Transcript Carbon-Zinc Dry Cell

Topic
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Table of Contents
Topic
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Topic 23: Electrochemistry
Basic Concepts
Additional Concepts
Electrochemistry: Basic Concepts
Topic
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Reviewing Redox
• Suppose you could separate the oxidation and
reduction parts of a redox reaction and cause
the electrons to flow through a wire.
• The flow of electrons in a particular direction
is called an electrical current.
Electrochemistry: Basic Concepts
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Reviewing Redox
• In other words, you are using a redox
reaction to produce an electrical current.
• This is what occurs in
a battery—one form of
an electrochemical cell
in which chemical
energy is converted to
electrical energy.
Electrochemistry: Basic Concepts
Topic
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Reviewing Redox
• You can reverse the process and use a current
to cause a redox
reaction to occur.
Electrochemistry: Basic Concepts
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Electrolysis
An electrochemical cell consists of two
electrodes and a liquid electrolyte.
One electrode, the cathode, brings electrons
to the chemically reacting ions or atoms in
the liquid; the other electrode, the anode,
takes electrons away.
The electrons act as chemical reagents at the
electrode surface.
The liquid electrolyte acts as the chemical
reaction medium.
Electrochemistry: Basic Concepts
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Electrolysis
• You can remember that
reduction always occurs at
the cathode and oxidation
always occurs at the anode
by studying this diagram.
Electrochemistry: Basic Concepts
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The Electrolysis Process
• Electrolysis takes place in a type of
electrochemical cell called an electrolytic
cell, in which a source of electricity, such as
a battery, is added to an external circuit
connecting the electrodes.
• The electrolysis process occurs when the
electrons are transferred between the electronic
conductors—the metal electrodes—and the
ions or atoms at the electrode surfaces.
Electrochemistry: Basic Concepts
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The Electrolysis Process
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Electrochemistry: Basic Concepts
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Electrolytic Cell
• Electrolysis, the splitting of compounds by
electricity, occurs when two electrodes, an
anode and a cathode, are inserted into a liquid
electrolyte such as molten sodium chloride
and connected to a source of electrical energy
such as a battery.
Electrochemistry: Basic Concepts
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Electrolytic Cell
Electrochemistry: Basic Concepts
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Electroplating
• Reduction of silver ions onto cheaper metals
forms silverplate.
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Electrochemistry: Basic Concepts
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Electroplating
• The object to be plated is made the cathode.
• At the pure silver
anode, oxidation of
silver metal to
silver ions replaces
the silver ions
removed from the
solution by plating
at the cathode.
Electrochemistry: Basic Concepts
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Electrolytic Cleaning
Electrolysis can be used to clean objects by
pulling ionic dirt away from them.
The electrolysis cell for this cleaning process
includes a cathode that is the object itself, a
stainless steel anode, and an alkaline
electrolyte.
When an electric current is run through the
cell, the chloride ions are drawn out.
Hydrogen gas forms and bubbles out, helping
to loosen corrosion products.
Electrochemistry: Basic Concepts
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Electrophoresis
• Electrophoresis is another electrochemical
process that was used to restore some of the
ceramic and organic artifacts from the Titanic.
• Electrophoresis involves placing an
artifact in an electrolyte solution between
positive and negative electrodes and
applying a current.
Electrochemistry: Basic Concepts
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Electrophoresis
• The current breaks up salts, dirt, and other
particles as their charged components
migrate to the electrodes.
• Electrophoresis is also used in laboratories
to separate and identify large molecules.
Electrochemistry: Basic Concepts
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Potential Difference
• Why do the electrons travel in one direction
and not in the reverse?
• The electron pressure at the cathode is kept
low by the reduction reaction, and the
electrons flow from a region of high
pressure (negative potential at the anode) to
a region of low pressure (positive potential
at the cathode). This potential difference
between the electrodes causes an electrical
current to flow.
Electrochemistry: Basic Concepts
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Potential Difference
• In this model of a lemon battery, the level of
the electron sea is raised or lowered by the
chemical reactions at the electrode surfaces,
creating a potential
difference across
the battery.
Electrochemistry: Basic Concepts
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Potential Difference
• A spontaneous oxidation reaction raises the
electron pressure (potential) at the anode,
and a spontaneous reduction reaction
reduces the
pressure at the
cathode.
Electrochemistry: Basic Concepts
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Potential Difference
Electrochemistry: Basic Concepts
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Potential Difference
• Because the redox reactions that take place
during electrolysis are not spontaneous, a
battery is needed to pump electrons from an
area of low potential to one of high potential.
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Electrochemistry: Basic Concepts
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Potential Difference
Electrochemistry: Basic Concepts
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Galvanic Cells
An electrochemical cell in which an
oxidation-reduction reaction occurs
spontaneously to produce a potential
difference is called a galvanic cell.
In a Galvanic cell, chemical energy is
converted into electrical energy.
Galvanic cells are sometimes called voltaic
cells; both terms refer to the same device.
A galvanic cell that has been packaged as a
portable power source is often called a battery.
Electrochemistry: Basic Concepts
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Batteries Perform Work
• When a simple galvanic cell does useful work,
it is called a battery.
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Electrochemistry: Basic Concepts
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Batteries Perform Work
• If the external circuit is connected with a wire,
electrons flow from the
site of oxidation at the
magnesium strip and
through the LED to the
surface of the copper
strip, where reduction of
Cu2+ ions takes place.
Electrochemistry: Basic Concepts
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Batteries Perform Work
• The voltage pushes electrons through the LED,
causing it to light up.
Electrochemistry: Basic Concepts
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Batteries Perform Work
Electrochemistry: Basic Concepts
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Modern Batteries
• Modern batteries come in a wide variety
of sizes, shapes, and strengths.
• Each type of battery serves a different
purpose.
Electrochemistry: Basic Concepts
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Modern Batteries
Electrochemistry: Basic Concepts
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Modern Batteries
• Although the term battery usually refers to
a series of galvanic cells connected together,
some batteries have only one such cell.
• Other batteries may have a dozen or more
cells.
Electrochemistry: Basic Concepts
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Modern Batteries
• When you put a battery into a flashlight,
radio, or CD player, you complete the
electrical circuit of a galvanic cell(s),
providing a path for the electrons to flow
through as they move from the reducing
agent (the site of oxidation) to the oxidizing
agent the site of (the reduction).
Electrochemistry: Basic Concepts
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Carbon-Zinc Dry Cell
• Whenever you put two or more common D
batteries into a flashlight, you are connecting
them in series.
• They have to be placed in the correct order
so that electrons flow through both cells.
Electrochemistry: Basic Concepts
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Carbon-Zinc Dry Cell
• These relatively inexpensive batteries are
carbon-zinc galvanic cells, and they come in
several types, including standard, heavy-duty,
and alkaline.
• This type of battery is often called a dry cell
because there is no aqueous electrolyte
solution; a semisolid paste serves that role.
Electrochemistry: Basic Concepts
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Carbon-Zinc Dry Cell
• A standard D battery is shown both whole
and cut in half to reveal the structure of the
carbon-zinc dry cell.
• Beneath the outside paper
cover of the battery is a
cylinder casing made of
zinc.
• The zinc serves as the
anode and will be oxidized
in the redox reaction.
Electrochemistry: Basic Concepts
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Carbon-Zinc Dry Cell
• The carbon rod in the center of the cylinder—
surrounded by a moist, black paste of
manganese (IV)
oxide (MnO2) and
carbon black—acts
as a cathode.
• Ammonium chloride
(NH4Cl) and zinc
chloride (ZnCl2) serve
as electrolytes.
Electrochemistry: Basic Concepts
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Carbon-Zinc Dry Cell
• Alkaline batteries contain potassium
hydroxide (KOH) in place of the ammonium
chloride electrolyte,
and they maintain a
high voltage for a
longer period of time.
Electrochemistry: Basic Concepts
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Carbon-Zinc Dry Cell
• The flow of electrons from the zinc cylinder
through the electrical circuits of an appliance
and back into the battery provides the
electricity needed to power a flashlight, radio,
CD player, toy, clock, or other item.
• When electrons leave the casing, zinc metal
is oxidized.
Electrochemistry: Basic Concepts
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Carbon-Zinc Dry Cell
• The reactions in the carbon rod and the paste
are much more complex, but one major
reduction that takes place is that of
manganese in manganese (IV) oxide.
Electrochemistry: Basic Concepts
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Carbon-Zinc Dry Cell
• In this reaction, the oxidation number of
manganese is reduced from
.
• Adding the two half-reactions together gives
the major redox reaction taking place in a
carbon-zinc dry cell.
Electrochemistry: Basic Concepts
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Carbon-Zinc Dry Cell
Each galvanic cell in a lead-acid battery has
two electrodes—one made of a lead (IV)
oxide (PbO2) plate and the other of spongy
lead metal.
In each cell, lead metal is oxidized as lead
(IV) oxide is reduced.
The lead metal is oxidized to Pb2+ ions as it
releases two electrons at the anode.
The Pb4+ ions in lead oxide gain two
electrons, forming Pb2+ ions at the cathode.
Electrochemistry: Basic Concepts
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Carbon-Zinc Dry Cell
• The Pb2+ ions combine with SO42– ions from
the dissociated sulfuric acid in the
electrolyte solution to form lead (II) sulfate
at each electrode.
• Thus, the net reaction that takes place when
a lead-acid battery is discharged results in
the formation of lead sulfate at both of the
electrodes.
Electrochemistry: Basic Concepts
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Lead Storage Batteries
Electrochemistry: Basic Concepts
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Lead Storage Batteries
• The reaction that occurs during discharge of a
lead-acid battery is spontaneous and requires
no energy input.
• The reverse reaction, which recharges the
battery, is not spontaneous and requires an
input of electricity from the car’s alternator.
Electrochemistry: Basic Concepts
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Lead Storage Batteries
• Current enters the battery and provides
energy for the reaction in which lead sulfate
and water are converted into lead (IV) oxide,
lead metal, and sulfuric acid.
Electrochemistry: Basic Concepts
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Experimental Batteries
• Two new experimental types of batteries for
use in electric cars show early promise as
candidates.
• One is a rechargeable, nickel-metal hydride
or NiMH battery.
• This type of battery is less toxic and has a
higher storage capacity than the batteries
now used in electric cars.
Electrochemistry: Basic Concepts
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Experimental Batteries
• Another experimental battery is a lithium
battery with a water-based electrolyte.
• Lithium is more easily oxidized than any other
metal but has a drawback that has limited its
use in batteries: it explodes violently when it
comes into contact with water.
• Lithium is used in some batteries to power
camcorders, but they require an expensive,
nonaqueous electrolyte.
Electrochemistry: Basic Concepts
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Aqueous Lithium Battery
• How can a lithium battery have an aqueous
electrolyte? Two facets of the construction
of this new battery keep the lithium metal
from reacting with water.
• First, the lithium is in the form of
individual atoms embedded in a material
such as manganese (IV) oxide, rather than
as a solid metal.
Electrochemistry: Basic Concepts
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Aqueous Lithium Battery
• Second, the electrolyte is full of dissolved
lithium salts, so the lithium ions that are
produced travel to the site of reduction
without reacting with water.
Electrochemistry: Basic Concepts
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Aqueous Lithium Battery
Basic Assessment Questions
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Question 1
What term describes a battery that is not
rechargeable?
Basic Assessment Questions
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Answer
primary battery
Basic Assessment Questions
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Question 2
What element is oxidized in most dry cells?
Basic Assessment Questions
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zinc
Answer
Basic Assessment Questions
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Question 3
Distinguish between a voltaic cell and an
electrolytic cell.
Basic Assessment Questions
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Answer
spontaneous redox reaction;
nonspontaneous redox reaction by electrolysis
Electrochemistry: Additional Concepts
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Additional Concepts
Electrochemistry: Additional Concepts
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Calculating Cell Potential
• The two reduction half-reactions in this
example represent the half-cells of a
voltaic cell.
• The standard reduction potentials for each
half-reaction are given.
Electrochemistry: Additional Concepts
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Calculating Cell Potential
Determine the overall cell reaction and the
standard cell potential.
Write the cell chemistry using cell notation
with vertical lines separating components.
Note that reduction of iodine has the higher
reduction potential.
This half-reaction will proceed in the forward
direction as a reduction.
The iron half-reaction will proceed in the
reverse direction as an oxidation.
Electrochemistry: Additional Concepts
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Calculating Cell Potential
• Rewrite the half-reactions in the correct
direction.
Electrochemistry: Additional Concepts
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Calculating Cell Potential
• Balance the reaction if necessary. Note
that this reaction is balanced as written.
• Calculate cell standard potential.
Electrochemistry: Additional Concepts
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Calculating Cell Potential
• Write the reaction using cell notation.
• When representing a reaction in cell
notation, the species in the oxidation halfreaction are written first in the following
order:
or in this case,
.
Electrochemistry: Additional Concepts
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Calculating Cell Potential
• The species in the reduction half-reaction
are written next in the order
or in this case,
.
• Therefore, the complete cell is represented
as
.
Electrochemistry: Additional Concepts
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Predicting the Spontaneity
of a Reaction
• Predict whether the following redox reaction
will occur spontaneously.
Electrochemistry: Additional Concepts
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Predicting the Spontaneity
of a Reaction
• Write the half-reactions. Note that the
coefficients are simplified.
• Find the standard cell potential, using E0
values from Table 21-1 in your textbook.
Electrochemistry: Additional Concepts
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Predicting the Spontaneity
of a Reaction
• The voltage is negative, so the reaction is
not spontaneous.
• The reverse reaction will occur spontaneously.
Additional Assessment Questions
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Question 1
For each pair of half-reactions, write the
balanced equation for the overall cell reaction,
calculate the standard cell potential, and
express the reaction using cell notation. Use E0
values from Table 21-1 in your textbook.
Mg2+(aq) + 2e– → Mg(s)
Pd2+(aq) + 2e– → Pd(s)
Additional Assessment Questions
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Answer
Additional Assessment Questions
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Question 2
Calculate the cell potential to determine if each
of these redox reactions is spontaneous.
Additional Assessment Questions
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Question 2a
2Ag+(aq) + Co(s) → Co2+(aq) + 2Ag(s)
Answer 2a
+ 1.08V; spontaneous
Additional Assessment Questions
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Question 2b
Cu(s) + Cu2+(aq) → 2Cu+(aq)
Answer 2b
– 0.368V; not spontaneous
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