Transcript 3202 Chapter 19 - Eric G. Lambert School

Unit 4:
Electrochemistry
Chemistry 3202
CHAPTER 19
Electrochemical Cells (p. 757)

An electrochemical cell changes chemical
energy into electricity

AKA: Galvanic cell or Voltaic cell

These cells require spontaneous redox
reactions
Definitions

Electrode:


Cathode:


The electrode where oxidization occuds
Electrolyte


The electrode where reduction occurs
Anode


A conductor that carries electrons into and out of a cell
A solute that conduct electric current in an aqueous
solution
Salt bridge

A device containing a salt that completes a circuit between
half cells; maintains electrical neutrality without interfering
Electrochemical Cells
The ½ reactions must be separate such
that electron transfer occurs through an
external circuit or wire.
 A salt bridge containing an electrolyte
solution connects the ½ cells to complete
the circuit.

Demo:
Zn(s) + Cu2+(aq) → Zn2+(aq) + Cu(s)
Oxidation ½ rxn
Zn(s) → Zn2+(aq) + 2e-
Demo:
Zn(s) + Cu2+(aq) → Zn2+(aq) + Cu(s)
Cu2+(aq) + 2 e- → Cu(s)
Reduction ½ rxn
Demo:
Zn(s) + Cu2+(aq) → Zn2+(aq) + Cu(s)
+
-

KEY POINTS
 Oxidation
occurs at the anode
 Reduction occurs at the cathode
 Electrons move through the wire from the
anode to the cathode
 Cations move toward the cathode
 Anions move toward the anode
Electrochemical Cell Notation
anode|anodic ion(s)|| cathodic ion(s)|cathode
eg. Zn(s) + Cu2+(aq) → Zn2+(aq) + Cu(s)
Zn(s) | Zn2+(aq) | | Cu2+(aq) | Cu(s)
Electrochemical Cells

Sample Problem:
 Draw
a labeled electrochemical cell for the
reaction below. Indicate the cathode, anode,
direction of electron flow and the direction of
ion flow. Use cell notation to represent the
cell.
 Pb2+(aq) + 2 Ag(s) → Pb(s) + 2 Ag+(aq)
Inert Electrodes

Some redox reactions involve substances
that can’t act as electrodes
 (gases,
mixtures of ions)
These cells must use an inert electrode an electrode made from a material that is
neither a reactant nor a product.
 Common inert electrodes include graphite
(C(s)) and platinum (Pt(s))

Salt bridge
Inert Electrodes
eg.
Pb(s) |Pb2+(aq) || Fe3+(aq), Fe2+(aq) |Pt(s)
anode
cathode

P. 761 #’s 1 – 4
Standard Reduction Potentials
The difference in the potential energy of
the electrons at the anode and the
cathode is the electric potential, E, of the
cell
 electric potential is usually called cell
voltage or cell potential
 E is measured in volts

Standard Reduction Potentials
 The
potential of each ½ cell is found by
comparing many redox reactions
 The
table of standard reduction
potentials gives the voltage for each ½
reaction in the standard state
(1 mol/L and 25 °C)
Standard Reduction Potentials
The potentials in the table use hydrogen
as a reference
 All cell voltages are derived from the result
of two half reactions


DO NOT MULTIPLY VOLTAGES when
multiplying equations to cancel electrons!!
Standard Reduction Potentials
eg. Calculate the standard cell potential
for the electrochemical cell below:
2 I-
(aq)
+ Br2(l)
I2(s) + 2 Br
→
Br2(l) + 2 e-
→
2 I-
I2(s) + 2 e-
(aq) →
2 Br
-
(aq)
-
(aq)
+1.07 V
-0.54 V
+0.53 V
Standard Reduction Potentials
eg. Calculate the standard cell potential for
the electrochemical cell below:
2 Fe3+(aq) + Sn2+(aq)
2 Fe3+(aq) + 2 e -
Sn2+(aq)
→
→
→
Sn4+(aq) + 2 Fe2+(aq)
2 Fe2+(aq) + 0.77 V
Sn4+(aq) + 2 e-
- 0.15 V
+ 0.62 V
Standard Reduction Potentials
Standard reduction potentials can be
used to predict whether a reaction is
spontaneous
 Spontaneous redox reactions have a
POSITIVE cell potential
OR
 OA is higher in the table than RA

Standard Reduction Potentials
eg. Predict the cell voltage for:
a)
Cu(s) | Cu2+(aq) | | Ag+(aq) | Ag(s)
b)
2 Fe2+(aq) + Sn2+(aq) → 2 Fe3+(aq) + Sn(s)
p. 773
#’s 5 – 8
 Handout Electro #4

Identify the
anode and the
cathode
in this
electrochemical
cell.
[June 2004]
# 47
Predicting Redox
All redox reactions occur between the
strongest oxidizing agent (SOA) and the
strongest reducing agent (SRA).
 Spontaneous reactions will produce the
calculated voltage.
 Non-spontaneous reactions will require
slightly more than the calculated voltage.

Electrolytic Cells
An electrolytic cell converts electrical
energy to chemical energy
 The process that takes place in an
electrolytic cell is called electrolysis
 These cells are used to make nonspontaneous reactions occur

 (ie.
negative cell voltages)
Electrochemical Cell
electrons →
Anode
Cathode
(-)
(+)
cations
→
anions
←
Electrolytic Cell
+ Battery -
electrons 
Anode
Cathode
(-)
(+)
cations
→
anions
←
Electrolytic Cells
An electrolytic cell uses an external source
of electricity such as a battery rather than
a voltmeter.
 The battery forces a non-spontaneous
reaction to occur

Electrolytic cell
Carbon
electrode
Carbon
electrode
Cl-
Na+
Zn(s) + Cu2+(aq) → Zn2+(aq) + Cu(s)
V
-
Eo = +1.10 V
+
Electrochemical Cell
p. 780
Zn2+(aq) + Cu(s) → Zn(s) + Cu2+(aq)
Eo = -1.10 V
-
+
-
+
Electrolytic Cell
FIX p. 780
(connected to positive)
(connected to negative)
Electrolysis
 Molten
Salts (p. 776)
 Water
 Aqueous
NaCl
 Aqueous NiCl2(aq)
Handout Electro #5
Electroplating

Electrolytic cells may be used to
electroplate a thin layer of a more
desirable metal over a less desirable metal
 eg.
gold plated jewelry
 tin cans
-
CATHODE
Reduction
(+)
+
ANODE
Oxidation
(-)
Electroplating

object to be plated or covered is the
cathode
 (connected

to the negative terminal)
metal to be plated or form a thin layer is
the anode
 (connected
to the positive terminal)
 Pretty to positive – ugly to negative
p.795
Faraday’s Law

Faraday’s Law can be used to determine
the amount of substance produced or
consumed in an electrolytic cell

The amount of substance is directly
proportional to the quantity of electricity
that flows through the cell (or the electric
charge).
Faraday’s Law
The Coulomb (C) is the unit used to
measure the quantity of electricity (Q)
flowing in a circuit (or the electric charge).
 The charge on one mole of electrons is
96,500 Coulombs.
 This charge is Faraday’s constant:

F
= 96,500 C/mol
Faraday’s Law

FORMULA: Q = neF
Q
= charge in coulombs (C)
 ne = # of moles of electrons (mol)
 F = 96,500 C/mol
Faraday’s Law

one coulomb is the quantity of electric
charge that flows through a circuit in one
second if the current is one ampere.
 ie.

1 A = 1 C/sec
FORMULA: Q = I t
Q
= charge in coulombs (C)
 I = current in amperes (A)
 t = time in seconds (s)
Faraday’s Law


From Q = I t
and
I t = neF
Q = neF

eg. A spoon was copper plated using
CuSO4(aq) in an electrolytic cell that has
3.10 amps of electricity passing through it
for 2.50 h. What mass of copper will be
deposited?
I t = neF
(3.10)(2.5 x 60 x 60) = ne x (96,500)
0.289 mol = ne
Reaction:
Cu2+(aq) + 2 e- → Cu(s)
nCu = 0.289 mol e- x
mCu
1 Cu = 0.145 mol Cu
2 e= 0.145 mol Cu x 63.55 g/mol
= 9.21 g Cu
Faraday’s Law
eg. 0.423 g of solid Au is deposited on a
bracelet using a solution of gold(III) nitrate
and 2.00 A of current.
 For how long was the current applied?

n = 0.423g /196.97 g/mol
= 0.00215 mol Au
Au3+(ag) + 3e- → Au(s)
ne = 0.00215 x 3 e-/1 Au
= 0.00645 mol eI t = neF
(2.00)(t) = (0.00645)(96,500)
t = 311 s

Page 793 #’s 21 - 24

An electrolytic cell has a zinc strip anode
and a zinc strip cathode placed in a
solution of zinc sulfate. A current of 0.500
A is supplied for 900.0 seconds. What
mass of zinc is electroplated?

In an electrolytic cell, 0.061 g of Zn(s) was
plated in 10.0 minutes from a solution of
ZnCl2(aq). What current was used?

The titanium cathode in an electrolytic cell
increases in mass by 2.35 g in 36.5 min at
a current of 6.50 A. What is the charge on
the titanium ion? (Show workings.)

Handout Electro #6
Applications
Pp. 764 – 766, 787, 788
Dry cell
Battery
Primary battery
Secondary battery
Button cell battery Car battery (lead acid)


Voisey’s Bay STSE
Extra Practice (identifying redox)
Last one!!
5 CH3OH(l) + 2 MnO4-(aq) + 6 H+(aq) →
5 CH2O(l) + 2 Mn2+(aq) + 8 H2O(l)
Redox Stoichiometry
A
redox titration could be used to find an
unknown concentration of an OA or RA.
 The colors of the reactants and products
serve as indicators
 Calculations are the same as those in
acid-base stoichiometry.
Redox Stoichiometry
(see p. 742)
5 H2O2 + 2 MnO4- + 6 H+ → 5 O2 + 2 Mn2+ + 8H2O
- purple MnO4- changes to the colorless
Mn2+ until all of the H2O2 reacts
- the endpoint occurs when the purple
color remains
Redox Stoichiometry 0.174 mol/L
Sample Problem:
25.00 mL of a Fe2+ solution was titrated
with acidic 0.02043 mol/L K2Cr2O7
solution. The endpoint was reached when
35.55 mL of K2Cr2O7 solution had been
added. What was the molar concentration
of Fe2+
Cr2O72- + Fe2+ → Cr3+ + Fe3+
(balanced??)