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Chemistry for Engineers
Learning Material 2
(4th Week)
Contents:
Electrochemical Energy
a. Oxidation: Reduction and Reaction
b. Batteries in Engineering Design
c. Chemical reactions responsible to corrosion
d. Electrolysis and Stoichiometry
-
Current and Charge
-
Calculations using current, mass, or time.
Prepared by:
Dr. Julieta P. Donato
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GUIDE ON HOW TO USE THE MODULE
A. For Faculty
The primary objective of this module is to address the learning of the students in
Blended learning approach. The author aims that through this module the school will meet
the expectation of the learners.
To the faculty of PLSP, you can use this module to supplement the learning of the
students. However, do not use this module to gain money from the students or sell this to
other instructors.
You can also use this module as a reference in preparing your instructional
materials like PowerPoint presentation and others, for synchronous/asynchronous learning
or for preparing the same module.
B. For Learners
This module is designed to engage you as a learner, and to guide you as you finish
this subject, Chemistry for Engineers.
In using this module, read every topic with comprehension. Make sure that you
have understand every lesson, because there are activities for you to answer after every
topic and all of the questions are answerable based on what you have understood in the
topic.
You can optionally share this module to your classmates or upload this online, just
make sure that it is for educational purposes and not for selling, and the credits must belong
to the author.
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Electrochemical Energy
Electrochemical Energy
Electrochemical energy is what we normally call the conversion of chemical energy into electrical energy
or vice versa. This includes reactions transferring electrons, redox reactions (reduction- oxidation).
Oxidation – Reduction Concept
All electrochemical processes involved the movement of electrons from one chemical species to another
in an oxidation- reduction (redox) reaction. in any redox process,





Oxidation is the loss of electrons, and reduction is the gain of electrons. Oxidation and reduction
occur simultaneously.
the oxidizing agent does the oxidizing by taking electrons from the substance being oxidized.
The reducing agent does the reducing by giving electrons the substance being reduced.
therefore, the oxidizing agent is reduced and the reducing agent is oxidized.
the oxidized substance ends up with a higher (more positive or less negative) oxidation number,
and the reduced substance ends up with a lower (less positive or more negative) oxidation
number.
the total number of electrons gained by the atoms/ ions of the oxidizing agent equals the total
number lost by atoms/ions of the reducing agent.
Reduction, when a substance receives one electron.
Oxidation when a substance gives away one electron.
There always has to be a balance of substances that give away and substance that receives electrons since
electrons cannot exist on their own without any bindings. This means that if a reduction is taking place also
an oxidation has to take place.
Photo taken from:
https://www.priyamstudycentre.com/wp-content/uploads/2021/09/redox-reaction.png
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Half-reaction method for balancing redox reactions
The half-reaction method for balancing redox reactions is commonly used for studying electrochemistry
for several reasons:



It separates the overall redox reaction into oxidation and reduction half-reactions, which reflect
their physical separation in electrochemical cells.
It is easier to apply to reactions in acidic or basic solution, which is common in cells.
It (usually) does not require assigning oxidation numbers (in cases where the half-reactions are not
obvious, we assign oxidation numbers to determine which atoms undergo a change and write half
reactions with the species that contain those atoms)
Steps in the Half Reaction Method
The balancing process begins with the “skeleton” ionic reaction that consists of only species that are
oxidized and reduced. Here are the steps in the half reaction method:
1. Divide the skeleton reaction into two half-reactions, each of which contains the oxidized and
reduced form of one of the species: if the oxidized form of a species is on the left side, the reduced
form must be on the right, and vice versa.
2.
Balance the atoms and charges in each half-reaction.

Atoms are balanced in this order: atoms other than O and H, then O, then H.
 Charge is balanced by adding electrons (𝑒 − )To the left side in the reduction half reaction
because the reactant gains them and to the right side in the oxidation half reaction because
the reactant loses them.
3. If necessary, multiply one or both half-reactions by an integer so that number of electrons gained
in reduction is equal to the number of electrons lost in oxidation.
4.
Add the balanced half reactions, and include states of matter.
5. Check that the atoms and charges are balanced.
Electrochemical cells
Electrochemical cells either generate electrical energy from chemical reactions or they use electrical energy
to cause chemical reactions.
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Anode is the negative or reducing electrode that releases electrons to the external circuit and oxidizes during
and electrochemical reaction.
Cathode is the positive or oxidizing electrode that acquires electrons from the external circuit and is
reduced during the electrochemical reaction.
There are basically to types of cells used for electrochemical conversion:
1) The galvanic cell (also called a voltaic cell) that converts chemical energy into electrical energy,
by a spontaneous reaction. A standard house hold battery contains one or more galvanic cells.
2) The electrolytic cell that converts
Electrical energy is used to fuel the reaction.
electrical
energy
into
chemical
energy.
Electrolysis
A process whereby electrical energy is converted directly into chemical energy is called electrolysis; i.e.,
an electrolytic process.
By virtue of their combined chemical energy, the products of an electrolytic process often react
spontaneously with one another, reproducing the substances that were reactants and were therefore
consumed during the electrolysis.
Photo taken from: https://useruploads.socratic.org/4FQ7RIXeSJm071mxtuzA_elecaluminum.GIF
Electrochemical energy storage
Electrochemical energy storage is a method used to store electricity in a chemical form. This storage
technique benefits from the fact that both electrical and chemical energy share the same carrier, the electron.
This common point allows limiting the losses due to the conversion from one form to another.
Common forms for electrochemical storage and conversion
 Batteries
 Fuel cells
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Electrochemical Processes in Batteries
Because of their compactness and mobility, by, and in our increasingly wireless world, that role is
growing. In general, a battery consists of self-contained voltaic cells arranged in series (plus-to-minus-toplus- and so on), so that their individual voltages are added.
Primary (Non-rechargeable) Batteries
A primary battery cannot be recharged, so it is discarded when the cell reaction has reached
equilibrium, that is when the battery is “dead”.
1.
Alkaline battery - the ubiquitous alkaline battery has a zinc anode case that houses a mixture of
mno2 (the oxidizing agent) and an alkaline taste of KOH and water. The cathode is an inactive
graphite rod. The alkaline battery powers portable radios, toys, flashlights, and so on, is safe, has a
long shelf life, and comes in many sizes.
2.
Mercury and Silver (button) batteries - both mercury and silver batteries use a zinc container as
the anode ( reducing agent) in a basic medium. The mercury battery employs HgO as the oxidizing agent,
the silver uses Ag2O2, and both use a steel can around the cathode. The solid reactants are compacted with
KOH and separated with moist paper. Both cells are manufactured as button-size batteries. The mercury
cell is used in calculators and the silver cell in watches, cameras, heart pacemakers, and hearing aids because
it is a very steady output. Disadvantages are the toxicity of discarded mercury and the high cost of silver.
Secondary (Rechargeable) batteries
In contrast to a primary battery, a secondary battery is rechargeable; when it runs down, Electrical
energy is supplied to reverse the cell reaction and for more reactant. In other words, in a secondary battery,
the vault excels are periodically converted to electrolytic cells to restore equilibrium concentrations of the
cell components.
1.
Lead-Acid battery - a typical lead-acid battery has six cells connected in series, each of which
delivers about 2.1 V for a total of about 12 V. Each cell contains two lead grids packed with high
surface-surface-area (spongy) Pb in the anode and high surface area pbo2 in the cathode.

Discharging. When the cells discharge as voltaic cells, it generates electrical energy.

Recharging. When the cell recharges as an electronic sell it uses electrical energy (supplied
by the vehicle's charging system) and the half cell and overall reactions are reversed.
Car and truck owners have relied on the lead acid battery for over a century to provide the
large burst of current needed to start the engine– and to do it in hot and cold weather for years. The
main problems with lead-acid batteries or loss of capacity due to corrosion of the positive Pb grid,
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detachment of the active material due to normal mechanical bumping, and formation of large pbso4,
crystals that hinder recharging.
2.
Nickel-Metal Hydride battery - concerns about the toxicity of cadmium in the nickel-cadmium
(nicad) battery have led to its replacement by the nickel-metal hydride battery. The anode half-reaction
oxidizes the hydrogen absorbed within a metal alloy in a basic KOH electrolyte, while nickel (III) in the
form of Nio(OH) is reduced at the cathode. Nickel metal hydride battery is common in cordless razors,
camera flash units, and power tools. It is lightweight, has high power, and is non-toxic, but it discharges
significantly during storage.
3.
Lithium-Ion battery - has an anode of Li atoms that lie between sheets of graphite. The cathode
is a lithium metal oxide such as limn2o4 or licoo2, And a typical electrolyte is 1 M lipf6 in an organic
solvent, such as dimethyl carbonate mixed with methyl ethyl carbonate. Electrons flow through the circuit
while solvated positive Li ions flow from anode to cathode within the cell. The cell reaction is reversed
during discharge. The lithium-ion battery hours countless laptop computers, cell phones, & camcorders. If
the drawbacks are cost and flammability of the organic solvent.
Fuel cells
In contrast to primary and secondary batteries, a fuel cell, sometimes called a flow battery, is not
self-contained. The reactants (usually fuel and oxygen) enter the cell, and the products leave, generating
electricity through controlled combustion of fuel. The fuel does not burn because, as in other voltaic cells,
the half-reactions are separated, and circuit. Hydrogen fuel cells have been used for years to provide
electricity and water during space flights. Similar once have begun to supply electric power for residential
needs, and every major car manufacturer has a fuel-cell prototype. Fuel cells produce no pollutants and
convert about 75% of a fuel’s bond energy into power, compared to 40% for a coal-fired power plant, and
25% for a gasoline engine.
Corrosion
Corrosion is a galvanic process by which metals deteriorate through oxidation—usually but not
always to their oxides. For example, when exposed to air, iron rusts, silver tarnishes, and copper and brass
acquire a bluish-green surface called a patina. Of the various metals subject to corrosion, iron is by far the
most important commercially.
Under ambient conditions, the oxidation of most metals is thermodynamically spontaneous, with
the notable exception of gold and platinum. Hence it is actually somewhat surprising that any metals are
useful at all in Earth’s moist, oxygen-rich atmosphere. Some metals, however, are resistant to corrosion for
kinetic reasons. For example, aluminum in soft-drink cans and airplanes is protected by a thin coating of
metal oxide that forms on the surface of the metal and acts as an impenetrable barrier that prevents further
destruction. Aluminum cans also have a thin plastic layer to prevent reaction of the oxide with acid in the
soft drink. Chromium, magnesium, and nickel also form protective oxide films. Stainless steels are
remarkably resistant to corrosion because they usually contain a significant proportion of chromium, nickel,
or both.
In contrast to these metals, when iron corrodes, it forms a red-brown hydrated metal oxide
(Fe2O3⋅xH2OFe2O3⋅xH2O), commonly known as rust, that does not provide a tight protective film.
Instead, the rust continually flakes off to expose a fresh metal surface vulnerable to reaction with oxygen
and water. Because both oxygen and water are required for rust to form, an iron nail immersed in
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deoxygenated water will not rust—even over a period of several weeks. Similarly, a nail immersed in an
organic solvent such as kerosene or mineral oil will not rust because of the absence of water even if the
solvent is saturated with oxygen.
Figure: Rust, the Result of Corrosion of Metallic Iron. Iron is oxidized to Fe2+(aq) at an anodic site on the
surface of the iron, which is often an impurity or a lattice defect. Oxygen is reduced to water at a different
site on the surface of the iron, which acts as the cathode. Electrons are transferred from the anode to the
cathode through the electrically conductive metal. Water is a solvent for the Fe 2+ that is produced initially
and acts as a salt bridge. Rust (Fe2O3•xH2O) is formed by the subsequent oxidation of Fe2+ by atmospheric
oxygen.
The corrosion of iron
The most common and economically destructive form of corrosion is the rusting of iron. About
25% each year for replacing steel in which the iron has corroded. Rust is not a direct product of the reaction
between iron and oxygen but arises through a complex electromechanical process. Let us look at the facts
of iron corrosion and then use the features of voltaic cell to explain them:
1.
2.
3.
4.
5.
6.
Iron does not rust in dry air; moist must be present.
Iron does not rust in air-free water; oxygen must be present.
Iron loss and rust formation occur at different places on the same object.
Iron rusts more quickly at low ph (high H+)
Iron rusts more quickly in ionic solutions.
Iron rusts more quickly in contact with a less active metal (such as Cu) and more slowly in contact
with a more active metal (such as Zn)
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Protecting Against the Corrosion of Iron
Corrosion is prevented by eliminating corrosive factors. Washing road salt off auto bodies removes
ions. Painting an object keeps out O2 and moisture. Plating chromium on plumbing fixtures is a more
pertinent method, as is “blueing” of gun barrels and other steel objects, in which a coating of Fe3O4
(magnetite) is bonded to the surface. The final point regarding corrosion concerns the relative activity of
other metals in contact with iron. The essential idea is that when functions as both anode and cathode in the
rusting process, it is lost only at the anode. Thus,
1. Corrosion increases when iron behaves more like the anode. When, such as copper, it loses
electrons more readily. For example, when iron plumbing is connected directly to copper plumbing,
the Iron pumps road rapidly. Non-conducting rubber or plastic spacers are placed between the
metals to avoid this problem.
2.
Corrosion decreases when iron behaves more like the cathode. In cathodic protection, the
corrosion, iron makes contact with a more active metal, such as zinc. The iron becomes the cathode
and remains intact, while the zinc acts as the anode and loses electrons. Coating steel with a
“sacrificial” layer of Zinc is called galvanizing. In addition to blocking physical contact with H2O
and O2, the zinc (or other active metal) is “sacrificed” (oxidized) instead of the iron. Sacrificial
anodes are used underwater and underground to protect iron and steel pipes, tanks, oil rigs, and so
on. Magnesium and aluminum are often used because they are much more active than iron, and
thus, act as the anode. Moreover, they form adherent oxide coatings, which slow their corrosion.
Figure: The Use of a Sacrificial Electrode to Protect Against Corrosion. Connecting a magnesium rod to an
underground steel pipeline protects the pipeline from corrosion. Because magnesium (E° = −2.37 V) is
much more easily oxidized than iron (E° = −0.45 V), the Mg rod acts as the anode in a galvanic cell. The
pipeline is therefore forced to act as the cathode at which oxygen is reduced. The soil between the anode
and the cathode acts as a salt bridge that completes the electrical circuit and maintains electrical neutrality.
As Mg(s) is oxidized to Mg2+ at the anode, anions in the soil, such as nitrate, diffuse toward the anode to
neutralize the positive charge. Simultaneously, cations in the soil, such as H+ or NH4+, diffuse toward the
cathode, where they replenish the protons that are consumed as oxygen is reduced. A similar strategy uses
many miles of somewhat less reactive zinc wire to protect the Alaska oil pipeline.
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Prophylactic Protection
One of the most common techniques used to prevent the corrosion of iron is applying a protective
coating of another metal that is more difficult to oxidize. Faucets and some external parts of automobiles,
for example, are often coated with a thin layer of chromium using an electrolytic process. With the increased
use of polymeric materials in cars, however, the use of chrome-plated steel has diminished in recent years.
Similarly, the “tin cans” that hold soups and other foods are actually consist of steel container that is coated
with a thin layer of tin. While neither chromium nor tin metals are intrinsically resistant to corrosion, they
both form protective oxide coatings that hinder access of oxygen and water to the underlying steel (iron
alloy).
Figure: Galvanic Corrosion. If iron is in contact with a more corrosion-resistant metal such as tin, copper,
or lead, the other metal can act as a large cathode that greatly increases the rate of reduction of oxygen.
Because the reduction of oxygen is coupled to the oxidation of iron, this can result in a dramatic increase
in the rate at which iron is oxidized at the anode. Galvanic corrosion is likely to occur whenever two
dissimilar metals are connected directly, allowing electrons to be transferred from one to the other.
Electrolysis
process by which electric current is passed through a substance to effect a chemical change. The
chemical change is one in which the substance loses or gains an electron (oxidation or reduction). The
process is carried out in an electrolytic cell, an apparatus consisting of positive and negative electrodes held
apart and dipped into a solution containing positively and negatively charged ions. The substance to be
transformed may form the electrode, may constitute the solution, or may be dissolved in the solution.
Electric current (i.e., electrons) enters through the negatively charged electrode (cathode);
components of the solution travel to this electrode, combine with the electrons, and are transformed
(reduced). The products can be neutral elements or new molecules. Components of the solution also travel
to the other electrode (anode), give up their electrons, and are transformed (oxidized) to neutral elements
or new molecules. If the substance to be transformed is the electrode, the reaction is often one in which the
electrode dissolves by giving up electrons.
Electrolysis is used extensively in metallurgical processes, such as in extraction (electrowinning)
or purification (electrorefining) of metals from ores or compounds and in deposition of metals from
solution (electroplating). Metallic sodium and chlorine gas are produced by the electrolysis of molten
sodium chloride; electrolysis of an aqueous solution of sodium chloride yields sodium hydroxide and
chlorine gas. Hydrogen and oxygen are produced by the electrolysis of water.
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Stoichiometry
Stoichiometry is a section of chemistry that involves using relationships between reactants and/or
products in a chemical reaction to determine desired quantitative data. In Greek, stoikhein means element
and metron means measure, so stoichiometry literally translated means the measure of elements. In order
to use stoichiometry to run calculations about chemical reactions, it is important to first understand the
relationships that exist between products and reactants and why they exist, which require understanding
how to balance reactions.
Stoichiometry of Electrolysis: The Relation Between Amounts of Charge and Products
Since charge flowing through an electrolytic cell yields products at the electrodes. It makes sense
that the more charge that flows, the more product that forms. In fact, this relationship was first
determined experimentally in the 1830s by Michael Faraday and is referred to as– faraday’s law of
electrolysis: the amount of substance produced at each electrode is directly proportional to the quantity of
charge flowing through the cell.
Each balanced half-reaction shows the amount (mol) of reactant, electrons, and product involved
in the change, so it contains the information we need to answer such questions as “how much material
will form from a given quantity of charge?” Or conversely “How much charge is needed to produce a
given amount of material?”
To apply Faraday's law,
1.
2.
3.
Balance the reaction to find the amount (mol) of electrons needed per mole of product.
𝐶
Use the faraday constant (𝐹 = 9.65 𝑥 104
𝑒 − ) to find the quantity of charge.
𝑚𝑜𝑙
Use the molar mass to find the charge needed for a given mass of product.
Measuring current to find charge
We cannot measure charge directly but we can measure current, the charge flowing per unit time.
The SI unit of current is the ampere (A), which is defined as a charge of 1 coulomb flowing through a
conductor and 1 second:
1 Ampere = 1coloumb/second
or
1A = 1C/s
The current multiplied by the time gives the charge:
Current x time = charge
or
A x s = C/s x s = C
Therefore, by measuring the current and the time during which the current flows, we find the charge,
which relates to the amount of product.
Problems involving stoichiometry of Electrolysis
Problems based on faraday's law of an ask you to calculate the current, mass of material, or time.
As we said, the electrode half-reaction provides the key to solving these problems because it is related to
the mass for a certain quantity of charge.
Here's a typical problem and practical electrolysis: how long does it take to produce 3.0g of
Cl2(g) by electrolysis of aqueous NaCl using a power supply with a current of 12A? The problem asks for
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the time needed to produce a certain mass, so let us first relate mass to amount (mol) of electrons to find
the charge needed, then we will relate the charge to the current to find the time.
The half-reaction tells us that 2 mol of electrons are lost to form 1 mol of Cl2 and we will use this
relationship as a conversion factor:
𝟐𝑪𝒍− (𝒂𝒒) → 𝑪𝒍𝟐 (𝒈) + 𝟐𝒆−
We convert the given mass of 𝐶𝑙2 to the amount of 𝐶𝑙2 , use the conversion factor from half-reaction, and
multiply by the Faraday constant to find the total charge:
𝑪𝒉𝒂𝒓𝒈𝒆 (𝑪) = 𝟑. 𝟎𝒈 𝑪𝒍𝟐 𝒙
𝟏 𝒎𝒐𝒍 𝑪𝒍𝟐
𝟐 𝒎𝒐𝒍 𝒆−
𝟗. 𝟔𝟓 𝒙 𝟏𝟎𝟒 𝑪
𝒙
𝒙
= 𝟖. 𝟐 𝒙 𝟏𝟎𝟑 𝑪
𝟕𝟎. 𝟗𝟎 𝒈 𝑪𝒍𝟐 𝟏 𝒎𝒐𝒍 𝑪𝒍𝟐
𝟏 𝒎𝒐𝒍 𝒆−
Now, we use the relationship between charge and current to find the time needed:
𝑻𝒊𝒎𝒆 (𝒔) =
𝒄𝒉𝒂𝒓𝒈𝒆 (𝑪)
𝑪
𝒄𝒖𝒓𝒓𝒆𝒏𝒕 (𝑨 𝒐𝒓 )
𝒔
= 𝟖. 𝟐 𝒙 𝟏𝟎𝟑 𝑪 𝒙
𝟏𝒔
= 𝟔. 𝟖 𝒙 𝟏𝟎𝟐 (~𝟏𝟏 𝒎𝒊𝒏)
𝟏𝟐𝑪
12