxidation number
  of -2
• Except with F (obvious)
• Except in certain other combinations (peroxides
  -1)
• H in combined form will have
• an oxidation number of -1 when combined with
  elements with en less than H (metals!)
• an oxidation number of 1 when combined with
  elements with en greater than H (nonmetals!)
• Group 1, 2, 3, 11, 12, 13 (1, 2, 3
  respectively)
• Oxidation number corresponds to charge on
  monatomic ion
• Exception of Cu (1 and 2)
• Other oxidation numbers can usually be determined
  by difference by summing the known oxidation
  numbers in the formula

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Other oxidation numbers for the elements

• Groups 4 and 14
• Group 4 and Si (in group 14) only 4
• C has lots of oxidation numbers determined by
  difference
• Rest of Group 14 also has 2
• Groups 5, 6, 7 (metals)
• Have additional oxidation numbers of 3 and 2
  (as well as higher values in higher rows)
• Correspond to charges on monatomic ions
• Group 15 (nonmetal)
• N has many oxidation s - determined by
  difference
• Other elements in Group have 3 (in addition to
  5)
• Group 16
• All elements (other than O) have 4 (in addition
  to 6)
• Group 17
• All elements (other than F) have 5, 3 and 1
  (in addition to 7)
• Groups 8, 9, 10
• Oxidation numbers of 2 and 3 (as well as higher
  values in higher rows)

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Conservation of Oxidation Numbers

• In any formula, the sum of the oxidation numbers
  has to equal the overall charge on the formula
• For neutral substances this is 0
• 1(4) 1(-4) 0
• 1(4) 4(-1) 0

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-4
Y is more electronegative than X
B is more electronegative than A
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Exercise in Determining Oxidation Numbers
H 1 O -2
?Cl 1

• HClO
• Na2CrO4
• H3PO3
• CoCl3
• Fe(NO2)3
• ClO2-
• SO32-
• BrO3-
• N2O
• NO2-

Na 1(2) O -2(4)
?Cr 6
H 1(3) O -2(3)
?P 3
Cl -1(3)
?Co 3
O -2(6) Fe has to be 3
?N 3 (9/3)
O -2(2) less -1
?Cl 3
O -2(3) less -2
?S 4
O -2(3) less -1
?Br 5
O -2
?N 1
O -2(2) less -1
?N 3
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Components of oxidation-reduction reactions

• Oxidation half-reaction loss of electrons
• Acid half reaction donate (lose) H (proton)
• Reduction half-reaction gain of electrons
• Base half reaction accept (gain) H (proton)

0
0
2
2-
Oxidation half-reaction (lose e-)
Reduction half-reaction (gain e-)
Electrons lost Electrons gained
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Dynamic View of Redox Process
Oxidation half-reaction upper hemisphere of
electron transfer
Balance Electron Transfer
Reduction half-reaction lower hemisphere of
electron transfer
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Additional Important Redox Concepts

• Reducing/Oxidizing Agents
• Reducing agent the species that is oxidized
• Can refer to the entire species OR to the
  elemental component for which the oxidation
  number is changing
• Oxidizing agent the species that is reduced
• Species OR the elemental component
• Complete electron transfer redox
• Formation of ions from elements OR elements from
  ions
• Electron polarization redox
• Movement of electrons toward the more
  electronegative element in a bond

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Complete Electron Transfer
Electron Polarization
C
d-
d-
2 Ca
2 Ca2
4 e-
OCO
4 e-
d
O2
O2
2 O2-

C
O2
CO2

2 Ca
O2
2 CaO
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Writing (and balancing) redox reactions

• Redox reactions usually written in net ionic form
• Removes extraneous nonparticipating species
• Balancing overall redox reaction
• Often have incomplete information
• Use half reactions to balance
• Table of half reactions provided
• All half reactions written as reduction reactions
• Use one reduction half reaction and one reversed
  reduction half reaction (oxidation reaction)
• Half reaction written in order of reduction
  potential
• Reduction potential ? electronegativity
• The more a species wants electrons, the higher
  the potential

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Sense of Reduction Potentials

• F most electronegative wants to be F-
• O wants to be -2
• Ozone peroxides are exceptions
• High oxidation states are not stable
• Element has given up too many electrons
• MnO4- Mn 7 CrO42- Cr 6 HNO3 N 5
• Halogens want to be halides
• Tendency is by electronegativity
• Metals want to lose electrons
• Least electronegative have highest negative
  reduction potentials want to be oxidized!

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Predicting Reduction Potentials

• Estimate Eo for Sr 2 2e- ? Sr given the
  following reduction potentials
• Ba 2 2e- ? Ba Eo - 2.91
• Mg 2 2e- ? Mg Eo - 2.38
• Estimate Eo for the following reaction 2ClO3-
  12 H 10 e- ? Cl2 6 H2O given the table of
  reduction potentials
• Explain Cr 3 3e- ? Cr Eo
  - 0.74
• CrO42- 8 H 3 e- ? Cr 3 4 H2O Eo
  1.20

-2.89
1.47
The metal, Cr, wants to lose electrons but in
chromate it has given up too many e- and is less
stable than as 3
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Balancing Redox Equations Neutral Solution
Balance the net ionic reaction of Al and bromine
to give aluminum bromide.

• Write the unbalanced equation in net ionic form.

• Separate the equation into two half-reactions.

2 e-
Br2
2 Br-
Al
Al3
3e-
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Balancing Redox Equations Neutral Solution
(cont)
3. Balance the atoms and electrons in the half
reactions
4. Combine the half reactions and sum REMOVE
ELECTRONS
6 e-
2
2
6
6
3
6
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Balancing Redox Equations Acid Solution
Balance the oxidation of Fe2 to Fe3 by Cr2O72-
in acid solution that also produces Cr3.

• Write the unbalanced equation for the net ionic
  form.

• Separate the equation into two half-reactions.

?
Cr2O72-
Cr 3
Fe 2
Fe 3
e-
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Balancing Redox Equations Acid Solution (cont)
3. Balance the atoms and electrons in the half
reactions Balance non H or O first
4. Add H to make waters with all available O
atoms
6
12
6 e-
Cr2O72-
Cr 3
2Cr 3
7 H2O
14 H
4. Add electrons to balance charge
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Balancing Redox Equations Acid Solution (cont)
4. Balance electrons, combine the half reactions
and sum REMOVE ELECTRONS
6
6
6
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Balancing Redox Equations - Basic Solution
Write a balanced ionic equation to represent the
oxidation of iodide ion by permanganate ion in
basic solution to yield iodine and manganese(IV)
oxide.

• Write the unbalanced equation for the net ionic
  form.

• Separate the equation into two half-reactions.

?
MnO4-
MnO2
I -
I2
2
2 e-
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Balancing Redox Equations Basic Solution (cont)
3. Balance the atoms and electrons in the half
reactions Balance non H or O first
4. Add H2O to side with extra O to make OH- with
all available O atoms!
3 e-
4 OH-
2 H2O
7. Add electrons to balance charge
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Balancing Redox Equations Basic Solution (cont)
4. Balance electrons, combine the half reactions
and sum REMOVE ELECTRONS
6
3
6
6
2
2
8
4
2 MnO2
6 I -




4 H2O
2 MnO4-
3 I2
8 OH-
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Balancing Redox Equations Practice
Balance each of the following equations
NO3- (aq) Fe2 (aq) ? NO (g) Fe3 (aq) (acid)
MnO4- (aq) IO3- (aq) ? Mn 2 (aq) IO4- (aq)
(base)
Cr 3 (aq) MnO4- (aq) ? Cr2O72- (aq) Mn 2
(aq) (acid)