Saint Louis University

SCHOOL OF ENGINEERING AND ARCHITECTURE

Department of Chemical Engineering

MODULE 2
This Module includes the following Units:

Unit 1: Oxidation-Reduction (Redox)

Unit 2: Rusting of Metals
Unit 3: Mechanical Properties of Chocolate – How Strong is your Chocolate

Unit 1: Oxidation-Reduction (Redox)

UNIT LEARNING OUTCOME

● Execute proficiency to demonstrate, calculate, analyze, and balance chemical
equations involving oxidation-reduction reactions

Engage/Explain

What is Oxidation-Reduction (Redox) Reaction?

In a restricted sense, the term “oxidation” refers to a reaction which involves the
combination of other substances with oxygen. The term “reduction”, on the other hand,
refers to the removal of oxygen from its compounds.
In the broadest sense, the concept of oxidation and reduction is associated with the
electrical state of the element. Oxidation refers to a reaction in which an element increases
in oxidation state due to loss of electrons. Reduction refers to a reaction in which an element
decreases in oxidation state due to gain of electrons. In many reactions, the oxidation states
of elements do not change, but in many others, the oxidation states of elements do change.
These changes in oxidation states are a consequence of electron transfer from the structure
of one atom to that of another. Therefore, oxidation and reduction must occur
simultaneously. Reactions wherein oxidation and reduction, according to the broad
definition, takes place simultaneously are called OXIDATION – REDUCTION reactions,
sometimes abbreviated as REDOX.
In a given reaction, the substance responsible for oxidation is called the oxidizing
agent and the substance responsible for reduction is called the reducing agent. The oxidizing
agent contains an element capable of taking up electrons. The oxidizing agent causes the
oxidation of a given element by removing electrons from that element but in so doing is itself
reduced. The reducing reagent causes the reduction of a given element by giving up
electrons to that element but in so doing is itself oxidized. In oxidation-reduction, the transfer
of electron is from the reducing agent to the oxidizing agent. For example, the typical
oxidation-reduction represented in the ionic form:
𝑍𝑛0 𝐶𝑢2+ 𝑍𝑛2+ 𝐶𝑢0
+ → +
𝑅𝑒𝑑𝑢𝑐𝑖𝑛𝑔 𝑎𝑔𝑒𝑛𝑡 𝑂𝑥𝑖𝑑𝑖𝑧𝑖𝑛𝑔 𝑎𝑔𝑒𝑛𝑡 𝑂𝑥𝑖𝑑𝑎𝑡𝑖𝑜𝑛 𝑝𝑟𝑜𝑑𝑢𝑐𝑡 𝑅𝑒𝑑𝑢𝑐𝑡𝑖𝑜𝑛 𝑝𝑟𝑜𝑑𝑢𝑐𝑡

1
Property of and for the exclusive use of SLU. Reproduction, storing in a retrieval system, distributing, uploading or posting online, or transmitting in any form or by any
means, electronic, mechanical, photocopying, recording, or otherwise of any part of this document, without the prior written permission of SLU, is strictly prohibited.
 Saint Louis University
SCHOOL OF ENGINEERING AND ARCHITECTURE
Department of Chemical Engineering

The reaction involved the transfer of two electrons from one zinc ion to one copper ion.
Oxidation-reduction reactions are subject to influence by the same factors that have
bearing upon the rates of reactions in general: temperature, concentration, and catalyst. In
addition, they are governed largely by the inherent characteristics of the particular oxidizing
and reducing agents employed. Since atoms or ions differ in their affinity for electrons they
will differ in their ability to take electrons from other atoms or in their ability to get ri d of their
own electrons. To be able to predict whether a reaction will actually occur upon bringing
together an oxidizing agent and a reducing agent, one must know whether the oxidizing
agent has sufficient oxidizing power to take electrons from the reducing agent. Relative
oxidizing and reducing capabilities of atoms or ions are summarized in tables known as
ELECTROMOTIVE SERIES or POTENTIAL SERIES. Potential series may either be a table of standard
oxidation potentials or a table of standard reduction potentials. In a standard oxidation
potential series, the reactants are arranged in their decreasing order of their power as
oxidizing agents. The reversible half-reactions are written in such a way that the reaction
toward the right is an oxidation, and the reaction toward the left is a reduction. A reaction
proceeds spontaneously if the half-reaction of the oxidizing agent is higher in the list than
that of the half-reaction of the reducing agent. The following will be helpful in figuring out the
most probable products of reactions between the most common oxidizing agents and
reducing agents.

Table B.1 Common Oxidizing Agents and Their Usual Products

Oxidizing Agent Product(s)
HNO3, conc NO2 + H2O
HNO2, dil NO + H2O
MnO4 (acid solution)
- Mn2+ + H2O
MnO4 (basic solution)
- MnO2
MnO4 (neutral solution)
- MnO42-
Cr2O7 Cr3+ + H2O
CrO42- Cr3+ + H2O
F2, Cl2, Br2, I2 F-, Cl-, Br-, I-
Fe3+ Fe2+
MnO2 Mn2+
KClO3, KBrO3 KCl, KBr
O2 or O3 H2O or O2-
H2O2 H2O
H2SO4, conc SO2
HClO4 Cl2
K2S2O8 SO42-
KIO4 IO32-
NaBiO3 Bi 3+
PbO2 Pb2+

2
Property of and for the exclusive use of SLU. Reproduction, storing in a retrieval system, distributing, uploading or posting online, or transmitting in any form or by any
means, electronic, mechanical, photocopying, recording, or otherwise of any part of this document, without the prior written permission of SLU, is strictly prohibited.
 Saint Louis University
SCHOOL OF ENGINEERING AND ARCHITECTURE
Department of Chemical Engineering

Table B.2 Common Reducing Agents and Their Usual Products

Reducing Agent Product(s)
Metal Metallic ions (cations)
H2S S or possibly SO2 or SO42-
S SO2 or SO42-
HCl, HBr, HI Free halogen
Fe2+ Fe3+
Sn2+ Sn4+
C2O4 2- CO2 + H2O
H2 H2O or H+
CO CO2
SO2, SO3, HSO3 H2SO4 or SO42-
Na2S2O4 (acid solution) H2SO3
Na2S2O4 (basic solution) SO32-

Explore

Balancing Redox Reactions

EXERCISE B
Watch video on Balancing Redox Reaction:
Link: https://www.youtube.com/watch?v=v5sDNmYCaqo

 Note: If you cannot access the video, review your lecture notes on Balancing Redox
Reactions

3
Property of and for the exclusive use of SLU. Reproduction, storing in a retrieval system, distributing, uploading or posting online, or transmitting in any form or by any
means, electronic, mechanical, photocopying, recording, or otherwise of any part of this document, without the prior written permission of SLU, is strictly prohibited.
 Saint Louis University
SCHOOL OF ENGINEERING AND ARCHITECTURE
Department of Chemical Engineering

Evaluate

Assignment #5: (Hand-in your answers in the Google Classroom)

Balancing Redox Reactions:
1. Fe + HCl → FeCl3 + H2
2. HNO3 + H2S → S + NO + H2O
3. KMnO4 + LiCl + H2SO4 → Cl2 + MnSO4 + K2SO4 + Li2SO4 + H2O
4. K2Cr2O7 + KI + H3PO4 → I2 + CrPO4 + K3PO4 + H2O
5. K2Cr2O7 + FeSO4 + H2SO4 → Cr2(SO4)3 + Fe2(SO4)3 + K2SO4 + H2O
6. MnO4 - + Fe +2 + H+ → Mn++ + Fe3+ + H2 O
7. MnO4 -
+ C2O4 -2
+ H +
→ Mn ++
+ CO2 + H2 O
8. FeCl3 + SO2 + H2O → FeCl2 + HCl + H2SO4
9. Na2S2O3 + I2 → NaI + Na2S4O6
10. Mn(NO3)2 + 5BiO2 + HNO3 → HMNO4 + Bi(NO3)3 + H2O

4
Property of and for the exclusive use of SLU. Reproduction, storing in a retrieval system, distributing, uploading or posting online, or transmitting in any form or by any
means, electronic, mechanical, photocopying, recording, or otherwise of any part of this document, without the prior written permission of SLU, is strictly prohibited.
 Saint Louis University
SCHOOL OF ENGINEERING AND ARCHITECTURE
Department of Chemical Engineering

Unit 2: Rusting of Metals

UNIT LEARNING OUTCOME

● Execute proper calculations and analysis of chemical corrosion reactions and
methods to lower rate of metals

Engage

Why is the Statue of Liberty green?

Statue of Liberty is a gift of France to the United States. It is made of copper metal. To
see the color of copper metal, you can check the metal inside the unused wire. You can
notice that the color is somewhat brownish with luster.
Why then that the Statue of Liberty is colored green (or bluish green)? It is because of
the exposure of the statue to the atmosphere (air and water). This causes a reaction that
imparts the distinct color of the statue, specifically the Copper Oxide (CuO).

Explore
Watch the videos of EXPERIMENT 3 and accomplish the Report Sheet at the end of this
section. You are advised to read the procedures below so you can follow the video. Also,
you are tasked to do the necessary observations and record the data based on what will be
showed in the video of the activity.
EXPERIMENT 3

Part I: Rusting of Steel Using the Salt Drop Technique. (First described in 1926 by U. R. Evans.
See Scully, J. C., The Fundamentals of Corrosion, 2nd Ed., Pergamon. 1975. p. 57.)
Procedure
Click on the link to watch the video:
https://drive.google.com/file/d/1S4PM7pUFQRjGMJSaMDkAXmkz9XI6wYUn/view?usp=shari
ng
1. Plain Steel

Obtain 100 mL of salt solution and add 10 drops of phenolphthalein. On a section of

mild steel, combine 4 drops of this solution and 3 drops of potassium ferri cyanide and
cover with a watchglass. Observe for at least five minutes. What changes occur?

On the same bar do as above except use ferrocyanide. Observe for at least 5 minutes.
What changes occur? Which chemical reagent (ferro or ferri) would you use to check for
rust on iron?

5
Property of and for the exclusive use of SLU. Reproduction, storing in a retrieval system, distributing, uploading or posting online, or transmitting in any form or by any
means, electronic, mechanical, photocopying, recording, or otherwise of any part of this document, without the prior written permission of SLU, is strictly prohibited.
 Saint Louis University
SCHOOL OF ENGINEERING AND ARCHITECTURE
Department of Chemical Engineering

Ions are spatially separated in this salt drop experiment because the drop is thicker in the
middle than at the edges. Electrochemical reduction reactions that produce OH − occur at
the edges due to readily available oxygen from the air. Electrochemical oxidation reactions
occur at the middle of the drop due to the lack of oxygen. See Figure 3.2.

Figure 3.2

2. Polymer Coated Steel

Using the file, place a deep scratch on one area of polymer coated steel can lid.
Place 3 drops of ferricyanide and 4 drops of salt solution on the scratch. On a second
area of the polymer coated lid, place the drops as above and cover with a watch glass.
Observe both areas of the lid for at least 5 minutes. What changes occur? See Figure 3.3.

Unscratched Iron Scratched Iron

Figure 3.3
3. Tin Coated Steel
Repeat Procedure 2 using tin plated steel can side, tin side up. Observe for at least 8
minutes. What changes occur?

4. Zinc Plated Steel

Repeat Procedure 2 using a piece of galvanized steel. Observe for 5 minutes. What
changes occur? Is iron rusting?
Observe that an intense pink color forms, indicating a reaction is taking place and
OH ions are produced. No blue is seen in the drop-indicating that the iron is not rusting.
Metals such as zinc are used because these sacrificial anodes are more willing to give
up electrons (oxidize) than the iron and thus protect the iron from oxidation.

5. 25-centavo coin, 1-peso coin

Following Procedures 2 in Part I. Use the salt drop technique on each of the two
coins with a deep scratch on each. What do you think will happen?

Part II: Galvanic Series (batteries)

6
Property of and for the exclusive use of SLU. Reproduction, storing in a retrieval system, distributing, uploading or posting online, or transmitting in any form or by any
means, electronic, mechanical, photocopying, recording, or otherwise of any part of this document, without the prior written permission of SLU, is strictly prohibited.
 Saint Louis University
SCHOOL OF ENGINEERING AND ARCHITECTURE
Department of Chemical Engineering

Procedure
Click on the link to watch the video:
https://drive.google.com/file/d/1cIcyws80DsYfbMR7t-c3zZWuYYsJndmg/view?usp=sharing

1. Voltmeter Ranking of Metals

Fill a wide mouth bottle with salt solution. Hang a copper strip over the side of the jar,
and stopper the jar. Abrade all metal strips with sandpaper. Clip one lead from a
voltmeter to the copper strip and the second lead to a metal strip into the solution
through the hole in your stopper and record the voltage on Table 3.1. Obtain two more
sets of readings from other students; average and calculate the standard deviation. Rank
your metals in ascending order of voltage.

Table 3.1: Measured Voltage vs. Copper Foil

2. Galvanic Couples of Metals

Place 2 strips of metal from Table 3.2 on each other and fold one end of the strips over
each other several times. Flip one metal out so that both metals are visible (see
figure 3.4). Place several drops of your salt solution on the junction of the 2 metals.
Observe and record which metal turns pink on Table 3.2. In using Mg, if both metals turn
pink, ignore the Mg.

Figure 3.4

7
Property of and for the exclusive use of SLU. Reproduction, storing in a retrieval system, distributing, uploading or posting online, or transmitting in any form or by any
means, electronic, mechanical, photocopying, recording, or otherwise of any part of this document, without the prior written permission of SLU, is strictly prohibited.
 Saint Louis University
SCHOOL OF ENGINEERING AND ARCHITECTURE
Department of Chemical Engineering

The metal acting as a cathode turns pink therefore the other metal must be the anode
and is corroding (rusting). How do the results in Table 3.2 compare with the voltage
ranking on Table 3.1?

Table 3.2: Galvanic Couples

8
Property of and for the exclusive use of SLU. Reproduction, storing in a retrieval system, distributing, uploading or posting online, or transmitting in any form or by any
means, electronic, mechanical, photocopying, recording, or otherwise of any part of this document, without the prior written permission of SLU, is strictly prohibited.
 Saint Louis University
SCHOOL OF ENGINEERING AND ARCHITECTURE
Department of Chemical Engineering

Experiment 03 (Accomplish this Report Sheet)

Rusting of Metals
Name: ___________________________________________

Date: __________________________Laboratory Instructor: ______________________________

REPORT SHEET

Part I:
1. Plain Steel

2. Polymer Coated Steel

3. Tin Coated Steel

4. Zinc Plated Steel

5. 25-centavo coin, 1-peso coin

9
Property of and for the exclusive use of SLU. Reproduction, storing in a retrieval system, distributing, uploading or posting online, or transmitting in any form or by any
means, electronic, mechanical, photocopying, recording, or otherwise of any part of this document, without the prior written permission of SLU, is strictly prohibited.
 Saint Louis University
SCHOOL OF ENGINEERING AND ARCHITECTURE
Department of Chemical Engineering

Part II:
1. Voltmeter Ranking of Metals
Measured Voltage vs. Copper Foil
Metal Your Other Other Average Standard Rank
Data Data 1 Data 2 Deviation

Zn

Cu

Mg

Al

Pb

Sn

2. Galvanic Couples of Metals

Galvanic Couples
Sn Pb Al Mg Cu

Zn

Cu

Mg

Al

Pb

Comparison of the results of galvanic couples and voltage ranking:

10
Property of and for the exclusive use of SLU. Reproduction, storing in a retrieval system, distributing, uploading or posting online, or transmitting in any form or by any
means, electronic, mechanical, photocopying, recording, or otherwise of any part of this document, without the prior written permission of SLU, is strictly prohibited.
 Saint Louis University
SCHOOL OF ENGINEERING AND ARCHITECTURE
Department of Chemical Engineering

Explain

Rusting of metals is a special case of metal oxidation. Iron will oxidize to form rust.
Water will cause metals to rust. This reaction can be accelerated by adding salt. In the
corrosion process, metals get oxidized. For example in mild steel (which is greater than 99%
iron) the metal corrodes according to the following: 𝐹𝑒 → 𝐹𝑒 +2 +
2𝑒 − (𝑡ℎ𝑒 𝑟𝑒𝑚𝑜𝑣𝑎𝑙 𝑜𝑓 𝑒𝑙𝑒𝑐𝑡𝑟𝑜𝑛𝑠)
These electrons are consumed by reacting with another substance (usually oxygen but it can
be H+ in acids) in reduction as in 𝑂2 + 4𝑒 − = 2𝐻2 𝑂 → 4𝑂𝐻− (𝑡ℎ𝑒 𝑔𝑎𝑖𝑛𝑖𝑛𝑔 𝑜𝑓 𝑒𝑙𝑒𝑐𝑡𝑟𝑜𝑛𝑠)
In an acid solution, the reduction is 2𝐻+ + 2𝑒 − → 𝐻2
These equations indicate that in order for metals to corrode (rust), two reactions occur; an
oxidation that converts metal to metal ions and electrons and a second reaction which
consumes those electrons by converting oxygen and water to hydroxide ions. In order for
these reactions to occur, the electrons must be transported from the place where the metal
dissolves to the place where the oxygen is consumed and an ionic current must also flow
between the sites to complete the circuit. This ionic current flows more easily through water
containing electrolytes (i.e., NaCl). This accounts for the rapid rusting of unprotected steel in
a salty environment.
The final product of iron oxidation (rust) is usually a ferric oxide (often hematite Fe 2O3).
The initial corrosion product of the anodic reaction is ferrous (Fe 2+) ion. This is subsequently
oxidized to Fe3+ by exposure to oxygen. In this experiment we are looking at the initial product
only.
In the experiment we can watch the corrosion reaction by using substances that
produce a color change when they react with the products of the iron oxidation or oxygen
reduction. Recall that phenolphthalein turns pink in the presence of hydroxide and
ferricyanide turns a deep blue in the presence of iron II++ (rust).
The corrosion process may be slowed by coating the metals with other metals or
polymers in order to protect the metal from the corrosive environment. Examples of this can
be seen in food cans which have a polymer coating and in galvanized steel where iron is
coated with zinc.
When we put two metals in direct contact, one can oxidize (rust) while the other
reduces oxygen. This reaction sets up a voltage and is the primary reaction in a battery. By
measuring this voltage, it is possible to construct a list ranking the metal's oxidation
tendencies. If metals which are far apart in oxidation tendencies are placed in contact with
each other and with an electrolyte solution, severe corrosion of one metal can occur.

11
Property of and for the exclusive use of SLU. Reproduction, storing in a retrieval system, distributing, uploading or posting online, or transmitting in any form or by any
means, electronic, mechanical, photocopying, recording, or otherwise of any part of this document, without the prior written permission of SLU, is strictly prohibited.
 Saint Louis University
SCHOOL OF ENGINEERING AND ARCHITECTURE
Department of Chemical Engineering

Evaluate

Assignment #6: (Hand-in your answers in the Google Classroom)

1. Explain your observations and conclusions from the coins experiment .

2. Why does grapefruit juice left in an open can taste metallic?
3. If nerves respond to electrical currents, why do you think putting aluminum foil
on an amalgam (gray) filled tooth hurts? Dental amalgam is a mixture of Ag,
Sn, and Hg.
4. Why do they put magnesium rods in a steel hot water heater? (Hint: Think
about galvanized steel.)
5. If pipes feeding a water fountain were made of copper with lead solder at
the junctions, which metal dissolves more readily? Explain.
6. Tarnished silver can be restored by contact with magnesium in a salt solution.
In this reaction, the tarnished silver is reduced. What is oxidized?

12
Property of and for the exclusive use of SLU. Reproduction, storing in a retrieval system, distributing, uploading or posting online, or transmitting in any form or by any
means, electronic, mechanical, photocopying, recording, or otherwise of any part of this document, without the prior written permission of SLU, is strictly prohibited.
 Saint Louis University
SCHOOL OF ENGINEERING AND ARCHITECTURE
Department of Chemical Engineering

Unit 3: Mechanical Properties of Chocolate –

How Strong is your Chocolate?
UNIT LEARNING OUTCOME
 Apply techniques and analytical measurements appropriate for the analysis and
calculations of the mechanical properties and chemistry of engineering materials

Engage
Materials we encounter
As future engineers, you cannot get away with different types of materials. These
include metals, concrete, plastics, and all other types of materials for various applications.
These materials should be suitable for the intended application, thus has to be tested for their
mechanical integrity.
Materials such as metals (aluminum, iron, copper, etc.), ceramics (silicon carbide,
porcelain) or polymers (milk jugs made of polyethylene) are tested by scientists and
engineers to reveal certain mechanical properties such as the maximum stress a material
can withstand. The stress at which a material breaks is a measure of its strength.
However, today you will be testing the strength of a delicious material you know as
CHOCOLATE!
One conventional method of mechanical testing is called a 3-point bend test, in which a
load (Mass) is applied to the center of a beam which has its edges restricted.

Figure 4.1: An ideal 3- point beam bending test

Figure 4.2: Experimental setup of a bar of chocolate

13
Property of and for the exclusive use of SLU. Reproduction, storing in a retrieval system, distributing, uploading or posting online, or transmitting in any form or by any
means, electronic, mechanical, photocopying, recording, or otherwise of any part of this document, without the prior written permission of SLU, is strictly prohibited.
 Saint Louis University
SCHOOL OF ENGINEERING AND ARCHITECTURE
Department of Chemical Engineering

Explore
Watch the video of EXPERIMENT 4 and accomplish the Report Sheet at the end of this section.
You are advised to read the procedures below so you can follow the video. Also, you are
tasked to do the necessary observations and record the data based on what will be showed
in the video of the activity.
EXPERIMENT 4
Procedure
Click on the lick to watch the video:
_(put the link here)_______________________________________________________

Note: during the actual experiment procedures, be sure to record all observations (i. e. – any
bending noted in chocolate bar, if the cup is moving around, how hard the coins are falling
into the cup, etc.)
1. Using the scissors, punch two small holes in the rim of the cup. The holes should be
opposite each other.
2. Cut a piece of string that is approximately 1.5 ft long. The string needs to be long
enough to tie to both ends of the cup, and hang approximately 4 – 6 inches below
the chocolate bar.
3. Tie one of the ends of the string to one of the holes in the cup. Tie the other end of the
string to the opposite hole.
4. Record the following dimensions (be sure to include units):
- Type of chocolate bar (milk chocolate, dark, etc):
- Length of chocolate bar
- Width of chocolate bar
- Thickness of chocolate bar
5. Place a mat on the floor to protect the chocolate when it falls.
6. Place the chocolate in between the two desks. Approximately ½ inch (or less) of the
chocolate bar should be touching each desk. Note which way the notches (or
lettering) is facing and try to remain consistent throughout the experiment.
7. Place the string and cup assembly across the middle of the chocolate bar.

8. Using the funnel, start placing the coins into the cup, one at a time. The coins should
be funneled in at a steady pace, ensuring that each coin lands in the cup before the

14
Property of and for the exclusive use of SLU. Reproduction, storing in a retrieval system, distributing, uploading or posting online, or transmitting in any form or by any
means, electronic, mechanical, photocopying, recording, or otherwise of any part of this document, without the prior written permission of SLU, is strictly prohibited.
 Saint Louis University
SCHOOL OF ENGINEERING AND ARCHITECTURE
Department of Chemical Engineering

next coin enters the cup. (2-3 coins a second is a good rate.) Try funneling the coins
in a way that they do not fall a large distance when they enter the cup.
9. Continue placing coins into the cup at the steady rate until the chocolate bar breaks.
10. Record the number of coins in the cup at the time of fracture.
11. Look at the fracture surface and write down any observations.

12. Find the mass of the cup, string, and the coins in the cup at fracture using the balance.
13. Repeat steps above for each chocolate bar to be tested.

Experiment 04 (Accomplish this Report Sheet)

15
Property of and for the exclusive use of SLU. Reproduction, storing in a retrieval system, distributing, uploading or posting online, or transmitting in any form or by any
means, electronic, mechanical, photocopying, recording, or otherwise of any part of this document, without the prior written permission of SLU, is strictly prohibited.
 Saint Louis University
SCHOOL OF ENGINEERING AND ARCHITECTURE
Department of Chemical Engineering

Mechanical Properties of Chocolate- How Strong is your Chocolate?

Name: _________________________________________________________________________

Date: __________________________Laboratory Instructor: ______________________________

REPORT SHEET
1. How was each of the chocolate bars different from each other? Describe physical characteristics
of each chocolate bar below:
a. choco bar 1

b. choco bar 2

c. choco bar 3

d. choco bar 4

2. Which choco bar broke first? Describe why you think that it broke first. How many coins did it
take to break the chocolate bar?

3. Which choco bar broke last? Describe why you think that it broke last. How many coins did it take
to break the chocolate bar?

4. Determine the flexural strength of the chocolate bars from strongest to weakest.
16
Property of and for the exclusive use of SLU. Reproduction, storing in a retrieval system, distributing, uploading or posting online, or transmitting in any form or by any
means, electronic, mechanical, photocopying, recording, or otherwise of any part of this document, without the prior written permission of SLU, is strictly prohibited.
 Saint Louis University
SCHOOL OF ENGINEERING AND ARCHITECTURE
Department of Chemical Engineering

5. By observing the chocolate bar as you added coins, were you able to predict when the choco bar
was about to break? Describe below why or why not.

6. Describe below what you think would happen in an experiment that used a choco bar twice the
thickness of the thickest choco bar used in this experiment.

7. What did the “breaks” in the choco bar look like? Do you think by examining the choco bar after
it broke that you could put it back together?

Evaluate
17
Property of and for the exclusive use of SLU. Reproduction, storing in a retrieval system, distributing, uploading or posting online, or transmitting in any form or by any
means, electronic, mechanical, photocopying, recording, or otherwise of any part of this document, without the prior written permission of SLU, is strictly prohibited.
 Saint Louis University
SCHOOL OF ENGINEERING AND ARCHITECTURE
Department of Chemical Engineering

Assignment #7: (Hand-in your answers in the Google Classroom)

1. Calculate the load (P) at which each chocolate bar broke.

Hint: use P (Newtons) = m (in Kg) * 9.81 m/s 2
2. We determined the strength of the chocolate bars by finding each bar’s
“breaking point.” Stress is the calculation of this breaking point and is defined
as force divided by area. Calculate the stress (σ) at which each chocolate
broke using the formula below. (σ is stress, w is the width of the bar (in meters),
t is the thickness of the bar (in meters), l is the length of the bar (in meters), P is
load (mass) applied (in Newtons).
1.5Pl
σ=
wt 2
3. Using the calculations you made in number 3, rank the different chocolates in
order of their flexural strength.
4. Why do the same types of chocolate fail with different number of coins in the
cup?
5. Why do different types of chocolate fail with different number of coins?
6. Would you expect the chocolate to fail at a lower or higher load if the grooves
were facing the other direction? Why?
7. What would you expect if the cross section was different (ie – the chocolate
bar is thicker)? Would you expect it to take more or less coins, and why?
8. The experiment was performed here in Saint Louis University, Baguio City.
Would you expect the same results if the experiment were done in Saint Louis
College, San Fernando City? Why or why not? (Chemistry laboratories in SLU
and SLC are not air conditioned.)

References:

18
Property of and for the exclusive use of SLU. Reproduction, storing in a retrieval system, distributing, uploading or posting online, or transmitting in any form or by any
means, electronic, mechanical, photocopying, recording, or otherwise of any part of this document, without the prior written permission of SLU, is strictly prohibited.
 Saint Louis University
SCHOOL OF ENGINEERING AND ARCHITECTURE
Department of Chemical Engineering

Journal

Chemistry in Context, 8th Edition. American Chemical Society., Mc-Graw Hill Higher Education, 2015

Textbooks

Brown, Le May and Bursten. (2010) Chemistry, The Central Science, 7th ed., USA: Prentice Hall
International
Chang, R.(2010) Chemistry, 10th ed., New York: McGraw Hill
Davis, Mackenzie L.,(2010), Water and Wastewater Engineering Design Principles and
Practice.,Professional Edition, McGraw Hill
Davis, Mackenzie L.,(2013), Cornwell, David A. Introduction to Environmental Engineering, 5th ed.,
McGraw Hill
Geankoplis, Christie J.(2010),Transport Processes and Unit Operations, 3rd edition
Manahan, Stanley E.(2013),Fundamentals of Environmental and Toxicological Chemistry:
Sustainable Science, 4th Ed., CRC Press
Marteel-Parish, Anne E., Abraham, Martin A.(2013) Green Chemistry and Engineering: A Pathway
to Sustainability. Wiley Publishing
Masterton, William L. et. al.(2018),Principles and Reactions: Chemistry for Engineering
Students,Philippine Ed.,C&E Publishing, Inc.
Petrucci, R.H. (2011) General Chemistry: Principles and Applications, 10th ed., Toronto Pearson
Canada
Silberberg, MS (2013), Principles of General Chemistry, 3rd ed., New York: Mc Graw Hill
Whitten, K.W., Raymond, E.D.,Peck, M.L., Stanley, G.G.,(2004) General Chemistry, 7th ed., USA:
Brooks/Cole
Yunus Cengel and John Cimbala. (2013)Fluid Mechanics Fundamentals and Applications.
Zumdahl, S. & Zumdahl, S. (2014), Chemistry, 9th ed., USA: Brooks/Cole.

19
Property of and for the exclusive use of SLU. Reproduction, storing in a retrieval system, distributing, uploading or posting online, or transmitting in any form or by any
means, electronic, mechanical, photocopying, recording, or otherwise of any part of this document, without the prior written permission of SLU, is strictly prohibited.