Transcript 11/1 Lecture

Chem. 1B – 11/1 Lecture
Announcements I
• Exam #2 - Results
– Average = 59.4
– Worst average so far for
any Chem 1B exam here
– Fraction of students
better than 90 was
reasonable, but many,
many students under 50
– Many questions like quiz
or last year’s exam
questions
Score Range
# Students
90-103
7
80s
14
70s
14
60s
27
50s
31
<50
40
Solutions for B version posted –
will post equivalent C version
question and % correct
Announcements II
• Post Exam 2 Grades
– Very few high scores
even if average is
better than exam 2
– Score is with 380 to
440 points (~40%)
– Cut-offs may change
slightly, but too early
to define
% Range
# Students
90-103
1
80s
14
70s
32
60s
50
50s
22
<50
16
Students in 60s have a chance to
improve (with more effort)
Announcements III
• Today’s Lecture
– Electrochemistry (Ch. 18 – Exam 3 material)
• Redox Reactions – various formats
• Voltaic (or Galvanic) Cells
• Definitions
• Standard Half-Cells and Cells
• Standard Reduction Potential
• Standard Cell Potentials
Chapter 18 Electrochemistry
• Electrochemical Reactions – Different Forms
– “Beaker” Reactions
•
•
•
•
Products form along with heat (assuming DH < 0)
Little control of reaction
Products co-mingled (from reduction and oxidation)
Example: nail “rusts” (oxidation of Fe, reduction of O2)
– Voltaic (Galvanic) Cells
• Oxidation and reduction reactions may be divided into
different parts (half-cells sometimes physically separated
through two reaction cells)
• Two electrodes are also needed
• Reaction can be “harnessed” through voltage/power
production
• Examples: batteries, pH measuring electrodes
Chapter 18 Electrochemistry
• Electrochemical Reactions – Different Forms
– Electrolytic Cell
• In this type of cell, external electrical energy is used to force
unfavorable reactions (e.g. 2H2O(l) ↔ 2H2(g) + O2(g)) to
occur
• Also requires two electrodes – but some differences from
electrodes of voltaic cells
• Examples: Production of Cl2 gas from NaCl(aq), production of
H2 gas from water (above), instruments that measure degree
of oxidation/reduction at specific voltages (analogous to
spectrometers)
Chapter 18 Electrochemistry
• Voltaic Cells - Description
of how example cell works
VOLTAIC CELL
voltmeter
– Reaction on anode =
Ag+ + e- → Ag(s)
oxidation
Zn(s)
– Anode = Zn electrode (as Zn
Ag(s)
has a greater tendency to
+
–
oxidize than Ag)
– So, reaction on cathode must
be reduction and involve Ag
– Oxidation produces e-, so
anode has (–) charge (voltaic
cells only); current runs from
cathode to anode
AgNO3(aq)
– Salt bridge allows
ZnSO4(aq)
replenishment of ions as
cations migrate to cathode Zn(s) → Zn2+ + 2e- Salt Bridge
and anions toward anodes
Chapter 18 Electrochemistry
• Basic Electrical Quantities
– Current: the flow of electrons (although defined
where a positive current has electrons moving
backwards)
– Current units: Amperes (A) with 1 A = 1 C/s and 1 C
= 1 Coulomb where 1 electron (elementary charge)
has a value of 1.60 x 10-19 C
– Potential or Voltage: The potential energy associated
with the movement of charge (e.g. to electrode of
opposite sign)
– Potential units: Volts (V) = 1 J/C
Chapter 18 Electrochemistry
• Basic Electrical Quantities – From Voltaic Cells
– Current: related to the flow of electrons
– Potential: related to the reaction occurring (more
energetic means higher potential)
– The ability of a metal (or other elements) to reduce
can be measured under standard conditions
– Example: Zn(s) + 2Ag+(aq) ↔ Zn2+(aq) + 2Ag(s)
If [Ag+] and [Zn2+] = 1 M, Ecellº = 1.56 V
Chapter 18 Electrochemistry
Voltaic Cells
• Cell notation
Voltaic CELL
– Example Cell:
voltmeter
Zn(s)|Zn2+(aq)||Ag+ (aq)|Ag(s)
Zn(s)
Ag(s)
“|” means phase
boundary
left side for
anode (right
side for
“||” means salt bridge
cathode)
AgNO3(aq)
ZnSO4(aq)
Salt Bridge
Chapter 18 Electrochemistry
• Example Questions
– Given the following cell, answer the following
question:
MnO2(s)|Mn2+(aq)||Cr3+(aq)|Cr(s)
– What compound is used for the anode?
– What compound is used for the cathode?
– Write out both half-cell reactions and a net reaction
Chapter 18 Electrochemistry
• Given the following
cell, write the cell
notation:
GALVANIC CELL
voltmeter – reads
+0.43 V
Pt(s)
+
Note: In this case the
Pt(s) is an “inert”
electrode (provides
electrons but doesn’t
react
–
Ag(s)
AgCl(s)
FeSO4 (aq),
Fe2(SO4)3(aq)
NaCl(aq)
Salt Bridge
Chapter 18 Electrochemistry
Standard Reduction Potential
• A cell used to determine the
standard reduction potential
consists of two half cells
• One half-cell, the anode, is the
standard hydrogen electrode
Pt(s)
+
(2H (aq) + 2e ↔ H2(g))
• Eanodeº = 0 (defined)
• Other is the test cell (compound
being reduced when half-cell is
coupled to standard hydrogen
electrode (oxidation electrode)
• Both cells under standard
H2(g)
conditions (1 M, 1 atm)
H+(aq)
• Ecellº = Ecathodeº
• The SHE is not actually used much
any more (just a reference for
relative potential)
Ag(s)
AgNO3(aq)
Chapter 18 Electrochemistry
Standard Reduction Potential
• Meaning of Values
– Half-cells that exhibit positive values
have electrodes with compounds that
easily reduce (e.g. Ag+(aq), MnO4-,
PbO2(s))
– Half-cells that exhibit negative values
have electrodes that easily oxidize
(e.g. alkali metals)
– What if we have two half-cells
(neither SHE), can we find Ecellº?
Example: Zn(s)|Zn2+(aq)||Ag+
(aq)|Ag(s) Ecellº = ?
Ag+ reduction
Eº = +0.80 V
Eº = 0
Zn2+
reduction
Eº = -0.76 V
Chapter 18 Electrochemistry
• Example Question
– An Ag/AgCl electrode is a common reference
electrode. What is the standard potential of a cell
made up of a Cu2+ solution being reduced to Cu(s)
and AgCl(s) being reduced to Ag(s)?
E°(Cu2+ + 2e- ↔ Cu(s)) = 0.34 V
E°(AgCl(s) + e- ↔ Ag(s) + Cl- (aq)) = 0.22 V
What is the balanced reaction and what species must
be present at 1 M?
Chapter 18 Electrochemistry
• Oxidizing/Reducing Agents
– Compounds with large positive or negative E°
(standard reduction) values are frequently used in
electrochemistry (or in redox titrations)
– Example: MnO4- - E° (MnO4-(aq) + 8H+(aq) + 5e-) =
1.51 V is frequently used in redox titrations
– Why? Because if E° is high, it strongly reduces,
which makes it useful for oxidizing a wide variety of
compounds (e.g. Cu(s))
– Such a compound is called an oxidizing agent
(oxidizes other compounds)
Chapter 18 Electrochemistry
• Oxidizing/Reducing Agents – cont.
– Products of reduction reactions with large negative E°
values (e.g. Li(s), K(s)) are easily oxidized and can
therefore reduce other compounds
– Example: Al(s) - E° (Al3+(aq) + 3e-) = -1.66 V is
capable of reducing transition metals (reaction with
iron oxide is in thermite reaction)
Chapter 18 Electrochemistry
• Reduction Potential and Oxidation of Metals by
Acids
– Just as we can see which metals will oxidize or reduce
when pairing two metals (Ag/Cu example), we also
can see which metals will react in acid to produce
H2(g)
– Metals with E° (standard reduction) < 0 will react
with H+
– Examples: Fe, Pb, Sn, Ni, Cr, Zn, Al
– Metals with E° (standard reduction) > 0 will not react
with acid (except with HNO3 which is a stronger
oxidizing agent)
Chapter 18 Electrochemistry
• Reducing Potential Questions
– Given the table below, which of the following
oxidizing agents is strong enough to oxidize Ag(s) to
Ag+(aq) (under standard conditions)?
a) H+(aq) b) Co2+(aq) c) Cu2+(aq) d) Co3+(aq) e) Br2(l)
Reaction
Eº (V)
Ag+(aq) + e- ↔ Ag(s)
+0.799
Co2+(aq) + 2e- ↔ Co(s)
-0.277
Cu2+(aq) + 2e- ↔ Cu(s)
+0.337
Co3+(aq) + e- ↔ Co2+(aq)
+1.808
Br2(l) + 2e- ↔ 2Br- (aq)
+1.065