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Honors Chemistry Lab Fall
By:
Mary McHale
Honors Chemistry Lab Fall
By:
Mary McHale
Online:
< http://cnx.org/content/col10456/1.16/ >
CONNEXIONS
Rice University, Houston, Texas
This selection and arrangement of content as a collection is copyrighted by Mary McHale. It is licensed under the
Creative Commons Attribution 2.0 license (http://creativecommons.org/licenses/by/2.0/).
Collection structure revised: November 15, 2007
PDF generated: October 26, 2012
For copyright and attribution information for the modules contained in this collection, see p. 111.
Table of Contents
1 Initial Lab: Avogradro and All That . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 1
2 Stoichiometry: Laws to Moles to Molarity . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 9
3 VSEPR: Molecular Shapes and Isomerism . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 17
4 Beer's Law and Data Analysis . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 23
5 Hydrogen and Fuel Cells . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 33
6 The Best Table in the World . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 49
7 Bonding 07 . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 57
8 Solid State and Superconductors . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 67
9 Organic Reactions . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 85
10 Transition Metals . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 95
11 Physical Properties of Gases . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 105
Attributions . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 111
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Chapter 1
1
Initial Lab: Avogradro and All That
Initial Lab: Avogadro and All ThatExperiment 1
Objective
•
The purpose of this laboratory exercise is to help you familiarize yourself with the layout of the
laboratory including safety aids and the equipment that you will be using this year.
•
Then, to make an order-of-magnitude estimate of the size of a carbon atom and of the number of atoms
in a mole of carbon based on simple assumptions about the spreading of a thin lm of stearic acid on
a water surface
Grading
•
•
•
Pre-lab not required for the rst lab
Lab Report (90%)
TA points (10%)
Before coming to lab. . .. . .
•
Read the following:
·
·
·
•
•
•
Lab instructions
Background Information
Concepts of the experiment
Print out the lab instructions and report form.
You may ll out the lab survey, due at the beginning of the lab, for extra credit if you wish.
Read and sign the equipment responsibility form and the safety rules, email Ms Duval at [email protected]
2 to conrm completing this requirement by noon on August 31st
Introduction
Since chemistry is an empirical (experimental) quantitative science, most of the experiments you will
do involve measurement.
Over the two semesters, you will measure many dierent types of quantities temperature, pH, absorbance, etc. but the most common quantity you will measure will be the amount of
a substance. The amount may be measured by (1) weight or mass (grams), (2) volume (milliliters or liters),
or (3) determining the number of moles. In this experiment we will review the methods of measuring mass
and volume and the calculations whereby number of moles are determined.
Experimental Procedure
We will start in the amphitheater of DBH (above DBH 180) for demonstrations: oxygen, hydrogen and
a mixture of the two in balloons and more besides.
1
2
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1
CHAPTER 1. INITIAL LAB: AVOGRADRO AND ALL THAT
2
Mandatory Safety talk by Kathryn Cavender, Director of Environmental Health and Safety at Rice.
1. Identication of Apparatus
On your benches, there are a number of dierent pieces of common equipment. With your TA's help,
identify each and sketch - I know this may sound a trivial exercise but it is necessary so that we are all on
the same page.
1. beaker
2. erlenmeyer ask
3. graduated (measuring) cylinder
4. pipette
5. burette
6. Bunsen burner
7. test tube
8. boiling tube
9. watch glass
2. Balance Use
In these general chemistry laboratories, we only use easy-to-read electronic balances saving you a lot of
time and the TA's a lot of headaches. However, it is important that you become adept at the use of them.
Three aspects of a balance are important:
1. The on/o switch. This is either on the front of the balance or on the back.
2. The "Zero" or "Tare" button. This resets the reading to zero.
3. CLEANLINESS. Before and after using a balance, ensure that the entire assembly is spotless. Dirt on
the weighing pan can cause erroneous measurements, and chemicals inside the machine can damage it.
4. Turn the balance on.
5. After the display reads zero, place a piece of weighing paper on the pan.
6. Read and record the mass. (2)
7. With a spatula, weigh approximately 0.2 g of a solid, common salt NaCl, the excess salt is discarded,
since returning the excess salt may contaminate the rest of the salt - in this exercise, this is not a big
deal but in strict analytical procedures it is.
8. Record the mass (1). To determine how much solid you actually have, simply subtract the mass of the
weighing paper(2) from the mass of the weighing paper and solid (1). Record this mass (3).You have
just determined the mass of an "unknown amount of solid."
9. Now place another piece of weighing paper on the balance and press the Zero or Tare button then
weigh out approximately 0.2 g of the salt (4).
Thus, the zero/tare button eliminates the need for
subtraction.
3. Measuring the volume of liquids
When working with liquids, we usually describe the quantity of the liquid in terms of volume, usual
units being milliliters (mL). We use three types of glassware to measure volume (1) burette, (2) volumetric
pipette, and (3) graduated cylinder.
•
Examine each piece of equipment. Note that the sides of each are graduated for the graduated cylinder
and the burette. You can read each to the accuracy of half a division.
•
Put some water into the graduated cylinder. Bend down and examine the side of the water level. Note
that it has a "curved shape." This is due to the water clinging to the glass sides and is called the
meniscus. When reading any liquid level, use the center of the meniscus as your reference point.
Graduated cylinder
1. Look at the graduations on the side of the cylinder.
increase upwards.
Note that they go from 0 on the bottom and
Thus, to get the mass of 10 mL of a liquid from a graduated cylinder, do the
following:
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3
2. Add water up to the 10 mL line as accurately as possible.
3. Dry a small beaker and weigh it (2).
4. Pour the 10 mL of water from the cylinder into the beaker. Reweigh (1).
5. Subtract the appropriate values to get the weight of the water (3).
Pipette
1. You may nd either that 0 is at the spout end or at the top of the pipette. You should be aware of
how these graduations go when using each pipette. Thus, to get the mass of 10 mL of a liquid from a
pipette, do the following:
2. Half-ll a beaker with water.
3. Squeeze the pipette bulb and attach to the top of the pipette.
Put the spout of the pipette under
water and release the bulb. It should expand, drawing the water into the pipette, do not let the water
be drawn into the bulb.
4. When the water level is past the last graduation, remove the bulb, replace with your nger, and then
remove the pipette from the water.
5. Removal of your nger will allow liquid to leave the pipette.
Always run some liquid into a waste
container in order to leave the level at an easy-to-read mark.
6. Add 10 mL of water to a pre-weighed dry beaker (5).
7. Weigh (4).
8. Subtract to get the weight of the water (6).
Burette
1. Examine the graduations. Note that 0 is at the top.
2. Using a funnel, add about 10 mL of water. To do this, rst lower the burette so that the top is easy
to reach.
3. Run a little water from the burette into a waste container. Then turn the burette upside down and
allow the rest of the water to run into the container (you will have to open the top to equalize the
pressure).
4. You have just "rinsed your burette." This should be done every time before using a burette rst rinse
with water, then repeat the process using whatever liquid is needed in the experiment.
5. Fill the burette to any convenient level (half-way is ne). It is a good technique to "overll" and then
allow liquid to run into a waste container until you reach the appropriate level so that you ll the space
from the top to the tip of the burette.
6. Dry a beaker and weigh (8).
7. Add 10 mL of water to a pre-weighed dry beaker (7).
8. Subtract to get the weight of the water (9).
4. Estimation of Avogadro's number
Briey, as a group with your TA, you will make an approximate (order of magnitude) estimate of Avogadro's number by determining the amount of stearic acid that it takes to form a single layer (called a
monolayer) on the surface of water. By making simple assumptions about the way the stearic acid molecules
pack together to form the monolayer, we can determine its thickness, and from that thickness we can estimate
the size of a carbon atom. Knowing the size of a carbon atom, we can compute its volume; and if we know
the volume occupied by a mole of carbon (in the form of a diamond), we can divide the volume of a mole of
carbon by the volume of an atom of carbon to get an estimate of Avogadro's number.
Procedure
Special Supplies: 14 cm watch glass; cm ruler; polyethylene transfer pipets; 1-mL syringes; pure distilled
water free of surface active materials; disposable rubber gloves (for cleaning own watch glasses in 0.1 M
NaOH in 50:50 methanol/water): 13 X 100 mm test tubes with rubber stoppers to t.
Chemicals: pure hexane, 0.108 g/L stearic acid (puried grade) solution in hexane. 0.1 M NaOH in 50:50
methanol/water used for washing the watch glasses.
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CHAPTER 1. INITIAL LAB: AVOGRADRO AND ALL THAT
4
SAFETY PRECAUTIONS: Hexane is ammable! There must be no open ames in the laboratory while
hexane is being used.
WASTE COLLECTION: At the end of the experiment, unused hexane solvent and stearic acid in hexane
solution should be placed in a waste container, marked "Waste hexane/stearic acid solution in hexane."
Measurement of the volume of stearic acid solution required to cover the water surface
Your TA will do this as a group demonstration:
1. Using a transfer pipette, obtain about 3-4 mL 0.108 g/L stearic acid solution in hexane in a clean, dry
13 X 100 mm test tube. Keep the tube corked when not in use.
2. Fill the clean watch glass to brim with deionized water. One recommended way to do this is to set up
your 25 mL burette on a ring stand. Wash and drain the burette with deionized water. (the deionized
water comes from the white handled spouts at each sink)
3. In a freshly cleaned and rinsed beaker, obtain more distilled water and ll the burette.
Place your
watch glass directly under the burette (about 1 inch or less from the tip) and dispense the water until
the entire watch glass is full. You may have to rell the burette 4 or 5 times to do this. With careful
dispensing, the surface tension of the water should allow you to ll the entire watch glass with relative
ease.
4. Carefully measure the diameter of the water surface with a centimeter ruler. It should be close to 14
cm, + or - a couple of millimeters. Next, rinse and ll your 1 mL syringe with stearic acid solution,
taking care to eliminate bubbles in the solution inside the syringe.
5. Read and record the initial volume of the syringe (1 mL is always a good place to start.)
6. Then add the stearic acid solution drop by drop to the water surface. Initially, the solution will spread
across the entire surface, and it will continue to do so until a complete monolayer of stearic acid has
been formed. If your rst few drops do not spread and evaporate quickly, either your water or watch
glass is still dirty.
As this point is approached, the spreading will become slower and slower, until
nally a drop will not spread out but will instead sit on the surface of the water (looking like a little
contact lens). If this "lens" persists for at least 30 s, you can safely conclude that you have added 1
drop more than is required to form a complete monolayer.
7. Now, read and record the nal volume reading of the syringe.Takes 10 min
8. Thoroughly clean the watch glass (or get another one), and repeat the experiment. Repeat until the
results agree to within 2 or 3 drops (0.04 ml).
When you have completed all of your measurements, rinse your syringe with pure hexane, and dispose of all
the hexane-containing solutions in the waste collection bottle provided.
Calculation Of Avogadro's Number
The calculation proceeds in several steps.
•
We calculate the volume of stearic acid solution in hexane required to deliver enough stearic acid to
form a monolayer.
•
All of the hexane evaporates, leaving only the thin monolayer lm of stearic acid, so we next calculate
the actual mass of pure stearic acid in the monolayer.
•
We calculate the thickness of the stearic acid monolayer, using the known density of stearic acid and
the area of the monolayer.
•
Assuming the stearic acid molecules are stacked on end and are tightly packed, and knowing that there
are 18 carbon atoms linked together in the stearic acid molecule, calculate the diameter and volume of
a carbon atom.
•
Calculate the volume of a mole of carbon atoms in diamond; divide the molar volume of carbon (diamond) by the volume of a single carbon atom to obtain an estimate of Avogadro's number. Remember
that the units of Avogadro's number are mol-1, so you can use unit analysis to check your answer.
Initial Lab: Avogadro and All ThatReport 1
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5
Note: In preparing this report you are free to use references and consult with others. However, you may
not copy from other students' work (including your laboratory partner) or misrepresent your own data (see
honor code).
Name(Print then sign): ___________________________________________________
Lab Day: ___________________Section: ________TA__________________________
Note: In preparing this report you are free to use references and consult with others. However, you may
not copy from other students' work (including your laboratory partner) or misrepresent your own data (see
honor code).
Demonstrations:
Balloons:
1. Oxygen
1. Hydrogen
2. Mixture of Hydrogen and Oxygen with relevant equation:
H2 +
O2
→
Thermite:
Include description and relevant equation: Fe2 O3 + Al
→
Dry Ice and Magnesium:
Include description and relevant equation: MgO + C
→
1. Identication of Apparatus
1. beaker
1. erlenmeyer ask
1. graduated (measuring) cylinder
1. pipette
1. burette
1. Bunsen burner
1. test tube
2. watch glass
2. Balance Use
1. Mass of weighing paper and solid, ________ g
2. Mass of weighing paper, __________ g
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CHAPTER 1. INITIAL LAB: AVOGRADRO AND ALL THAT
6
3. Mass of solid, ___________g
4. Mass of solid on tared weighing paper ____________g
3. Measuring the volume of a liquid
1. Mass of 50 mL beaker and water, g ______________
2. Mass of 50 mL beaker, g ______________________
3. Mass of water from graduated cylinder, g___________
4. Mass of 50 mL beaker and water, g ______________
5. Mass of 50 mL beaker, g ______________________
6. Mass of water from pipette, g ____________________
7. Mass of 50 mL beaker and water, g ______________
8. Mass of 50 mL beaker, g ______________________
9. Mass of water from burette, g ____________________
From a consideration of the masses of water measured above, and given that the density of water is 1 g/mL,
decide on an order of which is the most accurate method of volume measurement measuring cylinder,
pipette, or burette with (1) being the most accurate?
(1)
(2)
(3)
How precisely could each of the apparatus used be read?
(1) measuring cylinder
(2) pipette
(3) burette
4. Estimation of Avogadro's Number
Measurement of the volume of stearic acid solution required to cover the water surface
Record the diameter of the water
Trial 1
Trial 2
___________________
___________________
___________________
___________________
___________________
___________________
surface
Record the volume of stearic acid
solution required to cover the surface
Record the concentration of the
stearic acid solution
Table 1.1
Calculation Of Avogadro's Number
a. Calculation of the thickness of a monolayer of stearic acid
Trial 1
Trial 2
continued on next page
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From your data, the volume of
___________________
___________________
___________________
___________________
___________________
___________________
___________________
___________________
___________________
___________________
stearic acid solution required to
form a monolayer was
Calculate
the
mass
of
stearic
acid contained in that volume of
stearic acid solution (the concentration in grams per liter will be
given to you)
Calculate the volume, V, of pure
stearic acid in the monolayer on
the water surface. You will need
the density of solid stearic acid,
3
which is 0.85 g/ml (or g/cm ).
Calculate the area of the monolayer (A =
πr2 ,
r is the radius of
the water surface)
Calculate
the
thickness
of
the
monolayer (t = Volume/Area)
Table 1.2
b. Estimation of the size and volume of a carbon atom
A stearic acid molecule consists
Trial 1
Trial 2
___________________
___________________
___________________
___________________
of 18 carbon atoms linked together. Assuming that the thickness, t, of a monolayer is equal
to the length of the stearic acid
molecule, calculate the size of a
carbon atom, s = t/18
Assuming that a carbon atom is
a little cube, calculate the volume
of a carbon atom, volume =
s3
Table 1.3
c. Calculation of the volume of a mole of carbon atoms
Trial 1
Trial 2
continued on next page
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CHAPTER 1. INITIAL LAB: AVOGRADRO AND ALL THAT
8
Calculate the molar volume of
______________________________
_______________
carbon (diamond) by using the
density of diamond
3
3.51g/cm
)
and the atomic mass of a mole of
carbon
Is the volume of a mole of dia-
______________________________
_______________
mond the same as the actual volume of a mole of carbon atoms?
Table 1.4
d. Calculation of the volume of a mole of carbon (diamond) volume of a single carbon atom (Avogadro's
number)
Trial 1
Calculate
Avogado's
number
Trial 2
_________________________________
_____________
(NA) from the appropriate ratio
of volumes
Calculate the average value of
_________________________________
_____________
NA from your results
Express your results as a number
23 . Are you within a power of
________________________________
_____________
10
10 of the accepted value of 6.02
23 ?
10
Table 1.5
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Chapter 2
Stoichiometry: Laws to Moles to
1
Molarity
2.1 Experiment 2: Stoichiometry: Laws to Moles to Molarity
2.1.1 Objective
•
•
To determine the mass of a product of a chemical reaction
To make a solution of assigned molarity your accuracy will be tested by your TA by titration!
2.1.2 Grading
•
•
•
1
Pre-lab (10%)
Lab Report (80%)
TA points (10%)
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CHAPTER 2. STOICHIOMETRY: LAWS TO MOLES TO MOLARITY
10
2.1.3 Before Coming to Lab..
•
•
Read the lab instructions
Complete the pre-lab, due at the beginning of the lab
2.1.4 Introduction
The word stoichiometry derives from two Greek words: stoicheion (meaning "element") and metron (meaning
"measure"). Stoichiometry deals with calculations about the masses (sometimes volumes) of reactants and
products involved in a chemical reaction. Consequently, it is a very mathematical part of chemistry.
In the rst part of this lab, sodium bicarbonate is reacted with an excess of hydrochloric acid.
NaHCO3
(s) + HCl (aq) → NaCl (aq) + CO2 (g) + H2 O
By measuring the mass of NaHCO3 and balancing the equation (above), the mass of NaCl expected to
be produced can be calculated and then checked experimentally. Then, the actual amount of NaCl produced
can be compared to the predicted amount.
This process includes molar ratios, molar masses, balancing and interpreting equations, and conversions
between grams and moles and can be summarized as follows:
1. Check that the chemical equation is correctly balanced.
2. Using the molar mass of the given substance, convert the mass given in the problem to moles.
3. Construct a molar proportion (two molar ratios set equal to each other). Use it to convert to moles of
the unknown.
4. Using the molar mass of the unknown substance, convert the moles just calculated to mass.
In the second part of this lab, since a great deal of chemistry is done with solutions, a solution will be
prepared of allocated molarity. Molarity, or more correctly molar concentration, is dened to be the number
of moles of solute divided by the number of liters of solution:
cM = nVsubstance
solution
with units of [mole/L]. However molar concentration depends on the temperature so a higher temperature
would result in an increased volume with a consequential decrease in molar concentration. This can be a
º
signicant source of error, of the same order as the error in the volume measurements of a burette, when the
temperature increases more than 5
C.
Steps to preparing a solution of a certain concentration:
1. First need to know the formula for the solute, e.g. potassium chromate:
K2 CrO4 .
2. Need the molecular weight of the solute: by adding up the atomic weights of potassium, chromium
and oxygen: 39.10, 52.00 and 16.00 in the correct ratios:
3. 2
×
39.1, 52.0 and 4
×
16.00 = 194.2g/mole.
4. Then the volume of solution, usually deionised water: e.g.
for one liter of solution use a 1000 mL
volumetric ask. So a 1M solution would require 194.2g of solid
K2 CrO4
K2 CrO4
in 1 L, 0.1M 19.42g of solid
and so on.
5. Remember to ensure that all the solute is dissolved before nally lling to the mark on the volumetric
ask. If there is any undissolved solute present in the solution, the water level will go down slightly
below the mark, since the volume occupied by the solute diers from the actual volume it contributes
to the solution once it is dissolved.
Your teaching assistant will check the accuracy of the solution that you have made by titration, which
is a method of quantitatively determining the concentration of a solution.
A standard solution (known
concentration) is slowly added from a burette to a solution of the analyte (unknown concentration your
solution) until the reaction between them is judged to be complete equivalence point).
titration, some indicator must be used to locate the equivalence point.
In colorimetric
One example is the addition of
acid to base using phenolphthalein (indicator) to turn a pink solution colorless in order to determine the
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11
concentration of unknown acids and bases. Record your TAs value of the molarity of your solution on your
report form along with your percent error.
Figure 2.1
Figure 1: Reading the Burette
When an acid is neutralized by a base, since there is stoichiometrically equal amounts of acid and base
and the pH = 7, it is possible to accurately determine the concentration of either the acid or base solution.
Since:
Moles of a substance = Concentration of solution (moles/L) x Volume (L)
We can calculate the concentration of the acid or base in the solution using:
Balance Base
Bb
[U+E09E]
×
Ca
Bb
[U+E09F] ×
×
Moles of Acid = Moles of Base
Va
=
×
Balance Acid
Ba
×
[U+E09E]
Cb
Ba
[U+E09F]
×
2.1.4.1 Titration Calculations:
Step 1:Balance the neutralization equation. Determine Balance of Acid and Base.
Step 2:Determine what information is given.
Step 3:Determine what information is required.
Step 4:Solve using the equation below.
Bb × Ca × Va = Ba × Cb × Vb
2.1.4.2 Example:
Calculate the concentration of a nitric acid solution HNO3 if a 20 ml sample of the acid required an average
volume of 55 ml of a 0.047 mol/l solution of Ba[U+E09E]OH[U+E09F]2 to reach the endpoint of the titration.
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Vb
CHAPTER 2. STOICHIOMETRY: LAWS TO MOLES TO MOLARITY
12
Step 1:
2HNO3
+
Ba[U+E09E]OH[U+E09F]2
→
Ba[U+E09E]NO3 [U+E09F]2
+
2H2 O Balance Base =
1Balance Acid = 2
Step 2:Given informationVolume Acid = 20 mlVolume Base (average) = 55 ml Concentration of Base =
0.047 mol/l
Step 3: Required informationConcentration of AcidStep 4:Solve using the equation.
Ba × Cb × Vb 1 × Ca × 20ml = 2 × 0.047mol/1 × 55ml
Bb × C a × V a =
Ca = 0.2585 mol/l ( considering signicant gures 0.26
mol/l)
2.1.5 Experimental
2.1.5.1 Materials List
sodium bicarbonate
[U+E09E]NaHCO3 [U+E09F]
3M hydrochloric acid (HCl) solution
2.1.5.2 Procedure
2.1.5.3 Part 1
1. Weigh an empty 150-mL beaker on the electronic balance. Record this value in your data table.
2. Remove the beaker from the balance and add one spoonful of sodium bicarbonate (approximately 5
g). Re-weigh and record this value.
3. Pour approximately 20 mL of 3M hydrochloric acid into a 100-mL beaker. Rest a Pasteur pipette in
the beaker.
4. Add 3 drops of acid to the NaHCO3 beaker, moving the pipette so that no drops land on each other.
The key point is to spread out the adding of acid so as to hold all splatter within the walls of the
beaker.
5. Continue to add acid slowly drop by drop. As liquid begins to build up, gently swirl the beaker. This
is done to make sure any unreacted acid reaches any unreacted sodium bicarbonate. Do not add acid
while swirling.
6. Stop adding the hydrochloric acid when all bubbling has ceased. So that the minimum amount of HCl
has reacted with all of the sodium bicarbonate. Check when all the bubbling has ceased, by swirling
the beaker and to ensure that there is no more bubbling. When all the bubbling has ceased, add one
drop more of acid and swirl.
7. Weigh the beaker and contents, record.
8. Using a microwave oven, dry to constant weigh, initially for 1 min, when there is plenty of solution,
and then 10 second intervals thereafter. Measure weight to the nearest milligram.
2.1.5.4 Materials List
100 mls volumetric ask
3M hydrochloric acid (HCl) solution
sodium bicarbonate
[U+E09E]NaHCO3 [U+E09F]
methyl orange indicator
2.1.5.5 Part 2
1. Ask you TA for your assigned molarity it will range from 0.7 M to 1.2 M.
2. First need to know the formula for the solute.
3. Need the molecular weight of the solute in g/mole.
4. The volume of solution, 100 mLs.
5. Remember to ensure that all the solute is dissolved before nally lling with deionised water to the
mark on the volumetric ask.
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13
6. Take your solution to your TA to check the molarity by titration, record value on your report form
and your percent error.
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CHAPTER 2. STOICHIOMETRY: LAWS TO MOLES TO MOLARITY
14
2.2 Pre-Lab 2: Stoichiometry
(Total 10 points)
Click here
2 to print the Pre-Lab Note: In preparing this Pre-Lab you are free to use references and
consult with others.
However, you may not copy from other students' work (including your laboratory
partner) or misrepresent your own data (see honor code).
Name(Print then sign): ___________________________________________________
Lab Day: ___________________Section: ________TA__________________________
Circle the correct answer:
1) Which one of the following is a correct expression for molarity?
A) mol solute/L solvent
B) mol solute/mL solvent
C) mmol solute/mL solution
D) mol solute/kg solvent
E)
µmol
solute/L solution
2) What is the concentration (M) of KCl in a solution made by mixing 25.0 mL of 0.100 M KCl with
50.0 mL of 0.100 M KCl?
A) 0.100
B) 0.0500
C) 0.0333
D) 0.0250
E) 125
3) How many grams of CH3 OH must be added to water to prepare 150 mL of a solution that is 2.0 M
CH3 OH?
A) 9.6
B) 4.3
× 103
× 102
C) 2.4
D) 9.6
E) 4.3
4) The concentration of species in 500 mL of a 2.104 M solution of sodium sulfate is __________
M sodium ion and __________ M sulfate ion.
A) 2.104, 1.052
B) 2.104, 2.104
C) 2.104, 4.208
D) 1.052, 1.052
E) 4.208, 2.104
5) Oxalic acid is a diprotic acid. Calculate the percent of oxalic acid
H2 C2 O4
in a solid given that a
0.7984 g sample of that solid required 37.98 mL of 0.2283 M NaOH for neutralization.
A) 48.89
B) 97.78
C) 28.59
D) 1.086
E) 22.83
6) A 31.5 mL aliquot of
H2 SO4 (aq)
of unknown concentration was titrated with 0.0134 M NaOH (aq).
It took 23.9 mL of the base to reach the endpoint of the titration. The concentration (M) of the acid was
__________.
A) 0.0102
B) 0.0051
C) 0.0204
D) 0.102
2
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15
E) 0.227
3+
7) What are the respective concentrations (M) of Fe
and
I−
aorded by dissolving 0.200 mol FeI3 in
water and diluting to 725 mL?
A) 0.276 and 0.828
B) 0.828 and 0.276
C) 0.276 and 0.276
D) 0.145 and 0.435
E) 0.145 and 0.0483
8) A 36.3 mL aliquot of 0.0529 M
H2 SO4 (aq)
is to be titrated with 0.0411 M NaOH (aq). What volume
(mL) of base will it take to reach the equivalence point?
A) 93.6
B) 46.8
C) 187
D) 1.92
E) 3.84
9) A 13.8 mL aliquot of 0.176 M
H3 PO4 (aq)
is to be titrated with 0.110 M NaOH (aq). What volume
(mL) of base will it take to reach the equivalence point?
A) 7.29
B) 22.1
C) 199
D) 66.2
E) 20.9
10) A solution is prepared by adding 1.60 g of solid NaCl to 50.0 mL of 0.100 M CaCl2 . What is the
molarity of chloride ion in the nal solution? Assume that the volume of the nal solution is 50.0 mL.
A) 0.747
B) 0.647
C) 0.132
D) 0.232
E) 0.547
2.3 Report 2: Stoichiometry
(Total 80 points)
Note: In preparing this report you are free to use references and consult with others. However, you may
not copy from other students' work (including your laboratory partner) or misrepresent your own data (see
honor code). This is only an advisory template of what needs to be include in your complete lab write-up.
Name(Print then sign): ___________________________________________________
Lab Day: ___________________Section: ________TA__________________________
2.3.1 Part 1
2.3.2 Data Table
Mass
Grams
empty 150-mL beaker
NaHCO3 in beaker
Mass of NaHCO3
Table 2.1
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CHAPTER 2. STOICHIOMETRY: LAWS TO MOLES TO MOLARITY
16
Mass
Grams
NaCl plus beaker rst weighing
NaCl plus beaker second weighing
NaCl plus beaker third weighing
Table 2.2
1) The grams of NaHCO3 you had in your beaker was ________
2) Calculate how many moles of NaHCO3 the mass is ________
3) Write the molar ratio for the NaHCO3 / NaCl ratio _______
4) Write the number of moles of NaCl you predict were produced in your experiment.
5) Calculate the mass of NaCl you predict will be produced.
6) Determine, by subtraction, the actual mass of NaCl produced in your experiment.
a) rst weighing
b) second weighing
c) third weighing
7) Calculate your percentage yield.
2.3.3 Discussion Questions
1. Compare the numerical value of the observed ratio for maximum yield to the best ratio
2.3.4 Part 2
Record your TAs value of the molarity of your solution.
Calculate your percent error from your assigned value.
Complete the equation for the titration of
NaHCO3[U+E09E]aq[U+E09F]
+ HCl[U+E09E]aq[U+E09F] →
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Chapter 3
VSEPR: Molecular Shapes and
1
Isomerism
Molecular Shapes & Isomerism
Objectives
•
•
•
•
Understand the 3-dimensional nature of molecules
Learn about Molecular Symmetry
Be able to identify the various isomers possible for one molecular formula
Be able to identify enantiomers
3.1 Grading
•
•
Quiz (10%).
Lab Report Form (90%).
Before Coming to Lab . . .
Look over the following to make sure you have a basic understanding of the topics presented.
•
•
•
•
Drawing Lewis Structures
Determining the Shapes of Molecules from their Lewis Structures
Some Basic Aspects of Bonding
Model Kits
Introduction
The shape of a molecule is extremely important in determining its physical properties and reactivity. A
multitude of shapes are possible, and in today's lab, you will be looking at several.
In Part 1, you will be exploring the various symmetry elements that can be present in molecules. The
symmetry elements you will be looking for are mirror planes, rotation axes, and inversion centers. Being
able to determine which symmetry elements are present in a molecule help in understanding its chemistry. If
there is a plane present in the molecule that has the exact same arrangement of atoms on either side of the
mirror plane (σ). It is important to note that a molecule can have more
Rotation axes are represented as Cn (n = 1, 2, 3 . . .). The subscript indicates
plane, then the molecule has a
than one mirror plane.
o
how many degrees of rotation (360 /n) are needed in order to return to the same orientation of atoms with
o
which you started. So if there is a C2 axis, the rotation would be 180 . An example of a molecule having a
C2 axis is H2 O.
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18
CHAPTER 3. VSEPR: MOLECULAR SHAPES AND ISOMERISM
Figure 3.1
The third symmetry element is an
inversion center (i).
In such molecules, starting at any position
and drawing a line through the center and an equal distance to the opposite side of the molecule, you will
end up at a position with an identical environment to the one you started from.
Figure 3.2
Part 2 of the lab introduces the concept of
enantiomers.
Enantiomers are molecules sharing the
same molecular conguration, but they are non-superimposable images of each other. This concept should
become clearer as you build the models for this part of the lab. Enantiomers share many of the same physical
properties. The property which distinguishes them is the direction in which they rotate plane-polarized light.
plane-polarized light is just light in
which all wave vibrations have been ltered out except for those in one plane ). If both enantiomers are
They will rotate the light in equal amounts but in dierent directions (
present in a 1:1 ratio, the eects of the rotation of light cancel and no net rotation is observed.
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Such a
19
mixture of isomers is known as a
racemic mixture or as a racemate. Because these isomers rotate planeoptical isomers. Compounds that form optical isomers are said to
polarized light, they are also known as
be
chiral.
The chemistry of enantiomers is of great importance in the eld of medicine. It has been discovered that
with many drugs, one enantiomer will be biologically active while the other will be inactive or even produce
undesired side eects. For this reason, it has become advantageous for pharmaceutical companies to try to
synthesize the active enantiomer exclusively.
The next part of the lab deals with isomers. Isomers are molecules having the same molecular formula,
but the atoms are arranged in a dierent manner, while still obeying the rules of bonding. There are dierent
classications for isomers. For example,
structural isomers dier from one another in the order in which
the atoms are bound to each other (connectivity is dierent). On the other hand, geometrical isomers have
the same order of atoms, but the spatial arrangement of atoms is dierent (connectivity is the same). A
common example of geometrical isomers is the
cis
and
trans
forms of double bonds:
Figure 3.3
** NOTE: Remember that molecules having single carbon-carbon bonds cannot have cis/trans isomers
because there is free rotation about single bonds.
By building the models of various molecules during this lab, you will be able to better understand
molecular symmetry and isomers. Building models is not dicult; however, the chemical principles involved
are very important and you may nd some surprises in how atoms can be t together.
Finally, in Part 4, you will be applying your knowledge of VSEPR (Valence Shell Electron Pair Repulsion)
Theory in order to determine the geometry of several dierent molecules. VSEPR theory is useful in helping
to determine how atoms will orient themselves in molecules.
Basically, the idea is that the arrangement
adopted by a molecule will be the one in which the repulsions among the various
electron domains are
minimized. The two kinds of electron domains are bonding (electron pair shared by two atoms) and nonbonding (electron density centralized on one atom) pairs of electrons.
Experimental Procedure
For Parts 1 & 2: You and your lab partner are to work with one other lab group in preparing these
models (no more than 3 - 4 students). Your TA will assign each group a certain set of molecules to make
and answer questions pertaining to those molecules. Each group will then present their answers to the class.
These models will need to be completed and answers determined within 30 minutes so that we can continue
to Parts 3 & 4 as soon as possible.
For Parts 1-4, the work should be divided among the group members. Be sure to discuss the questions
and answers among yourselves, but put your own conclusions on the Report Form.
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CHAPTER 3. VSEPR: MOLECULAR SHAPES AND ISOMERISM
20
1. Symmetry Elements
Using the Molecular Framework models, make models of the following compounds:
a. CH4
b. CH3 Cl
c. CH2 Cl2
d. CHCl3
e. CH2 ClF
f. CHBrClF
g. BF3
h. BF2 Cl
i. PH3
j. PH2 Cl
Choose a color to represent each atom. For example, make all C atoms black, all H atoms white, etc.
Once the models are created, look for
symmetry elements that may be present.
Ask yourselves the
following questions:
•
Does the molecule contain a
mirror plane (σ)?
In other words, is there a plane within the molecule
which results in one half being a mirror image of the other half ?
•
Does the molecule contain a
two-fold rotation axis (C2)?
Remember from the Introduction that the
subscript indicates the degrees of rotation necessary to reach a conguration that is indistinguishable
o
from the original one. In this case, the rotation will be 180 .
•
•
•
Does the molecule contain any higher-order rotation axes?
C3 rotation by 120o
C4 rotation by 90o
C∞ (innity rotation axis) rotation of any amount will result in an indistinguishable orientation
Does the molecule have an inversion center (i)?
·
·
·
Determine which of these symmetry elements are present in your assigned molecules.
All of the columns
of the table on the report form should be lled out. If you have any diculty determining whether such
symmetry elements are present in the molecules you are assigned, your TA can provide examples of each
symmetry element.
Extra credit points can be earned by indicating in the table how many of each symmetry element are
present for each molecule (i.e. How many mirror planes are present?).
2.Mirror Images
Using the model kits, build models which are the mirror images of the models you were assigned to build
(b, c, d, e, f, g, h, i and j) in Part 1. With the two mirror images in hand, try to place the models on top of
one another, atom for atom.
superimposable mirror images of one another.
nonsuperimposable mirror images. Nonsuperimposable
If you can do this, the model and its mirror image are
If not, the molecule and its mirror image form
mirror images are also known as
enantiomers.
For each compound, decide whether the mirror image is superimposable or nonsuperimposable. Can you
make a generalization about when to expect molecules to have nonsuperimposable mirror images?
3.Isomers
In this exercise you will build models of compounds which are structural and/or geometrical isomers of
one another.
Make the following models:
A. Structural Isomers
1. Make a model(s) of C2 H5 Cl. How many dierent structural isomers are possible?
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21
2. Make a model(s) of C3 H7 Cl. How many dierent structural isomers are possible?
3. Make a model(s) of C3 H6 Cl2 . How many dierent structural isomers are possible?
B. Geometrical Isomers
1. Make a model(s) of C2 H3 Cl. How many dierent structural and geometrical isomers are possible?
2. Make a model(s) of C2 H2 Cl2 . How many dierent structural and geometrical isomers are possible?
3. Make a model(s) of cyclobutane (C4 H8 ). HINT: cyclo = ring of C atoms
4. Now make dichlorocyclobutane (C4 H6 Cl2 ) by replacing two H atoms on cyclopropane with Cl atoms.
How many dierent structural and geometrical isomers are possible for dichlorocyclobutane? You may
wish to make a couple of cyclobutane molecules so that you can compare the structures. Do any of the
isomers have nonsuperimposable mirror images?
C. Aromatic Ring Compounds
1. Make a model of benzene, C6 H6 . Even though your model will contain alternating double and single
bonds, remember that in the real molecules of benzene all the C-C bonds are equivalent.
What
symmetry elements does benzene possess?
2. Make a model(s) of chlorobenzene, C6 H5 Cl. How many dierent structural and geometrical isomers
are possible?
3. Make a model(s) of dichlorobenzene, C6 H4 Cl2 . How many dierent structural and geometrical isomers
are possible?
4. Make a model(s) of trichlorobenzene, C6 H3 Cl3 . How many dierent structural and geometrical isomers
are possible?
4. Hypervalent Structures
Hypervalent compounds are those that have
more than an octet of electrons
around them. Such
compounds are formed commonly with the heavier main group elements such as Si, Ge, Sn, Pb, P, As, Sb,
Bi, S, Se, Te, etc. but rarely with C, N or O. Refer to the rules for Electron Domain theory in order to assign
Lewis structures to the following molecules. Describe possible isomeric forms and the bond angles between
the atoms. How many lone pairs of electrons are present on the central atom of each molecule, if any? (**
Your book will be very useful in aiding you with these structures **)
a. PF5
b. PF3 Cl2
c. SF4
d. XeF2
e. BrF3
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22
CHAPTER 3. VSEPR: MOLECULAR SHAPES AND ISOMERISM
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Chapter 4
1
Beer's Law and Data Analysis
4.1 Beer's Law and Data Analysis
4.1.1 Objectives
•
Learn or review typical data analysis procedures plotting data with excel, performing linear regression
analysis, etc.
•
Explore the concepts and applications of spectrophotometry
4.1.2 Grading
•
•
•
Pre-lab (10%)
Lab Report Form including plot (80%)
TA points + Pop Quiz (10%)
4.1.3 Before coming to lab. . .
•
•
•
Read the lab instructions
Print out the lab instructions and report form.
Complete the pre-lab, due at the beginning of the lab
4.1.4 Introduction
When describing chemical compounds, scientists rely on their chemical and physical properties. In lab, we
might observe that a metal reacts violently with water, that a reactant is liquid at room temperature, or
that a powder is yellow. Chemical and physical properties can be used qualitatively to identify a material
or to predict its behavior, or quantitatively to determine how much of that material is present in a solution.
In this lab, we will develop a scheme to determine the concentration of copper sulfate in aqueous solution
using spectrophotometry.
To start, we will consider light and its interaction with matter.
of colors, especially when they contain transition metal ions.
Chemicals exhibit a diverse range
In order for a compound to have color, it
must absorb visible light. Visible light consists of electromagnetic radiation with wavelengths ranging from
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CHAPTER 4. BEER'S LAW AND DATA ANALYSIS
24
approximately 400 nm to 700 nm, a small section of the electromagnetic radiation spectrum shown below.
Light is characterized by its frequency (
space per second, or by its wavelength (
λ),
ν ), the number of times the crest of the wave passes some point in
the distance between two successive crests. These two quantities
λν = c = 3× 108 m/s. Planck related the frequency
−34
is Planck's constant, h = 6.626 × 10
J/s.
are related by the speed of light, a fundamental constant:
of light to its energy (E) according to
E = hν ,
where h
A compound will absorb light when the radiation posesses the energy needed to move an electron from
its lowest energy (ground) state to some excited state. The particular energies of radiation that a substance
absorbs dictate the colors that it exhibits. Conversely the color of a compound can help us to determine its
electronic conguration.
White
ple
sum
(like
of
light
contains
most
the
aqueous
remaining
all
wavelengths
solutions)
colors
that
in
this
absorbs
are
visible
visible
transmitted
region.
light,
by
the
the
When
color
object
a
we
and
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transparent
perceive
strike
is
our
samthe
eyes.
25
If
an object absorbs all wavelengths of visible light, none reaches our eyes, and it appears black. If it absorbs
no visible light, it will look white or colorless. If it absorbs all but orange, the material will appear orange.
We also perceive an orange color when visible light of all colors except blue strikes our eyes. Orange and
blue are complementary colors; the removal of blue from white light makes the light look orange, and vice
versa. Thus, an object has a particular color for one of two reasons: It transmits light of only that color or
it absorbs light of the complementary color.
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CHAPTER 4. BEER'S LAW AND DATA ANALYSIS
26
Figure 4.1
Complementary colors can be determined using an artist's color wheel. The wheel shows the colors of
the visible spectrum, from red to violet. Complementary colors, such as orange and blue, appear as wedges
opposite each other on the wheel.
With our eye, we can make qualitative judgments about the color(s) of light a sample absorbs. However,
given a red solution of
[Ti (H2 O)6 ]
3+
we can not determine if it absorbs green light or if it absorbs all
colors of light but red. To quantitatively determine the amount of light absorbed by a sample as a function
of wavelength, we will measure its absorption spectrum using a UV-visible spectrophotometer.
absorption spectra of aqueous [Ti (H2 O)6
]
3+
solutions are shown below.
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Typical
27
Notice the absorption maximum is at 490 nm. Because the sample absorbs more strongly in the green and yellow
regions of the visible spectrum, it appears red-violet. Measuring the absorption spectrum of a second, more
dilute solution demonstrates that the spectrum changes as a function of the concentration of the solution.
To understand how to use the absorption spectrum as a quantitative tool for chemical analysis, read on!
Spectrophotmetric Basics
The essential components of a spectrophotometer consist of a radiation source, a wavelength selector
(monochromator), a photodetector and read-out device.
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CHAPTER 4. BEER'S LAW AND DATA ANALYSIS
28
Figure 4.2
The incident light from a tungsten (visible light source) or deuterium (UV light source) lamp is focused
by a lens and passes through an entrance slit. By passing the beam through the monochromator (either a
prism or a diraction grating) it is separated into monochromatic (i.e., one-color or single-wavelength) light.
One particular wavelength of monochromatic light is selected and allowed to pass through the exit slit into
the sample. Light transmitted through the sample is detected by a photodetector which converts the signal
to an electrical current which is measured by a galvanometer and sent to a recording device, typically a
computer.
The measurement of transmittance (T) is made by determining the ratio of the intensity of incident (
I0 )
and transmitted (I) light passing through pure solvent and sample solutions as a function of wavelength.
[Note: The percent transmittance (%T) is obtained by multiplication of T by 100.] The logarithm of the
reciprocal of the transmittance is called the absorbance (A),
A = log (1 / T)
Care must be taken when small values of transmittance are being measured as stray light from either the
room or scattering within the instrument can cause large errors in your readings!
4.1.4.1 Extracting Quantitative Information
The Beer-Lambert law relates the amount of light being absorbed to the concentration of the substance
absorbing the light and the pathlength through which the light passes:
A = εbc.
In this equation, the measured absorbance (A) is related to the molar absorptivity constant (
), the path
length (b), and the molar concentration (c) of the absorbing. The concentration is directly proportional to
absorbance.
The single largest application of the spectrophotometer is for quantitative analysis. The prerequisite for
such analysis is a known absorption spectrum of the compound under investigation. Of particular importance
is the maximum absorption (at
λmax )
[Why choose the maximum? Could the choice alter the precision of
our experiment? the accuracy?], which can be easily obtained by plotting absorbance vs. wavelength at a
xed concentration. Next, a series of solutions of known concentration are prepared and their absorbance is
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29
measured at
λmax .
Plotting absorbance vs. concentration, a calibration curve can be determined and t using
linear regression (least-squares t). An unknown concentration can be deduced by measuring absorbance at
the absorption maximum and comparing it to the standard curve. Caution: The Beer-Lambert Law is only
obeyed (the standard curve is linear) for reasonably dilute solutions. Only those points in the linear range
of the standard curve may be used for accurate concentration determination.
Typical results are shown for the absorbance of [Ti (H2 O)6
]
3+
measured at 490 nm.
Concentration (mg/mL)
%Transmittance
Absorbance
0
100.
0
1
50.0
0.301
2
25.0
0.602
3
12.5
0.903
Table 4.1
Figure 4.3
Over the studied range the solutions obey Beer's Law. If a solution has a measured absorbance of 0.450,
we can calculate its concentration to be 1.5 mg/mL.
4.2 Experimental Procedure
In this experiment, each lab pair will measure the absorbance of CuSO4 at six concentrations.
You will
create a calibration curve to correlate copper sulfate concentration to absorbance. This curve will be used
next week to determine the concentration of an unknown copper sulfate solution and, in turn, the percent
yield of a series of chemical reactions.
Materials CuSO4
· 5H2 O
distilled water pipette bulb 1cm cuvette 4 - 25 mL volumetric ask for your
dilutions
Note: You will be borrowing these and must collect them from your TA. Do not forget to return the ask
at the end of the lab). All students will lose 3 points in that lab section if any go missing!
100 mL volumetric ask for the parent solution (in your drawer)10 mL volumetric pipette or 10 mL
graduated cylinder
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CHAPTER 4. BEER'S LAW AND DATA ANALYSIS
30
1. Measure out an appropriate mass of CuSO4
· 5H2 O
to get 100ml of
∼0.1M
solution and record the
mass on your report form. Show your calculation to the TA before making the solution. This is your
parent solution. Calculate the molarity using the actual mass measured and record it.
2. Do the following dilutions and calculate the concentrations for each.
Dilution (ml parent : ml total)
0:25 (DI
H2 O )
5:25
10:25
15:25
20:25
25:25 (parent solution)
Table 4.2
1. Measure the absorbance of the 6 solutions you have prepared and the unknown given to you by your
TA.
4.2.1 Analysis
Plot the concentration as a function of absorbance for your six solutions. Perform a linear regression analysis
and determine the equation of a best-t line.
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31
4.3 PreLab: Spectrophotometry and Data Analysis Beer's Law
2 for the Pre-Lab
Hopefully here
Name(Print then sign): ___________________________________________________
Lab Day: ___________________Section: ________TA__________________________
This assignment must be completed individually and turned in to your TA at the beginning of lab. You
will not be allowed to begin the lab until you have completed this assignment.
In many of the experiments that you will do throughout the duration of this course you will be asked
®
to analyze your data by making plots and calculating the best t line through your data.
commonly used to analyze data in this fashion is Microsoft Excel
One program
. The following exercise will help you
through the process used to obtain a plot and linear regression for a set of data.
Suppose you go for a 5 mile run and you tabulate the after each mile as in the following table.
Distance Traveled (miles)
Time (sec)
1
510
2
1026
3
1548
4
2077
5
2612
Table 4.3
4.3.1 Questions:
1. Plot the distance traveled in miles vs. the time in seconds.
2. Use linear regression to obtain a trendline and give the equation obtained in terms of the variables
distance traveled and time and the R-squared value. Comment on the meaning of the R-squared value
and its signicance when doing data analysis.
3. Using the equation you obtained by doing linear regression, estimate how long the 6th mile will take
you to run.
4. Assuming this linear trend persists, how far have you run if you nish in 2900 sec.
EXCEL INSTRUCTIONS:
•
In order to plot this data in excel, you should enter the data exactly as above in to column A (rows
1-6) and column B (rows 1-6).
•
To plot the data you will need to go to Insert on the tool bar and then click Chart. A Chart Wizard
will appear. Select XY(Scatter) as the Chart type and choose the sub-type that does not have any
lines connecting the points, then click next.
•
•
On Step 2, click on the series tab near the top of the screen and click Add.
You do not need to name the series unless you have multiple plots on one graph, but you can type in
a name if you wish.
•
•
To insert the X data, click on the icon at the far left of the x series box.
Select the x values by clicking on the rst one and while the left mouse button is down dragging the
mouse down to the last value.
•
When the values have been selected click the icon again and repeat for the y values. You may also
manually enter the values by separating them with a comma. Don't forget you need to remember what
units you are using when answering questions.
2
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CHAPTER 4. BEER'S LAW AND DATA ANALYSIS
32
•
•
Click next when you have x and y values entered correctly.
The next step is just entering title information and changing the appearance of your plot; click next
when nished.
•
Choose your chart location and then nish. You now need to place a linear regression line on your
plot.
•
•
•
Right click on a data point and pick add trendline.
Choose linear as the type and click the options tab.
Check the boxes to get the equation and R-squared value and click ok.
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Chapter 5
1
Hydrogen and Fuel Cells
5.1 Hydrogen and Fuel Cells Experiment
5.1.1 Objective
•
Build a fuel cell in order to appreciate practically a range of important chemical and physical principles,
such as galvanic cells, energy conversion, energy quality, combustion reactions, water electrolysis and
bio-fuels.
•
Crituique the design in order to improve the eciency of the fuel cell and to accomplish it practically
5.1.2 Grading
•
•
•
1
Pre-lab (10%)
Lab Report (80%)
TA points (10%)
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CHAPTER 5. HYDROGEN AND FUEL CELLS
34
5.1.3 Before Coming to Lab. . .
•
Read the lab instructions
5.1.4 Introduction
As the world's reserves of fossil fuels are diminishing and our awareness of environmental protection is
increasing, we strive to develop alternative ways of energy production. Thus in many countries research into
construction of stable and ecient fuel cells has been given high priority.
Indeed, President Bush in his
January 28th 2003 State of the Union address, proposed a $1.2 billion fuel-cell research and development
program.
Fuel cells are used for direction conversion of the energy of combustion reactions to electrical energy. A
possible fuel is hydrogen, which can be produced from water in electrolysis plants driven by solar cells or
windmills. A future interesting fuel source for operation fuel cells might be bio-fuels i.e fuels produced from
non-fossil organic material such as methane from biogas plants, alcohol produced by fermentation of sugar
or hydrolyzed starch (or, in the not so distant future, perhaps also from enzymatically hydrolyzed cellulose).
Conventional power plants turn approximately 40% of the fuel energy into electricity; we say that the
eciency of the plant is 40%. (Although, in some modern plants surplus heat is reused for district heating
thus increasing the actually eciency somewhat).
However, with fuel cells the eciency of chemical-to-
electric energy conversion is unsurpassed, namely about 70% (or even high in some experimental plants).
U.S. energy dependence is higher today than it was during the oil shock of the 1970's, and oil imports
are project to increase. Passenger vehicles alone consume 6 million barrels of oil every day, equivalent 85%
of oil imports.
•
•
If just 20% of cars used fuel cells, we could cut oil imports by 1.5 million barrels a day.
If every new vehicle sold in the U.S. next year was equipped with a 60kW fuel cell, we would double
the amount of the country's available electricity supply.
•
10,000 fuel cell vehicles running on non-petroleum feul would reduce oil consumption by 6.98 million
gallons per year.
Fuel cells could dramatically reduce urban air pollution, decrease oil imports, reduce the trade decit and
produce American jobs. The U.S. Department of Energy projects that if a mere 10% of automobiles nationwide were powered by fuel cells, regulated air pollutants would be cut by one million tons per year and 60
million tons of the greenhouse gas carbon dioxide would be eliminated. DOE projects that the same number
of feel cell cars would cut oil imports by 800,000 barrels a day about 13% of total imports. Since fuel cells
run on hydrogen derived from a renewable source, the fuel cell emissions will be nothing but water vapor.
5.2 The Chemistry of a Fuel Cell
A fuel cell is a galvanic cell in which electricity is generated by a combustion reaction. The fuel cell consists
of two electrodes between which electrical contact is established by means of an electrolyte. Oxygen or just
plain atmospheric air is fed continuously to the cathode and the fuel is fed continuously to the anode.
The fuel could be any of a vast number of combustible materials, e.g. methane, ethane or ethanol (all
organic fuels) hydrogen, hydrazine or sodium borohydride (inorganic fuels). With the hydrogen burning cell
as an example we can describe the chemistry of the cell by the following reactions:
Anode at which oxidation of the fuel takes place:
H2 + 2OH- -> 2H2O + 2e-
½
Cathode at which reduction of oxygen takes places
O2 + H2O + 2e- -> 2OH-
½
The next reaction for the cell:
H2 +
O2 -> H2O
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With ethanol as the fuel the matter becomes somewhat more complicated, since ethanol is oxidized in
steps to ethanal, ethanoic acid and carbon dioxide respectively. In an ideally working fuel cell we assume
that ethanal and ethanoic acid are further oxidized so that the only carbon compound of the overall process
is carbon dioxide. We have not succeed ( by simple chemical tests) to detect either ethanol or ethanoic acid
(or rather ethanoate due to the strongly basic electrolyte solution) as intermediate products in our own cells.
However we still suggest a three-step oxidation of ethanol(and at the same time admitting that the last step
is dubious):
Anode:
Step 1: CH3CH2OH + 2 OH- -> CH3CHO + 2 H2O + 2eStep 2: CH3HO + 2 OH- -> CH3COOH + H2O + 2 eStep 3: CH3COOH +8 OH- -> 2 CO2 + 6 H2O + 8 eSum: CH3CH2OH + 12OH- -> 2 CO2 + 9 H2O + 12 eCathode:
3O2 + 6 H2O + 12 e- -> 12OHOverall reaction:
CH3CH2OH + 3O2 -> 2CO2 + 3H2O
Sodium borohydride can power a cell in either a direct or indirect manner. Indirectly sodium boroydride
will decompose in water to produce NaBO2 (borax) and hydrogen
NaBH4 + 2H2O -> NaBO2 + 4H2
This hydrogen will then fuel the cell as shown above. However, sodium borohydride can directly power
a cell with higher energy yields.
Anode:
NaBH4 + 8OH- -> NaBO2 + 6H2O + 8eCathode:
2O2 + 4H2O + 8e- -> 8OHWhile sodium borohydride costs
∼$50
per kilogram, it has projected that mass production and borax
recycling could reduce that price to as low as $1 per kilogram.
5.3 Experimental
5.4 Caution!!! Plastic can burn.
To get good results, very careful measurements are required. Be sure to wear suitable eye protection.
Materials:
•
•
2X 50-60mL disposable hypodermic syringes without needles and pistons.
3X pieces of nickel net (2 cut to cover the anges of the syringe cylinders approximately 2cm X 10cm
+ 1 extra piece) The net should be a very ne mesh.
•
•
•
•
•
•
•
•
•
•
•
•
•
2X machine screws with nuts and waters (all brass)
2X 20cm pieces of insulated 1mm copper wire with
∼1.5cm
insulation removed from each end
Heating plate
Aluminum plate 4-6mm thick with 7-8mm hole drilled through center
Baking paper
Screwdriver, drill, spanner, at bit, scissors, wooden board and small saw
tape
Lab stand with clamps
600mL beakers
1.5V electric motor
Red LED
digital mulitimeter
balloons
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CHAPTER 5. HYDROGEN AND FUEL CELLS
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•
•
•
•
•
•
•
•
•
electrical leads with alligator clips
1M sodium hydroxide solution
4M nitric acid
ethanol
methanol
Palladium chloride solution (very expensive and should be recycled)
NaBH4
Oxygen gas
Hydrogen gas
Building an electrode (each group should build 2)
•
•
Cut a piece of nickel mesh to cover the ange of the syringe cylinder completely
Place an aluminum plate on a heating plate. Place the baking paper on the Al plate and the nickel
net on the paper.
•
•
Heat the plate to a temperature that will melt the plastic but not burn it.
4. Place the ange of the syringe on the nickel net on the heating plate. Press down rmly so that the
nickel net is melted onto the ange. Make sure that the net is sealed tight to the whole of the ange
surface, but take care not to melt so much plastic that the cylinder hold itself is covered with molten
plastic.
•
•
Remove the syringe and net form the eating plate and allow to cool.
At one of the sides of the ange drill a hole through the ange using the electric drill. Place a piece
of wood beneath to prevent drilling into the lab bench. (see picture) Push the machine screw through
the hole and fasten using a washer and nut. (see picture)
•
Mount a piece of insulated copper wire around the machine screw by twisting an end into a loop with a
at bit and fastening it with the nut. Tighten it so that good electrical contact is established between
the wire and the nickel net. Use tape to attach the wire to the syringe cylinder.
•
•
Cut o excess nickel net around the ange.
Clean the nickel net by immersing the electrode in 4M nitric acid for at least ve minutes. Also clean
the extra piece of nickel net in this manner. This much be carried out in the fume hood since poisonous
umes may evolve. Rinse thoroughly with water.
•
Place the nickel net of the electrode in a solution of palladium chloride for 30 minutes and then gently
rinse with water. Be sure to put the extra piece of nickel net in the palladium chloride solution as well.
The electrode is now ready. You should have something that resembles the picture.
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Figure 5.1
Figure 1: Drilling holes in ange.
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CHAPTER 5. HYDROGEN AND FUEL CELLS
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Figure 5.2
Figure 2: Wire connection assembly.
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CHAPTER 5. HYDROGEN AND FUEL CELLS
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Figure 3: Final assembled cell
Building the cell
•
•
First cut top o of one of the syringes. This will be the electrode you introduce the liquid/solid fuel.
Place your two electrodes into a 600mL beaker containing 1M NaOH solution.
The nickel meshing
should be completely submerged in solution.
•
Fill a balloon with oxygen gas (from gas cylinder) and connect using rubber hosing to the syringe that
was not cut. The oxygen may bubble slowly through the syringe.
•
•
Roll up the extra piece of nickel mesh and place into the cut syringe.
Add
∼20mg of NaBH4 to the syringe with the extra piece of nickel mesh.
If time permits you may test
other fuels later.
Figure 5.4
Figure 4: Functional cell layout.
Testing the cell
•
Measure the voltage generated by your cell by taking a digital multimeter and setting it to DC voltage.
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Connect one probe to each wire of the cell. The reading may continue to grow for a while and then
stabilize. Record this stable voltage. It should read between 0.8V-1V.
•
Measure the current your cell sources by keeping the probes connected and switching to current mode.
This reading should be between 30mA-50mA.
Powering an LED
•
One fuel cell does not generate enough voltage to power anything of interest.
Just like you would
connect 2 or 4 AA batteries in series to power a portable CD player, it is necessary to connect multiple
fuel cells to generate larger voltages.
•
Pair up with another group and connect the positive terminal of one cell to the negative terminal
of the other cell using the wires with alligator clips.
Now connect the unwired positive terminal to
the positive (longer lead) of the red LED. The unwired negative terminal should be connected to the
negative (short lead) of the red LED. At this point the LED should be lit. If you do not see any light,
you should use the multimeter to check the voltage generated by the two cells in series and verify
that it is greater that 1.5V. If you do not measure any voltage verify that you have wired everything
correctly.
Figure 5.5
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CHAPTER 5. HYDROGEN AND FUEL CELLS
42
Figure 5.6
Figure 5: Powering an LED circuit.
Powering a small motor
•
While two cells in series generate the proper voltage to operate the motor, they cannot source enough
current to run a motor longer than a few seconds. By putting cells in parallel more current can be
obtained.
•
You will need two sets of two cells in series as described in the Powering an LED section (4 groups
are needed for this part).
•
Take the positive connection from each series cell and connect to one terminal of the electric motor.
Take the negative connection from each series cell and connect to the other terminal on the electric
motor. At this point the motor shaft should begin to turn. If not, check the wiring and verify that
you are applying at least 1.5V. It is also possible that 2 parallel cells will not generate enough current.
Additional cells can be added in parallel to generate more current.
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Figure 5.7
Figure 6: Powering a motor circuit.
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CHAPTER 5. HYDROGEN AND FUEL CELLS
44
5.4.1 Pre-Lab: (Total 10 Points)
Name(Print then sign): ___________________________________________________
Lab Day: ___________________Section: ________TA__________________________
This assignment must be completed individually and turned in to your TA at the beginning of lab. You
will not be allowed to begin the lab until you have completed this assignment.
1.Fill in the blanks:
Fuel
cells
are
used
______________.
combustion reaction.
for
direction
conversion
of
the
energy
of
combustion
reactions
to
A fuel cell is a ______________ in which electricity is generated by a
A fuel cell provides a ______________ voltage that can be used to power
motors, lights or any number of electrical appliances.
2.T or F At the anode, oxidation of the fuel takes place.
3.T or F The fuel cell emissions will be nothing but water vapor.
4.T or F The eciency of fuel cells, chemical-to-electric energy conversion, is approximately 40%.
Review of series and parallel circuits:
In a series circuit, the electrons in the current have to pass through all the components, which are
arranged in a line. Consider a typical series circuit in which there are three resistors of value R1, R2, and
R3.
Figure 5.8
There are two key points about a series circuit:
•
•
The current throughout the circuit is the same.
The voltages add up to the battery voltage.
Therefore:
VT = V1 + V2 + V3
From Ohm's Law:
•
•
•
•
VT = IRT;
V1 = IR1;
V2 = IR2;
V3 = IR3
Þ IRT = IR1 + IR2 + IR3
Therefore:
RTot = R1 + R2 + R3
5.In the circuit below, the current is 100 mA.
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Figure 5.9
(a) What is the current in each resistor?
(b) What is the voltage across each resistor?
(c) What is the total resistance?
(d) What is the battery voltage?
Figure 5.10
Parallel circuits have their components in parallel branches so that an individual electron can go through
one of the branches but not the others.
The current splits into the available number of branches.
In this case, the current will split into three.
•
•
For a parallel circuit:
The voltage across each branch is the same.
The currents in each branch add up to the total current.
From this:
Itot = I1 + I2 + I3
From Ohm's Law:
I T = V ; I1 = V; I2 = V; I3 = V
RT R1 R2 R3
Þ V = V + V + V
RT R1 R2 R3
Þ 1/RTot = 1/R1 + 1/R2 + 1/R3
6.This question refers to the circuit below.
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CHAPTER 5. HYDROGEN AND FUEL CELLS
46
Figure 5.11
(a) What is the total resistance of the circuit?
(b) What is the current through each resistor?
(c) What is the total current?
Figure 5.12
For resistors in both series and parallel, follow these guidelines:
•
•
Work out the total resistance of the parallel combination.
Work out the total resistance of the circuit by adding your answer in the previous step to the values
of the series resistors.
7.What is the single resistor equivalent of this circuit?
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5.4.2 Report (80 points)
Note: In preparing this report you are free to use references and consult with others. However, you may
not copy from other students' work (including your laboratory partner) or misrepresent your own data (see
honor code).
The following tables and questions should be answered in your written report. Please put the information
in the relevant section of your report (i.e. observations and results, discussion)
What would happen if zinc screws were used instead of brass?
What is the purpose of the palladium coating on the anode?
What is the purpose of the palladium coating on the cathode?
What fuel cell worked best?
Explain, in detail, why you think that the best fuel cell worked better than the others?
Debate fuel cells.
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CHAPTER 5. HYDROGEN AND FUEL CELLS
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Chapter 6
1
The Best Table in the World
6.1 The Best Table in the World!
6.1.1 Objective
The goals of this experiment are:
•
•
to observe the reactions of several metals with cold water, hot water, acids and then other metal ions.
to prepare an activity series of the metals based on the observations from the above reactions.
6.1.2 Grading
You will be assessed on:
•
observations of the reactions of several metals with cold water, hot water, acids and then other metal
ions.
•
•
preparation of an activity series of the metals based on the observations from the above reactions.
answers to the post-lab questions.
6.1.3 Background Information
First, you are going to travel back to 1869 and marvel at how the rst periodic law and table were born
when only 63 elements had been discovered at the time. A 35 year old professor of general chemistry, Dmitri
Ivanovich Mendeleev at the University of St. Petersburg (now Lennigrad) in Russia was shuing his cards
with the properties of each element on each card trying to organize his thoughts for his soon-to-be famous
textbook on chemistry. When he realized that if the elements were arranged in the order of their atomic
weights, there was a trend in properties that repeated itself several times! His paper was delivered by his
graduate student, Nikolai Aleksandrivich Menchutkin before the Russian Chemical Society while Medeleev
was busy visiting cheesemaking cooperatives at the time!
In order to see and nd order among the elements, we must have some general acquaintance with them.
Elements are made of matter, and matter is dened as anything that has mass and occupies space. This
includes everything that you can see and a lot that you cannot. It follows that in order to distinguish between
dierent types of matter (in other words dierent elements) we have to assess their properties.
There are two types of properties: intensive and extensive. In the former case, intensive properties do
not depend on the how much of an element is present but do include state (whether a substance is a solid,
1
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CHAPTER 6. THE BEST TABLE IN THE WORLD
50
liquid or gas), color and chemical reactivity. Extensive properties depend on the quantity of matter present
- mass and volume are extensive properties.
Properties can be further categorized as either chemical or physical. A chemical change describes how
the substance may change composition, such as spontaneously by combustion or in combination with other
substances. On the other hand, physical changes are those properties that can be measured without changing
the composition of the matter. Condensation of steam to water is a physical change.
6.1.4 Introduction
What is there to know about the periodic table? Why is it important? Why does it appear in nearly every
science lecture room and labs? Is it just a portrait of an aspect of chemistry or does it serve a useful purpose?
Why is the name periodic appropriate? Why is the table arranged in such a way? What are the important
features of the table? Does it give order to the approximately 120 known elements?
6.2 Relative Reactivity of Metals and the Activity Series
A supercial glance at the Periodic Table will reveal that all known elements are listed by their chemical
symbols. An in depth glance at the Periodic Table yields information on the mass of an atom of the element
in atomic mass units (amu) for the molar mass of a mole (
symbol for each element.
6.02 × 1023 ) of atoms in grams below the chemical
Above the chemical symbol for each element, there is a second number listed,
the atomic number, which gives the number of protons (positively charged particles in the nucleus), or the
number of electrons (negatively charged outside the nucleus) for a neutral atom.
Mendeleev arranged the elements in the Periodic Table in order of increasing atomic number in horizontal
rows of such length that elements with similar properties recur periodically; that is to say, they fall directly
beneath each other in the Table.
The elements in a given vertical column are referred to as a family or
group. The physical and chemical properties of the elements in a given family change gradually as one goes
from one element in the column to the next. By observing the trends in properties, the elements can be
arranged in the order in which they appear in the Periodic Table.
6.3 Procedure
6.3.1 I. Activity Series
6.3.1.1 Part 1. Reactions of Metals with Water
CAUTION! Sodium reacts very rapidly with water to evolve hydrogen and heat. This is potentially dangerous
because of the possibility of the violent explosive reaction of
H2
(g) with
O2
(g) present in the air.
CAUTION! Sodium causes severe chemical burns when it comes into contact with the skin. Note: Metallic
sodium must be stored below the surface of an inert liquid such as kerosene to prevent oxidation by air.
1. I will demonstrate the reaction of sodium and then potassium with water. Observe the rate of evolution
of
H2
gas as I use tweezers to place a tiny pea-size piece of sodium then potassium into a 500-mL beaker
full of deionised water. Record your observation on the Report Form and write a balanced equation
for this reaction.
2.
Place 5 mL
H2 O
in each of four clean tubes and label them as follows:
A.
Mg
B.
Cu
C.
Zn
D.
Ca
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Table 6.1
1. Use sandpaper or steel wool to remove the oxide from the surfaces of Mg, Cu, and Zn.
2. Place several small pieces of Mg, Cu, and Zn in the correctly labeled test tube prepared above. Place
two or three (not more!) pieces of Ca turnings in the test tube labeled "Ca".
3. Watch for evidence of reaction by noting evolution of gas bubbles and any change in the color or size
of the metal. Record your observations and write net ionic equations for each reaction.
Note: Trapped air bubbles on the metal surfaces are not indicative of a reaction.
CAUTION:
H2
is FLAMMABLE!
CAUTION: Residual calcium should be discarded in a special container designated by your instructor.
Note: Net ionic equations must balance in mass (atoms) and in total charge on each side of the equation.
6.3.1.2 Part 2. Reactions of Metals with HCl
CAUTION: The reaction of Ca with HCl is not studied. Residual calcium should be discarded in a special
container designated by your instructor.
1. Decant the water from each test tube used in the procedure above and leave the pieces of metal that
remain unreacted in each test tube.
2. Place the test tubes in a test tube rack/holder.
3. Add 2 mL of 3 M HCl solution to each test tube.
CAUTION: Some of the test tubes may become very hot.
Leave them in the rack/holder while you are
making observations.
1. Observe relative rate of
H2
gas evolution for up to 10 minutes and record your observations on your
report form.
2. Based on the observations in the previous steps, list the elements that react in 3M HCl in order of
increasing strength as reducing agents and write net ionic equations for all reactions.
6.3.1.3 Part 3. Reactions of Metals with Other Metal Ions
1. Place a clean 1 inch-square of metal foil (sheet) of each of these metals Cu, Zn and Pb on a at surface.
2. Clean the metal surfaces by sanding them with ne sandpaper or steel wool.
3. Place one or two drops in spots of each of these solutions in a clockwise order on the metal surfaces:
+
A.
0.5 M Ag
B.
0.5 M Cu
C.
0.5 M Zn
D.
0.5 M Pb
2+
2+
2+
Table 6.2
1.
2+
NOTE: Do not test a cation of a metal on a square of the same metal such as Cu
ion and Cu metal.
2. Watch for color changes in each spot as evidence of reaction. If you are not sure whether the reaction
has occurred, rinse the plate with water. A distinct spot of a dierent color on the surface is good
evidence for the reaction.
3. Write net ionic equations for each reaction . Arrange Ag, Cu, Pb and Zn in order of their increasing
strength as reducing agents. If a metal A reacts with a cation of another metal B, metal A is a stronger
reducing agent, more reactive than metal B.
4. Rinse and dry each square of metal and return it to the correct beaker on the reagent shelf for other
students to use.
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CHAPTER 6. THE BEST TABLE IN THE WORLD
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6.3.1.4 Part 4. Flame Tests
6.3.1.5 One station set up that all sections will rotate through
Clean a spatula wire by dipping it into dilute hydrochloric acid (3M) and then holding it in a hot Bunsen
ame. Repeat this until the spatula doesn't produce any color in the ame.
When the spatula is clean, moisten it again with some of the acid and then dip it into a small amount of
the solid you are testing so that some sticks to the spatula. Place the spatula back in the ame again.
If the ame color is weak, it is often worthwhile to dip the spatula back in the acid again and put it back
into the ame as if you were cleaning it. You often get a very short but intense ash of color by doing that.
Chemicals/Materials:
1. Chloride salts of Li, Na, K, Rb, Cs, Ca, Ba, Cu, Pb, Fe (II) and Fe(III) Sr (nitrate salt).
2. Glass rods with loops of Pt wire.
3. Bunsen burner/clicker.
4. Concentrated nitric acid or hydrochloric acid.
Record your observations on your report form.
It should be noted that sodium is present as an impurity in many if not most metal salts. Because sodium
imparts an especially intense color to a ame, ashes of the sodium may be observed in nearly all solutions
tested.
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6.4 Pre-Lab 5: The Best Table in the World!
2 for the Pre-Lab
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Name(Print then sign): __________________________________________________
Lab Day: ___________________Section: ________TA________________________
This assignment must be completed individually and turned in to your TA at the beginning of lab. You
will not be allowed to begin the lab until you have completed this assignment.
1. The mass of an atom of the element in atomic mass units (amu) for the molar mass of a mole (
6.02 ×
10
23 ) of atoms in grams above or below the chemical symbol for each element?
Circle the
correct one.
2. The
second
symbol
listed
for
each
element
is
the
_______
__________,
symbol
________? Fill in the blanks.
3. The number in question 2 gives the number of
•
•
____________ or
the number of ________________ for a neutral atom. Fill in the blanks
4. The elements in a given vertical column are referred to as a _________ or __________. Fill
in the blanks.
5. The horizontal rows are called __________? Fill in the blank
6. The block of elements between groups II and III are called ___________ _________? Fill
in the blanks.
7. Elements 58 to 71 are known as ____________ or __________________? Fill in the
blanks.
8. Elements 90 to 103 are known as _________ ____________? Fill in the blanks.
9. Do elements with larger atomic numbers than 92 occur naturally? True or false? Circle the correct
one.
6.5 Report 5: The Best Table in the World!
3 for the Report Form
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Note: In preparing this report you are free to use references and consult with others. However, you may
not copy from other students' work (including your laboratory partner) or misrepresent your own data (see
honor code).
Name(Print then sign): __________________________________________________
Lab Day: ___________________Section: ________TA________________________
2
3
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CHAPTER 6. THE BEST TABLE IN THE WORLD
54
6.5.1 I. Activity Series
6.5.1.1 Part 1. Reactions of Metals with Water
Metal
Observations
Net Ionic Equations (If NoReaction Occurs, write N.R)
Na
K
Mg
Cu
Zn
Ca
Table 6.3
6.5.1.2 Part 2. Reactions with HCl
Metal
Observations
Net Ionic Equations (If No Reaction Occurs, Write N.R.)
Mg
Cu
Zn
Table 6.4
2. Based on your experimental results place Mg, Cu, Zn and Ca in order of increasing strength as reducing
agents.
6.5.1.3 Part 3. Reactions with Other Metal Ions
1. Write in the appropriate box either REACTION or NO REACTION.
Zn
Ag
Cu
Zn
2+
2+
Pb
Cu
Pb
+
Do not test
Do not test
2+
Do not test
Table 6.5
2. Write balanced equations to represent the results tabulated above.
3. Based on your experimental results, arrange Ag, Cu, Zn and Pb in order of increasing strength as
reducing agents.
+
4. Arrange Ag , Cu
5.
2+
, Zn
2+
and Pb
2+
in order of increasing strength as oxidizing agents.
Combine the results from Part 2 and Part 3.
Arrange Mg, Cu, Zn, Ca, Ag and Pb in order or
increasing strength as reducing agents.
6. Place Ni in this row, if it is found that Ni will deposit on Zn foil, but not on Pb foil when a drop of
NiSO4 is placed on both.
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55
6.5.1.4 Part 4. Flame Tests
Element
Color in ame
Li
Na
K
Rb
Cs
Ca
Sr
Ba
Cu
Pb
Table 6.6
What are the limitations of this test?
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CHAPTER 6. THE BEST TABLE IN THE WORLD
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Chapter 7
Bonding 07
1
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CHAPTER 7. BONDING 07
58
7.1 Lab 5: Bonding 07
7.2 Objective
•
•
•
To test various compounds and determine their conductivity and bonding.
To understand how electronegativity can predict bond type.
To learn the relationship between bonding and conductivity.
7.3 Grading
•
•
•
Pre-Lab (10%)
Lab Report Form (80%)
TA Points (10%)
7.4 Background Information
A chemical bond is a link between atoms that results from the mutual attraction of their nuclei for electrons.
Bonding occurs in order to lower the total potential energy of each atom or ion.
Throughout nature, changes
that decrease potential energy are favored.
The main types of bonds that we will be covering are ionic bonds, covalent bonds, and metallic bonds.
An ionic bond is the chemical bond that results from the electrostatic attraction between positive (cations)
and negative (anions) ions.
The ionic relationship is a give and take relationship.
One ion donates or
gives electrons, while the other ion receives or takes electrons.
A covalent bond is a chemical bond resulting from the sharing of electrons between two atoms.
two main types of covalent bonds.
The rst being non-polar covalent bonds.
There are
These are bonds in which the
bonding electrons are shared equally by the united atoms-with a balanced electrical charge.
Polar covalent
bonds are covalent bonds in which the united atoms have an unequal attraction for the shared electrons.
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59
Figure 7.1
The role of electrons in bonding has been well-studied.
The ability of an atom or element to attract
electrons to itself is known as the element's electronegativity. A scale was rst calculated by the Nobel laureate Linus Pauling and is commonly called the Pauling electronegativity scale. The actual electronegativity
values aren't as important as how they compare to a dierent element. In Part I of today's experiment, you
will compare electronegativity values to predict the type of bond that will exist between two elements.
In the solution state, ionic compounds dissociate to give a separation of charge. The separation of charge
allows for the ow of electrons through solution. The ow of electrons is classied as conductivity. A strong
electrolyte is a compound that when dissolved in water will completely ionize or dissociate into ions.
That
is, the compound exists in water only as individual ions, and there are no intact molecules at all.
This
solution conducts electricity well.
A weak electrolyte is a compound that when dissolved in water only
partially ionizes or dissociates into ions.
ions and intact molecules.
That is, the compound exists in water as a mixture of individual
This solution conducts electricity weakly. A nonelectrolyte is a compound that
when dissolved in water does not ionize or dissociate into ions at all.
as intact molecules.
In water, this compound exists entirely
The solution does not conduct electricity at all.
By measuring the conductivity of
a dissolved compound, we can classify it as a nonelectrolyte, weak electrolyte, or strong electrolyte and
determine its ability to dissociate into ions. There are four common compounds that you will encounter in
today's lab.
ACIDS are molecular compounds which ionize (turn into ions) in water.
always
H +.
The cation that is formed is
Therefore, in the formulas for simple acids, H is always the rst element listed. Some acids
are strong electrolytes and some acids are weak electrolytes.
because by denition an acid is a
H+
There are no acids which are nonelectrolytes
donor.
BASES can be molecular compounds or ionic compounds.
Some bases are soluble and some are not.
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The
CHAPTER 7. BONDING 07
60
−
soluble bases ionize or dissociate into ions in water, and the anion formed is always OH
have hydroxide ( OH
water.
−
) as the anion.
.
The ionic bases
If they are soluble, the ions simply separate (dissociate) in the
All of the ionic bases which are soluble are strong electrolytes.
SALTS are ionic compounds which are not acids or bases.
and the anion is not hydroxide.
In other words, the cation is not hydrogen
Some salts are soluble in water and some are not.
All of the salts which
are soluble are relatively strong electrolytes.
NONELECTROLYTES are compounds which dissolve in water but do not ionize or dissociate into
ions.
These would be molecular compounds other than the acids or bases already discussed.
7.5 Experimental Procedure
Caution:Acids and bases are corrosive and can cause burns.
7.5.1 Part I. Predicting bond type through electronegativity dierences.
Using the electronegativity table provided in the lab manual, predict the type of bond that each of the
following compounds will have by the following process:
•
•
•
Find the electronegativity for each element or ion in compound using electronegativity table provided.
Subtract the electronegativites (using absolute value).
If values are between:
4.0-1.7Ionic bond-50-100% ionic
1.7-0.3Polar Covalent bond-5-50% ionic
0.3-0.0Non-Polar Covalent-0-5% ionic
Determine the type of bonding in the following compounds: KCl, CO, CaBr2 , SiH4 , MgS.
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61
Figure 7.2
7.5.2 Part II. Weak and strong electrolytes
7.5.3 Chemicals
•
•
•
•
•
•
•
•
•
•
•
•
tap water
0.1 M hydrochloric acid, HCl
0.1 M acetic acid, HC2 H3 O2
0.1 M sulfuric acid,
H2 SO4
0.1 M sodium hydroxide, NaOH
0.1 M ammonia, NH3
0.1 M sodium acetate, NaC2 H3 O2
0.1 M sodium chloride, NaCl
0.1 M ammonium acetate, NH4 C2 H3 O2
0.1 M ammonium chloride, NH4 Cl
methanol, CH3 OH
ethanol,
C2 H5 OH
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CHAPTER 7. BONDING 07
62
•
sucrose solution,
C12 H22 O11
In today's lab, you will be using a MicroLab conductivity probe to determine how well electrons ow through
a given solution. First, you will need to calibrate the probe with a non-electrolyte (distilled water) and a
very strong electrolyte.
To quantify how well a solution conducts, we will assign numerical values to the
conductance probe. A non-conducting solution will have a conductance value of 0, a poor conducting solution
will have a reading of 0 to 1,000, and good conductors will have readings of 3,000 up.
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7.5.4 Instructions for MicroLab Conductivity Experiment
Open the MicroLab Program by clicking on the Shortcut to MicroLab.exe tab on the desktop.
On the Choose an Experiment Type Tab, enter a name for the experiment, and then double click on
the MicroLab Experiment icon
Click Add Sensor, Choose sensor = Conductivity Probe
Choose an input, click on the red box that corresponds to the port that your conductivity sensor is
connected to. Choose 20,000 microseconds
Choose a Sensor, click radial button that says Conductivity Probe. Click next.
Click Perform New Calibration
Click Add Calibration Point place the conductivity probe in the non-conductive standard solution,
while swirling wait until the value is constant and then enter 0.0 into the Actual Value box in MicroLab
and hit ok.
Again, Click Add Calibration Point place the conductivity probe in the conducting standard solution,
while swirling wait until the value is constant and then enter 1020 into the Actual Value box in MicroLab
and hit ok. Repeat for 3860 as the Actual Value.
Under Curve Fit Choices , click on First order (linear) and then Accept and Save this Calibration, when
prompted to Enter the units for this calibration, leave as is and click ok, save as your name-experiment-date.
Click nish.
In the sensor area, left click on the conductivity icon and drag it to the Y-axis over data source two,
also click and drag to column B on the spreadsheet and also click and drag to the digital display window.
When ready to obtain data, click start.
This is very important: Be sure to thoroughly since the probe with DI water between every use.
Beginning with the tap water, measure the conductance of each of the following solutions.
Using the
information provided in the lab manual, classify each solution as a non-, weak, or strong electrolyte. For
those solutions that are electrolytes, record the ions present in solution.
7.5.5 Part III. Electrolyte strength and reaction rate
7.5.6 Chemicals
•
•
•
•
•
calcium carbonate powder - shake once
1 M HCl - stopper it
1 M HC2 H3 O2
0.5 M
H2 SO4
Test tube gas collection apparatus - end at 20mL
Measure 2 g of powdered calcium carbonate ( CaCO3 ) onto a piece of weigh paper. Obtain 30 mL of 1 M
HCl in a graduated cylinder. Pour the acid into the test tube apparatus. Add the calcium carbonate to the
acid and QUICKLY stopper the tube to begin collecting gas. Record the time it takes to collect 20 mL of
gas. The acid may react very fast with the CaCO3 generating the gas very rapidly. Clean out the test tube
apparatus and repeat the experiment using 1 M HC2 H3 O2 and 0.5 M
H2 SO4 .
7.5.7 Part IV. Chemical reactions
7.5.8 Chemicals
•
•
0.01 M calcium hydroxide, Ca (OH)2
Plastic straws
Obtain
∼20
mL of saturated calcium hydroxide solution.
Make sure it is clear and colorless.
Place the
conductivity probe in the solution and begin monitoring it conductivity. With your straw, slowly exhale into
the solution. Note any observations in the solution and the conductivity.
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CHAPTER 7. BONDING 07
64
7.6 Pre-Lab 5: Bonding 07
7.7 (Total 10 Points)
2 for the Pre-Lab
Hopefully here
Name(Print then sign): ___________________________________________________
Lab Day: ___________________Section: ________TA__________________________
This assignment must be completed individually and turned in to your TA at the beginning of lab. You
will not be allowed to begin the lab until you have completed this assignment.
7.7.1 Part I. Bonding of chemicals in solution
1. Write out the formulas of the following acids:
•
phosphoric ____________________
•
perchloric ____________________
•
nitric ____________________
•
sulfuric __________________
•
hydrochloric ____________________
•
acetic ____________________
1. Write out the formulas of the following bases:
•
calcium hydroxide ____________________
•
potassium hydroxide ____________________
•
sodium hydroxide ____________________
•
ammonia ____________________
1. Write out the formulas of the following salts:
•
potassium chromate ____________________
•
potassium sulfate ____________________
•
copper(II) nitrate ____________________
•
calcium carbonate ____________________
•
potassium iodide ____________________
2
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7.8 Report 5: Bonding 07
3 for the Report Form
Hopefully here
Note: In preparing this report you are free to use references and consult with others. However, you may
not copy from other students' work (including your laboratory partner) or misrepresent your own data (see
honor code).
Name(Print then sign): ___________________________________________________
Lab Day: ___________________Section: ________TA__________________________
7.8.1 Part I. Predicting bond type through electronegativity dierences.
Chemical Formula
Electroneg (1)
Electroneg (2)
Di Electroneg
Type of bond
KCl
CO
CaBr2
SiH4
MgS
Table 7.1
7.8.2 Part II. Weak and strong electrolytes
Solution Tested
Numerical Output
Electrolyte Strength
Ions Present
0.1 M HCl
0.1 M HC2 H3 O2
0.1 M
H2 SO4
0.1 M NaOH
0.1 M NH3
0.1 M NaC2 H3 O2
0.1 M NaCl
0.1 M NH4 C2 H3 O2
0.1 M NH4 Cl
CH3 OH
C2 H5 OH
Sucrose
Tap water
Table 7.2
1. Why do we use deionized water instead of tap water when making solutions for conductivity measurements?
3
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CHAPTER 7. BONDING 07
66
7.8.3 Part III. Electrolyte strength and reaction rate
2. Time to collect 20 mL of gas using 1 M HCl _______________________. Write the reaction
of HCl with CaCO3 .
3. Time to collect 20 mL of gas using 1 M HC2 H3 O2 _______________________. Write the
reaction of HC2 H3 O2 with CaCO3 .
4. Time to collect 20 mL of gas using 0.5 M
the reaction of
H2 SO4
H2 SO4 _________________________.Write
with CaCO3 .
5. Why does it take dierent lengths of time to collect 20 mL of gas?
6. Based on the time it took to collect 20 mL of gas, rank the acids in the order of increasing strength.
7. Why did we use 0.5 M
H2 SO4
instead of 1.0 M
H2 SO4 ?
7.8.4 Part IV. Chemical reactions
8. Write the chemical reaction for calcium hydroxide with your exhaled breath.
9. Write your observations for the reaction that took place (i.e. appearance, conductivity, etc.)
10. When in separate solutions, aqueous ammonia, NH3 (aq) and acetic acid HC2 H3 O2 conduct electricity
equally well. However, when the two solutions are mixed a substantial increase in electrical conductivity is
observed. Explain.
11. Separately, ammonium sulfate and barium hydroxide solutions are very good conductors. When the
two solutions are mixed a substantial decrease in conductivity is observed. Rationalize this.
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Chapter 8
Solid State and Superconductors
1
8.1 Solid State Structures and Superconductors
8.1.1 Objectives
•
•
•
Build examples of: simple cubic, body centered cubic and face centered cubic cells.
Understand and familiarize with three-dimensionality of solid state structures.
Understand how binary ionic compounds (compounds made up of two dierent types of ions) pack in
a crystal lattice.
•
Observe
the
special
electromagnetic
characteristics
of
superconducting
materials
using
1,2,3-
superconductor YBa2 Cu3 O8− , discovered in 1986 by Dr. Paul Chu at the University of Houston.
8.1.2 Grading
Your grade will be determined according to the following
•
•
•
Pre-lab (10%)
Lab report form. (80%)
TA points (10%)
8.1.3 Before coming to lab:
•
•
Read introduction and model kits section
Complete prelab exercise
8.1.4 Introduction
From the three states of matter, the solid state is the one in which matter is highly condensed. In the solid
state, when atoms, molecules or ions pack in a regular arrangement which can be repeated "innitely" in
three dimensions, a crystal is formed. A crystalline solid, therefore, possesses long-range order; its atoms,
molecules, or ions occupy regular positions which repeat in three dimensions. On the other hand an amorphous solid does not possess any long-range order. Glass is an example of an amorphous solid. And even
though amorphous solids have very interesting properties in their own right that dier from those of crystalline materials, we will not consider their structures in this laboratory exercise.
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CHAPTER 8. SOLID STATE AND SUPERCONDUCTORS
68
The simplest example of a crystal is table salt, or as we chemists know it, sodium chloride (NaCl). A
+
crystal of sodium chloride is composed of sodium cations ( Na ) and chlorine anions ( Cl
−
) that are arranged
in a specic order and extend in three dimensions. The ions pack in a way that maximizes space and provides
the right coordination for each atom (ion). Crystals are three dimensional, and in theory, the perfect crystal
would be innite. Therefore instead of having a molecular formula, crystals have an empirical formula based
on stoichiometry. Crystalline structures are dened by a unit cell which is the smallest unit that contains
the stoichiometry and the spatial arrangement of the whole crystal. Therefore a unit cell can be seen as
the building block of a crystal.
The crystal lattice
In a crystal, the network of atoms, molecules, or ions is known as a crystal lattice or simply as a lattice.
In reality, no crystal extends innitely in three dimensions and the structure (and also properties) of the
solid will vary at the surface (boundaries) of the crystal.
However, the number of atoms located at the
surface of a crystal is very small compared to the number of atoms in the interior of the crystal, and so, to a
rst approximation, we can ignore the variations at the surface for much of our discussion of crystals. Any
location in a crystal lattice is known as a lattice point. Since the crystal lattice repeats in three dimensions,
there will be an entire set of lattice points which are identical. That means that if you were able to make
yourself small enough and stand at any such lattice point in the crystal lattice, you would not be able
to tell which lattice point of the set you were at the environment of a lattice point is identical to each
correspondent lattice point throughout the crystal.
Of course, you could move to a dierent site (a non-
correspondent lattice point) which would look dierent. This would constitute a dierent lattice point. For
example, when we examine the sodium chloride lattice later, you will notice that the environment of each
sodium ion is identical. If you were to stand at any sodium ion and look around, you would see the same
thing. If you stood at a chloride ion, you would see a dierent environment but that environment would be
the same at every chloride ion. Thus, the sodium ion locations form one set of lattice points and the chloride
ion locations form another set. However, lattice points not only exist in atom positions. We could easily
dene a set of lattice points at the midpoints between the sodium and chloride ions in the crystal lattice of
sodium chloride.
The unit cell
Since the crystal lattice is made up of a regular arrangement which repeats in three dimensions, we can
save ourselves a great deal of work by considering the simple repeating unit rather than the entire crystal
lattice. The basic repeating unit is known as the unit cell. Crystalline solids often have at, well-dened
faces that make denite angles with their neighbors and break cleanly when struck. These faces lie along
well-dened directions in the unit cell.
The unit cell is the smallest, most symmetrical repeating unit that, when translated in three dimensions,
will generate the entire crystal lattice.
It is possible to have a number of dierent choices for the unit cell. By convention, the unit cell that
reects the highest symmetry of the lattice is the one that is chosen. A unit cell may be thought of as being
like a brick which is used to build a building (a crystal). Many bricks are stacked together to create the entire
structure. Because the unit cell must translate in three dimensions, there are certain geometrical constraints
placed upon its shape. The main criterion is that the opposite faces of the unit cell must be parallel. Because
of this restriction there are only six parameters that we need to dene in order to dene the shape of the
unit cell. These include three edge lengths a, b, and c and three angles
α, β[U+F02C][U+F020]and γ .
Once
these are dened all other distances and angles in the unit cell are set. As a result of symmetry, some of
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69
these angles and edge lengths may be the same. There are only seven dierent shapes for unit cells possible.
These are given in the chart below.
Unit Cell Type
Triclinic
Restrictions on Unit Cell Param-
Highest Type of Symmetry Ele-
eters
ment Required
a is not equal to b is not equal
no symmetry is required, an in-
α[U+F020]is not equal
β[U+F020]is not equal to γ .
versioncenter may be present
to c;
Monoclinic
to
a is not equal to b is not equal
highest
to c[U+F03B][U+F020][U+F020]
lowed is aC2 axis or a mirror
symmetry
element
al-
α[U+F020]=
γ[U+F020][U+F03D][U+F020]
◦
90 [U+F020] β[U+F020]is not
plane
◦
equal to 90 .
Orthorhombic
a is not equal to b is not equal
has three mutually perpendicu-
to c[U+F03B][U+F020][U+F020]
larmirror planes and/or C2 axes
α[U+F020]=
β[U+F020]=
γ[U+F020][U+F03D][U+F020]
90
Tetragonal
a
◦
=b
is
not
equal
to
c
has one C4 axis
[U+F03B][U+F020][U+F020]
α[U+F020]=
β[U+F020]=
γ[U+F020][U+F03D] 90 ◦
Cubic
a
=b
=c
has C3 and C4 axes
[U+F03B][U+F020][U+F020]
α[U+F020]=
β[U+F020]=
γ[U+F020][U+F03D][U+F020]
◦
90 [U+F020]
Hexagonal, Trigonal
a
Rhombohedral*
=b
is
not
equal
to
c
C6
axis
(hexagonal);
[U+F03B][U+F020][U+F020]
α[U+F020]=
β[U+F020]=
◦
90 [U+F02C][U+F020][U+F020]
γ[U+F03D][U+F020] 120 ◦
(trigonal)
a
C3 axis (trigonal)
=b
=c
C3
axis
[U+F03B][U+F020][U+F020]
α[U+F020]=
β[U+F020]=
γ[U+F020]is not equal to 90 ◦
Table 8.1
*There is some discussion about whether the rhombohedral unit cell is a dierent group or is really a
subset of the trigonal/hexagonal types of unit cell.
8.1.5 Stoichiometry
You will be asked to count the number of atoms in each unit cell in order to determine the stoichiometry
(atom-to-atom ratio) or empirical formula of the compound. However, it is important to remember that solid
state structures are extended, that is, they extend out in all directions such that the atoms that lie on the
corners, faces, or edges of a unit cell will be shared with other unit cells, and therefore will only contribute a
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CHAPTER 8. SOLID STATE AND SUPERCONDUCTORS
70
fraction of that boundary atom. As you build crystal lattices in these exercises you will note that eight unit
cells come together at a corner. Thus, an atom which lies exactly at the corner of a unit cell will be shared
¼
by eight unit cells which means that only 1/8 of the atom contributes to the stoichiometry of any particular
unit cell. Likewise, if an atom is on an edge, only
two unit cells.
½
of the atom will be in a unit cell because four unit cells
share an edge. An atom on a face will only contribute
to each unit cell since the face is shared between
It is very important to understand that the stoichiometry of the atoms within the unit cell must reect
the composition of the bulk material.
8.1.5.1 Binding forces in a crystal
The forces which stabilize the crystal may be ionic (electrostatic) forces, covalent bonds, metallic bonds,
van der Waals forces, hydrogen bonds, or combination of these. The properties of the crystal will change
depending upon what types of bonding is involved in holding the atoms, molecules, or ions in the lattice.
The fundamental types of crystals based upon the types of forces that hold them together are: metallic in
which metal cations held together by a sea of electrons, ionic in which cations and anions held together by
predominantly electrostatic attractions, and network in which atoms bonded together covalently throughout
the solid (also known as covalent crystal or covalent network).
8.1.5.2 Close-packing
Close-packing of spheres is one example of an arrangement of objects that forms an extended structure.
Extended close-packing of spheres results in 74% of the available space being occupied by spheres (or atoms),
with the remainder attributed to the empty space between the spheres.
This is the highest space-lling
eciency of any sphere-packing arrangement. The nature of extended structures as well as close-packing,
which occurs in two forms called hexagonal close packing (hcp) and cubic close packing (ccp), will be
explored in this lab activity. Sixty-eight of the ninety naturally occurring elements are metallic elements.
Forty of these metals have three-dimensional submicroscopic structures that can be described in terms of
close-packing of spheres.
Another sixteen of the sixty-eight naturally occurring metallic elements can be
described in terms of a dierent type of extended structure that is not as ecient at space-lling.
This
structure occupies only 68% of the available space in the unit cell. This second largest subgroup exhibits a
sphere packing arrangement called body-centered cubic (bcc).
You should be able to calculate the % of void space using simple geometry.
8.1.5.3 Packing of more than one type of ion (binary compounds) in a crystal lattice
A very useful way to describe the extended structure of many substances, particularly ionic compounds, is
to assume that ions, which may be of dierent sizes, are spherical.
The structure then is based on some
type of sphere packing scheme exhibited by the larger ion, with the smaller ion occupying the unused space
(interstitial sites). In structures of this type, coordination number refers to the number of nearest neighbors
of opposite charge. Salts exhibiting these packing arrangements will be explored in this lab activity.
8.1.5.4 Coordination number and interstitial sites
When spherical objects of equal size are packed in some type of arrangement, the number of nearest
neighbors to any given sphere is dependent upon the eciency of space lling.
The number of nearest
neighbors is called the coordination number and abbreviated as CN. The sphere packing schemes with the
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highest space-lling eciency will have the highest CN. Coordination number will be explored in this lab
activity. A useful way to describe extended structures, is by using the unit cell which as discussed above is
the repeating three-dimensional pattern for extended structures. A unit cell has a pattern for the objects as
well as for the void spaces. The remaining unoccupied space in any sphere packing scheme is found as void
space. This void space occurs between the spheres and gives rise to so-called interstitial sites.
8.1.6 Synthesis of solid state materials
There exist many synthetic methods to make crystalline solids. Traditional solid state chemical reactions are
often slow and require high temperatures and long periods of time for reactants to form the desire compound.
They also require that reactants are mixed in the solid state by grinding two solids together. In this manner
the mixture formed is heterogeneous (i.e. not in the same phase), and high temperatures are required to
increase the mobility of the ions that are being formed into the new solid binary phase. Another approach
to get solid state binary structures is using a precursor material such as a metal carbonate, that upon
decomposition at high temperatures loses gaseous CO2 resulting in very ne particles of the corresponding
metal oxide (e.g., BaCO3(s)
→ BaO(s) + CO2(g) ).
8.1.7 X-ray crystallography
8.1.8
To determine the atomic or molecular structure of a crystal diraction of X-rays is used. It was observed
that visible light can be diracted by the use of optical grids, because these are arranged in a regular
manner. Energy sources such as X-rays have such small wavelengths that only grids the size of atoms will
be able to diract X-rays.
As mentioned before a crystal has regular molecular array, and therefore it is
possible, to use X-ray diraction to determine the location of the atoms in crystal lattice. When such an
experiment is carried out we say that we have determined the crystal structure of the substance. The study
of crystal structures is known as crystallography and it is one of the most powerful techniques used today
to characterize new compounds. You will discuss the principles behind X-ray diraction in the lecture part
of this course.
8.1.9 Superconductors
A superconductor is an element, or compound that will conduct electricity without resistance when it
is below a certain temperature. Without resistance the electrical current will ow continuously in a closed
loop as long as the material is kept below an specic temperature. Since the electrical resistance is zero,
supercurrents are generated in the material to exclude the magnetic elds from a magnet brought near it.
The currents which cancel the external eld produce magnetic poles opposite to the poles of the permanent
2 . In some countries (including USA) this
magnet, repelling them to provide the lift to levitate the magnet
magnet levitation has been used for transporation. Specically trains can take advantage of this levitation
to virtually eliminate friction between the vehicle and the tracks. A train levitated over a superconductor
can attain speeds over 300 mph!
8.2 Solid State Model Kits
In this experiment we will use the Institute for Chemical Education (ICE) Solid-State Model Kits which are
designed for creating a variety of common and important solid state structures. Please be careful with these
2
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72
materials as they are quite expensive. There is a list of kit components on the inside of the lid of each box.
Please make sure that you have all the listed pieces and that these are in their proper locations when you
nish using the kit.
The TAs will deduct points from your lab grade if the kits are not returned with all pieces present and
properly organized.
8.2.1 Use of the Solid State Model Kit:
The following instructions are abbreviated. Please consult the instruction manual found in the kits for more
details if you need assistance in building any of the structures given. Note that some of the model kits are
older than others and the manuals' and page numbers may not correspond.
There are four major part types in each model kit:
*2 o-white, thick plastic template bases with holes (one with a circle, the other a semicircle);
*cardboard templates (about 20 labeled A-T);
*metal rods (to be inserted in the holes to support the plastic spheres)
*plastic spheres in 4 sizes and colors.
The spheres can represent atoms, ions, or even molecules depending upon the kind of solid it is.
You will be given directions for the use of a specic base, template, placement of the rods, selection of
spheres, and arrangement of the spheres as you progress. The ICE model kits make use of Z-diagrams to
represent how the structure will be built up. Each type of sphere will be numbered with the z layer in which
it belongs.
As we build each structure in three-dimensional space, we will be drawing gures to represent the unit
cell structures.
Each level or layer of atoms, ions, or molecules in a unit cell can be represented by a
two-dimensional base, that is, a square, hexagon, parallelogram, etc.
To draw the Z-diagrams the bottom layer is referred to as z=0. We then proceed layer by layer up the
unit cell until we reach a layer which is identical to the z=0 layer.
This is z=1.
Since z=0 and z=1 are
identical by denition, we do not have to draw z=1, although you might want to do so as you are learning
how to work with solid state gures. The layers between top and bottom are given z designations according
to their positions in the crystal. So, for example, a unit cell with 4 layers (including z=0 and z=1) would
also have z=0.33 (1/3) and z=0.67 (2/3).
Each solid-state kit has two types of bases (one using rectangular coordinates, the other using polar
coordinates) indicated by a full circle or semicircle, or by color (yellow and green.)
You will rst build structures that involve only one type of atom, as you would nd in crystalline solids
of the elements, especially that of the metals. Then you will examine ionic compounds which are known
as binary solids. Binary solids are those composed of only two types of atoms, such as sodium chloride or
calcium uoride.
If time permits there is an extra credit exercise you can do. You may not do this extra credit exercise
until the report form has been completed nor may you receive credit for the extra credit assignment unless
you fully complete the report form.
8.2.2 Working groups and teams
You and your lab partner will constitute a group. Each group will receive one model kit and two groups
will work together as a team. Your TA will assign you the structures you have to do, and at the end each
team will discuss the structures assigned on front of the class. The number of teams and the assignments
the TA will give you will be decided based on the number of students in a particular laboratory session. The
laboratory is divided for six teams (A-F)
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8.2.3 Experimental Procedure
Every part of the experimental procedure has correspondent questions on the Report Form. Do not proceed
until ALL questions accompanying each section have been answered and recorded.
1. Demonstration of the 1,2,3-superconductor YBa2 Cu3 O8−
A pellet of the 1,2,3-superconductor YBa2 Cu3 O8− is placed on the top of an inverted paper cup. The
pellet is cooled down by carefully pouring liquid nitrogen over it until the bottom of the cup is lled up. After
approximately 10 seconds (when the bubbling stops) the pellet should reach the liquid nitrogen temperature.
Your TA will then place a very strong magnet over the pellet.
What happens to the magnet? What happens as the superconductor warms up? What is the Meissner
eect? (Write observations and answer these questions on your report form)
Warning- LIQUID NITROGEN CAN CAUSE FROST BITE! Do not directly touch anything that has
come into contact with the liquid nitrogen until it is warmed up to room temperature.
NOTE TO TA: to remove a levitating magnet, simply wait until the liquid nitrogen fully evaporates
or use another magnet to "grab" the oating magnet. Be careful not to lose or break these very tiny, yet
expensive, magnets!!!!
8.2.4
8.2.5 2. Cubic Cells
There are many types of fundamental unit cells, one of which is the cubic cell.
In turn, there are three
subclasses of the cubic cell:
a. simple or primitive cubic (P)
b. body-centered cubic (bcc, I*)
c. face-centered cubic (fcc, F)
*The I designation for body-centered cubic comes from the German word innenzentriert.
We do not have time to build models of all of the unit cells possible, so we will focus on the cubic structure
and its variations. Our investigation will include several aspects of each cell type:
•
•
•
•
A.
the number of atoms per unit cell
the eciency of the packing of atoms in the volume of each unit cell
the number of nearest neighbors (coordination number) for each type of atom
the stoichiometry (atom-to-atom ratio) of the compound
Simple Cubic Unit Cells or Primitive Cubic Unit Cells (P)
8.2.5.1 Team A
Group 1. Single Unit Cell
·
·
·
·
·
Construct a simple cubic cell using template A and its matching base.
Insert rods in the 4 circled holes in the shaded region of the template.
Build the rst layer (z = 0) by placing a colorless sphere on each rod in the shaded region.
Draw a picture of this layer as previously described.
Complete the unit cell by placing 4 colorless spheres on top of the rst layer.
This is the z=1 layer.
8.2.5.1.1 Group 2. Extended Structure
•
•
•
Construct an extended cubic cell using template A.
Insert rods in the circled holes of template A in the area enclosed by the dotted lines.
Construct a set of unit cells as described for making a single unit cell.
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Look closely at the structures generated by both groups. They are called simple (or primitive) cubic.
Considering all of the cells around it, answer the corresponding questions on the report form.
B.
Body-Centered Cubic Structure (BCC)
Team B
Group 1. Single Unit Cell
·
Construct a body-centered cubic (bcc) structure using template F.
·
Insert the rods in all 5 of the holes
in the shaded region.
·
·
·
Use the guide at left and place four colorless spheres in the rst layer (1) at the corners for z=0.
Place one colorless sphere in the second layer (2) on the center rod for z=0.5
Construct the z=1 layer.
Group 2.Extended Structure
•
•
•
•
•
Using template F, construct an extended body-centered cubic structure.
Insert rods in every hole of the template/block.
Using the guide which follows, place colorless spheres for z=0 on every rod labeled 1.
For z=0.5 place colorless spheres on each rod labeled 2.
Complete the z=1 layer and then place another two layers on top.
1. Face-Centered Cubic (FCC) Structure
Team C
Group 1. Single Unit Cell
·
Construct a single face-centered cubic cell using template C, colorless spheres and the layering as
illustrated. Only put rods and spheres on one of the squares dened by the internal lines.
Figure 8.1
Group 2. Extended Structure
·
Construct an extended face-centered cubic structure using template C (You can nd instructions on
how to do it in the manual that comes with the kit.)
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3. Close-Packing: Sphere Packing & Metallic Elements
Team D
Group 1. Construct the hexagonal close-packing unit cell (use the one requiring the C6 template)
Group 2. Construct the cubic close-packing unit cell (use the one requiring the C6 template)
Team E
Group 1. Add a 2' layer on top of the existing structure.
Group 2. Add a 2' layer on top of the existing structure.
Team F
Using only the shaded portion on the template, construct the face-centered cubic unit cell which uses the
C4 template.
Compare the structures of the face-centered cubic unit cell made on the C4 template to that made on
the C6 template.
8.2.5.2 4. Interstitial sites and coordination number (CN)
8.2.5.3 Team A
Group 1 - Construct CN 8, CN 6 and CN 4 (using the C4 template).
Group 2 - Construct CN 6, CN 4 (body diagonal) (using the C6 template).
8.2.6 5. Ionic Compounds
Now we will look at some real ionic compounds which crystallize in dierent cubic unit cells. We will use
the models to determine the stoichiometry ( atom-to-atom ratios) for a formula unit.
Team B
Cesium Chloride
·
Construct a model of cesium chloride on template A. This time use colorless spheres as layers 1 and 1'
and the green spheres for layer 2.
·
Start with the shaded area and then work your way outward to an extended structure. Consider both
simple and extended structures when answering the questions which follow.
8.2.6.1 Team C
8.2.6.2 Fluorite: Calcium uoride
·
Construct a model of uorite, which is calcium uoride, on template E.
·
·
Green spheres will be used for layers 1, 3, and 1' while colorless spheres go on layers 2 and 4.
Finish with a 1' layer on top. Build the structure by placing rods in all 13 holes in the area enclosed by
the internal line.
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CHAPTER 8. SOLID STATE AND SUPERCONDUCTORS
8.2.6.3 Team D
8.2.6.4 Lithium Nitride
8.2.6.5 · Use the L template and insert 6 rods in the parallelogram portion of the dotted lines.
8.2.6.6 · Construct the pattern shown below. Be sure to include a z=1 layer. 1 is a green
sphere while 1 and 2 are blue spheres. The 0 indicates a 4.0 mm spacer tube; the 2 is an
18.6 mm spacer.
Figure 8.2
8.2.6.7 Teams E and F
8.2.6.8 Zinc Blende and Wurtzite: Zinc Sulde
Team E. Zinc Blende: Use template D to construct the crystal pattern illustrated below. Numbers 2 and 4
are blue spheres while 1 and 3 are colorless spheres and 4
is a 16.1 mm spacer.
Team F. Wurtzite: Use template L to construct the Wurtzite lattice. Numbers 1, 3 and 1' are colorless
spheres and Numbers 2 and 4 are pink spheres.
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Figure 8.3
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78
8.3 Pre-Lab: Solid State and Superconductors
8.4 (Total 10 Points)
3 for the Pre-Lab
Hopefully here
Name(Print then sign): ___________________________________________________
Lab Day: ___________________Section: ________TA__________________________
This assignment must be completed individually and turned in to your TA at the beginning of lab. You
will not be allowed to begin the lab until you have completed this assignment.
1. List the existing crystal systems (unit cell types):
2. Which of these unit cells will we study in this laboratory exercise?
3. Which are the three subclasses of this type of unit cell?
4. Dene coordination number:
5. What is the volume of a sphere? Of a cube?
8.5 Report: Solid State and Superconductors
4 for the Report Form
Hopefully here
Note: In preparing this report you are free to use references and consult with others. However, you may
not copy from other students' work (including your laboratory partner) or misrepresent your own data (see
honor code).
Name(Print then sign): ___________________________________________________
Lab Day: ___________________Section: ________TA__________________________
Part I Demonstration and Unit cell theory
A. TA Demo of the superconductor
Describe and explain your observations (What happens with the magnet? Briey describe the Meissner
eect?)
B. The unit cell 1. A cube (see below) has _______ corners, _______ edges & _______ faces.
Figure 8.4
3
4
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2. Structure A below shows how a unit cell may be drawn where one choice of unit cell is shown in bold
lines. In Structures B, C and D below, draw the outline(s) of the simplest 2-D unit cells (two-dimensional
repeating patterns depicted by a parallelogram that encloses a portion of the structure).
If the unit cell is moved in the X,Y-plane in directions parallel to its sides and in distance increments
equal to the length of its sides, it has the property of duplicating the original structural pattern of circles as
well as spaces between circles. Can a structure have more than one type of unit cell? ________
Figure 8.5
Structure A
Structure B
Structure C
Structure D
Table 8.2
3. If the circle segments enclosed inside each of the bold-faced parallelograms shown below were cut out
and taped together, how many whole circles could be constructed for each one of the patterns:
Figure 8.6
Table 8.3
4.
Shown below is a 3-D unit cell for a structure of packed spheres.
The center of each of 8 spheres
is at a corner of the cube, and the part of each that lies in the interior of the cube is shown. If all of the
sphere segments enclosed inside the unit cell could be glued together, how many whole spheres could be
constructed?
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CHAPTER 8. SOLID STATE AND SUPERCONDUCTORS
80
number of whole spheres: ________
5. For each of the gures shown below, determine the number of corners and faces. Identify and name
each as one of the regular geometric solids.
Figure 8.7
AB
A
B
Number of corners
Number of faces
Name of the shape of this object
Table 8.4
Part II Experimental
1. Cubic Cells
8.5.1 A. Simple Cubic Unit Cells or Primitive Cubic Unit Cells (P)
a. How would you designate the simple cube stacking - aa, ab, abc, or some other?
b. If the radius of each atom in this cell is r, what is the equation that describes the volume of the cube
generated in terms of r? (Note that the face of the cube is dened by the position of the rods and does not
include the whole sphere.)
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c. Draw the z-diagram for the unit cell layers.
d. To how many dierent cells does a corner atom belong? What is the fractional contribution of a single
corner atom to a particular unit cell?
e. How many corner spheres does a single unit cell possess?
f. What is the net number of atoms in a unit cell? (Number of atoms multiplied by the fraction they
contribute)
g. Pick an interior sphere in the extended array. What is the coordination number (CN) of that atom?
In other words, how many spheres are touching it? .
h. What is the formula for the volume of a sphere with radius r?
i. Calculate the packing eciency of a simple cubic unit cell (the % volume or space occupied by atomic
material in the unit cell). Hint: Consider the net number of atoms per simple cubic unit cell (question g)
the volume of one sphere (question i), and the volume of the cube (question b).
8.5.2 B. Body-Centered Cubic (BCC) Structure
a. Draw the z diagrams for the layers.
b. Fill out the table below for a BCC unit cell
Atom type
Number
Fractional Contribution
Total Contribution
Coordination Number
Corner
Body
Table 8.5
c. What is the total number of atoms in the unit cell?
d. Look at the stacking of the layers. How are they arranged if we call the layers a, b, c, etc.?
e. If the radius of each atom in this cell is r, what is the formula for the volume of the cube generated in
terms of the radius of the atom? (See diagrams below.)
f. Calculate the packing eciency of the bcc cell: Find the volume occupied by the net number of spheres
per unit cell if the radius of each sphere is r; then calculate the volume of the cube using r of the sphere and
the Pythagoras theorem (
a2 + b2 = c2 )
to nd the diagonal of the cube.
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CHAPTER 8. SOLID STATE AND SUPERCONDUCTORS
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8.5.3 C. The Face Centered Cubic (FCC) Unit Cell
8.5.4 a. Fill out the following table for a FCC unit cell.
Atom type
Number
Fractional Contribution
Total Contribution
Coordination Number
Corner
Face
Table 8.6
b. What is the total number of atoms in the unit cell?
c.
Using a similar procedure to that applied in Part B above; calculate the packing eciency of the
face-centered cubic unit cell.
1. Close-Packing
a. Compare the hexagonal and cubic close-packed structures.
b. Locate the interior sphere in the layer of seven next to the new top layer. For this interior sphere,
determine the following:
Number of touching spheres:
hexagonal close-packed (hcp)
cubic close-packed (ccp)
on layer below
on the same layer
on layer above
TOTAL CN of the interior sphere
Table 8.7
c.
Sphere packing that has this number (write below) of adjacent and touching nearest neighbors is
referred to as close-packed. Non-close-packed structures will have lower coordination numbers.
d. Are the two unit cells the identical?
e. If they are the same, how are they related? If they are dierent, what makes them dierent?
f. Is the face-centered cubic unit cell aba or abc layering? Draw a z-diagram.
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III.Interstitial sites and coordination number (CN)
a. If the spheres are assumed to be ions, which of the spheres is most likely to be the anion and which
the cation, the colorless spheres or the colored spheres? Why?
b. Consider interstitial sites created by spheres of the same size. Rank the interstitial sites, as identied
by their coordination numbers, in order of increasing size (for example, which is biggest, the site with
coordination number 4, 6 or 8?).
c. Using basic principles of geometry and assuming that the colorless spheres are the same anion with
radius r A in all three cases, calculate in terms of rA the maximum radius, rC, of the cation that will t
inside a hole of CN 4, CN 6 and CN 8. Do this by calculating the ratio of the radius of to cation to the
radius of the anion:
rC /rA .
d. What terms are used to describe the shapes (coordination) of the interstitial sites of CN 4, CN 6 and
CN 8?
CN 4: ________________
CN 6: _______________
CN 8: ________________
8.5.5 IV.Ionic Solids
A. Cesium Chloride
1. Fill the table below for Cesium Chloride
Colorless spheres
Green spheres
Type of cubic structure
Atom represented
Table 8.8
2. Using the simplest unit cell described by the colorless spheres, how many net colorless and net green
spheres are contained within that unit cell?
3. Do the same for a unit cell bounded by green spheres as you did for the colorless spheres in question
4.
4. What is the ion-to-ion ratio of cesium to chloride in the simplest unit cell which contains both? (Does
it make sense? Do the charges agree?)
B. Calcium Fluoride
1. Draw the z diagrams for the layers (include both colorless and green spheres).
2. Fill the table below for Calcium Fluoride
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84
Colorless spheres
Green spheres
Type of cubic structure
Atom represented
Table 8.9
3. What is the formula for uorite (calcium uoride)?
C. Lithium Nitride
1. Draw the z diagrams for the atom layers which you have constructed.
2. What is the stoichiometric ratio of green to blue spheres?
3. Now consider that one sphere represents lithium and the other nitrogen. What is the formula?
4. How does this formula correspond to what might be predicted by the Periodic Table?
D. Zinc Blende and Wurtzite
Fill in the table below:
Zinc Blende
Stoichiometric ratio of colorless to pink spheres
Formula unit (one sphere represents and the other the sulde ion)
Compare to predicted from periodic table
Type of unit cell
Table 8.10
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Wurtzite
Chapter 9
1
Organic Reactions
9.1 Organic Reactions
9.1.1 Objectives
•
•
•
Synthesis of some important esters.
Oxidation of a primary alcohol rst to an aldehyde and then a carboxylic acid.
To saponify a typical vegetable oil.
9.1.2 Grading
You will be assessed on
•
•
detailed answers required in the lab report.
the correctness and thoroughness of your observations.
9.1.3 Introduction
Esters are an important class of organic compounds commonly prepared from the esterication reaction of
an organic acid with an alcohol in the presence of a strong mineral acid (usually
H2 SO4 ).
They are chiey
responsible for the pleasant aromas associated with various fruits, and as such are used in perfumes and
avorings.
Some esters also have useful physiological eects.
The best known example is the analgesic
("pain killing") and anti-pyretic ("fever reducing") drug acetylsalicylic acid, otherwise known by its trade
name aspirin.
Liniments used for topical relief of sore muscles contain the ester methyl salicylate ("oil of wintergreen"),
which is prepared from the reaction of methyl alcohol with the acid group of salicylic acid. Methyl salicylate
acts as an analgesic and is absorbed through the skin; however, methyl salicylate is also a skin irritant
(like many organic substances), which in this instance provides the benecial side eect of the sensation of
warming in the area of the skin where the liniment is applied.
Oxidation of a primary alcohol may yield either an aldehyde or a carboxylic acid, depending on the
reaction conditions.
For example, mild oxidation of ethanol produces acetaldehyde, which under more
vigorous conditions may be further oxidised to acetic acid.
The oxidation of ethanol to acetic acid is
responsible for causing wine to turn sour, producing vinegar.
A number of oxidising agents may be used. Acidied sodium dichromate (VI) solution at room temperature will oxidise primary alcohols to aldehydes and secondary alcohols to ketones. At higher temperatures
primary alcohols are oxides further to acids.
1
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CHAPTER 9. ORGANIC REACTIONS
86
Figure 9.1
2−
The dichromate solution turns from the orange color of the Cr2 O7
(aq) to the blue color of the Cr
3+
(aq). This color change is the basis for the "breathalyser test". The police can ask a motorist to exhale
through a tube containing some orange crystals. If the crystals turn blue, it shows that the breath contains
a considerable amount of ethanol vapor.
Soaps are produced by the reaction of metallic hydroxides with animal fats and vegetable oils. The major
components of these fats and oils are triglycerides. Triglycerides are esters of the trihydroxy alcohol called
glycerol and various long-chain fatty acids. Tristearin is a typical triglyceride. Upon reaction with sodium
hydroxide, the ester bonds of tristearin are broken.
The products of the reaction are the soap, sodium
stearate, and glycerol. This type of reaction is called saponication (Greek: sapon, soap) and it is depicted
below.
Figure 9.2
Soap is made commercially by heating beef tallow in large kettles with an excess of sodium hydroxide.
When sodium chloride is added to this mixture (called the "saponied" mixture), the sodium salts of the
fatty acids separate as a thick curd of crude soap. Glycerol is an important by-product of the reaction. It is
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87
recovered by evaporating the water layer. The crude soap is puried, and coloring agents and perfumes are
added to meet market demands.
9.1.3.1 EXPERIMENTAL PROCEDURE
CAUTION WEAR EYE PROTECTION!
CAUTION - Concentrated sulfuric acid will burn and stain the skin as well as damage clothing. In case
of skin or clothing contact, wash the area immediately with large amounts of water.
9.1.3.2 Synthesis of esters
1. Place approximately 2 g (or 2 mL if the substance is a liquid) of the organic acid and 2 mL of the
alcohol in a large test tube.
2. Add 5 - 7 drops of concentrated (18 M) sulfuric acid, mix the solution well with a glass stirring rod
◦
and then place the test tube in a hot water bath (largest beaker in your drawer) (∼ 80 C) for 5 - 10
minutes.
3. Remove the test tube from the hot water bath and cautiously pour the mixture into about 15 mL of
saturated sodium bicarbonate contained in a small beaker. The sodium bicarbonate will destroy any
unreacted acid.
4. Observe the aroma produced from each of the following esterication reactions. Write the structure of
the esters produced, and the balanced equations for the esterication and the acid/sodium bicarbonate
reactions:
Complete the following reactions using the procedure above and record your observations.
(1)
C7 H6 O3 + CH3 OH →
salicylic acid + methyl alcohol
(2) CH3 CH2 CH2 CH2 CH2 CH2 CH2 CH2 OH
+ CH3 COOH →
1 - octanol + glacial acetic acid
(3) CH3 CH2 CH2 CH2 CH2 OH
+ CH3 COOH →
amyl alcohol + glacial acetic acid
(4)
C2 H 5 OH + CH3 COOH →
ethanol + acetic acid
9.1.3.3 Oxidation of an alcohol with acidied potassium dichromate(VI) solution
•
Add 10 drops of dilute sulfuric acid (6M) and 5 drops of potassium dichromate(VI) solution (0.01M)
to 5 drops of ethanol. The oxidising agent is added slowly to the alcohol so that the temperature is
kept below that of the alcohol and above that of the carbonyl compound. (Carbonyl compounds are
more volatile than the corresponding alcohols). Usually the alcohol is in excess of the oxidant and the
aldehyde is distilled o to avoid further oxidation.
•
•
•
Note the color and smell cautiously (Royal Wave).
Warm the mixture and smell cautiously (Royal Wave).
Repeat the experiment using rst methanol and then propan-2-ol in place of ethanol.
Describe what happens and explain the color changes.
What conditions and techniques would favour the oxidation of ethanol to
a. ethanal rather than ethanoic acid.
b. ethanoic acid rather than ethanal?
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88
9.1.3.4 Oxidation of an alcohol with acidied potassium permanganate (VII) solution
•
Add 10 drops of dilute sulfuric acid and 5 drops of potassium permanganate (VII) solution (0.01M) to
5 drops of ethanol. Note the color and smell cautiously.
•
•
•
Warm the mixture and smell cautiously (Royal Wave).
Repeat the experiment using rst methanol and then propan-2-ol in place of ethanol.
Take the pH of your nal mixture using Universal indicator paper
Describe what happens and explain the color changes.
What is your nal product?
9.1.3.5 Saponication of a vegetable oil
CAUTION - Sodium hydroxide is a very caustic material that can cause severe skin burns. Eye burns caused
by sodium hydroxide are progressive: what at rst appears to be a minor irritation can develop into a severe
injury unless the chemical is completely ushed from the eye. If sodium hydroxide comes in contact with
the eye, ush the eye with running water continuously for at least 20 minutes. Notify your TA immediately.
If sodium hydroxide is spilled on some other parts of the body, ush the aected area with running water
continuously for at least 2-3 minutes. Notify your TA immediately.
Never handle sodium hydroxide pellets with your ngers.
Use weighing paper and a scoopula.
Solid
sodium hydroxide will absorb water from the atmosphere. It is hygroscopic. Do not leave the container of
sodium hydroxide open.
Keep ethanol and ethanol-water mixtures away from open ames.
Aqueous iron chloride will stain clothes permanently and is irritating to the skin. Avoid contact with this
material.
In this experiment, you will saponify a vegetable oil
1. Pour 5 mL (5.0 g) of vegetable oil into a 250-mL beaker.
2. Slowly dissolve 2.5 g of NaOH pellets in 15 mL of the 50% ethanol/water mixture in a 50-mL beaker.
3. Add 2-3 mL of the NaOH solution to the beaker containing the oil. Heat the mixture over a hot plate
with stirring. CAUTION: Keep your face away from the beaker and work at arm's length. Stirring
is required to prevent spattering.
Every few minutes, for the next 20 minutes, add portions of the
ethanol/water mixture while continuing to stir to prevent spattering. After about 10 more minutes of
heating and stirring, the oil should be dissolved and a homogenous solution should be obtained.
4. Add 25 mL of water to the hot solution. Using the hot grips, pour this solution into a 250 mL beaker
containing 150 mL of saturated NaCl solution. Stir this mixture gently and permit it to cool for a few
minutes.
5. Skim the soap layer o the top of the solution and place it in a 50-mL beaker.
6. Into a test tube, place a pea-sized lump of your soap. Place a similar amount of laundry detergent in a
second tube and a similar amount of laundry detergent in a second tube and a similar amount of hand
soap in a third tube. Add 10 mL of water to each tube. Stopper each tube and shake thoroughly.
7. Estimate the pH of the solution using wide-range indicator solution or wide-range test paper. Record
the results. Pour the contents of the test tubes into the sink and rinse the tubes with water.
Figure 9.3
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9.2 Pre-Lab: Introductory Organic Reactions
9.3 (Total 25 Points)
2 for the Pre-Lab
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Name(Print then sign): ___________________________________________________
Lab Day: ___________________Section: ________TA__________________________
This assignment must be completed individually and turned in to your TA at the beginning of lab. You
will not be allowed to begin the lab until you have completed this assignment.
For questions 1-4, draw the structural formulae of:
1) 2,2 - dimethylbutane
2) 3-ethyl-2,4-dimethylpentane
3) 2,3,4-trimethylhexane
4) 3-ethyl-2-methylheptane
For questions 5-8, give the names of
5) CH3 CH2 CH2 CH2 CH
6) CH3 CH
= C = CH2
7) CH3 CH
= CHCH3
8)
= CH2
(CH3 )2 C = CHCH3
For questions 9-11, give the structural formulae for:
9) hex-3-ene
10) 3-methylhex-1-ene
11) 2,5- dimethylhex-2-ene
For questions 12-14, give the names of:
12)
2
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CHAPTER 9. ORGANIC REACTIONS
90
Figure 9.4
13)
Figure 9.5
14)
Figure 9.6
For questions 15-19, give the stuctural formulae of:
15) trans-1,2-dibromoethene
16) trans-1-chloroprop-1-ene
17) cis- hex-2-ene
18) pent-1-yne
19) 3-methylbut-1-yne
For questions 20-25, name the following compounds:
20)
Figure 9.7
21)
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91
Figure 9.8
22)
Figure 9.9
23)
Figure 9.10
24)
Figure 9.11
25)
Figure 9.12
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CHAPTER 9. ORGANIC REACTIONS
92
9.4 Report: Organic Reactions
3 for the Report Form
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not copy from other students' work (including your laboratory partner) or misrepresent your own data (see
honor code).
Name(Print then sign): ___________________________________________________
Lab Day: ___________________Section: ________TA__________________________
9.4.1 Observations:
9.4.1.1 Synthesis of esters
Reagents
Product
C7 H6 O3 (salicylic
acid
)
Observations
+
CH3 OH(methyl alcohol)
CH3 CH2 CH2 CH2 CH2 CH2 CH2 CH2 OH
(1
-
octanol
+CH3 COOH
)
(glacial acetic acid)
CH3 CH2 CH2 CH2 CH2 OH(amyl
+CH3 COOH
alcohol)
(glacial
acetic acid)
C2 H5 OH
(ethanol)
+CH3 COOH
(acetic acid)
Table 9.1
9.4.2 Oxidation of an alcohol with acidied potassium dichromate(VI) solution.
Remember to describe what happens and explain the color changes.
What conditions and techniques would favour the oxidation of ethanol to
a. ethanal rather than ethanoic acid.
b. ethanoic acid rather than ethanal?
Add 10 drops of dil.
drops of
K2 Cr2 O7
H2 SO4
to 5
to the follow-
Observations:
cautiously!
color/smellSmell
Observations:
color/smell
warmingSmell cautiously!
ing alcohols
Ethanol
Methanol
continued on next page
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93
Propan-2-ol
Table 9.2
9.4.3 Oxidation of an alcohol with acidied potassium permanganate (VII) solution
Remember to describe what happens and explain the color changes.
What is your nal product?
Add 10 drops of dil.
H2 SO4
to 5
drops of KMnO4 to the following
Observations:
color/smellSmell
cautiously!
Observations:
color/smell
warmingSmell cautiously!
alcohols
Ethanol
Methanol
Propan-2-ol
Table 9.3
9.4.4
9.4.5 Saponication of a vegetable oil
Reagent
pH from indicator paper
Your soap
Laundry Detergent
Hand Soap
Table 9.4
w from the File menu, and then double-click your template.
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94
CHAPTER 9. ORGANIC REACTIONS
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Chapter 10
1
Transition Metals
10.1 Transitions Metals: Synthesis of an Inorganic Compound (transdinitrobis(ethylenediamine)cobalt(III) nitrate)
10.1.1 Objectives
•
•
•
To synthesize a transition metal complex of cobalt three, Co(III), and ethylenediamine.
To characterize the resulting metal complex spectroscopically.
To understand concept of limiting reactant.
10.1.2 Grading
Your will be determined according to the following:
•
•
•
prelab (10%)
lab report form (80%)
TA points (10%)
10.1.3 Introduction
The transition metals are the largest group (classication) of elements from the periodic table. These can
be found in nature as ores or in its elemental form, such as gold.
All transition metals have more than
one oxidation state. Most transition metals (TMs) can complex with other species (called ligands in TM
Complex jargon) by giving their electrons to them, forming a complex. These ligands, which are the nearest
neighbor atoms to the metal center, constitute the inner (or rst) coordination sphere. Complexes may be
either neutral or charged and have distinctive properties that may be quite unlike those associated with their
constituent molecules and ions, each of which is capable of independent existence. An example of a charged
complex is ferricyanide, [Fe (CN)6
−3
]
+3
. The Fe
−
and CN
ions found in the ferricyanide complex ion exist
as independent species and in other compounds. The transition metals are well known for forming a large
number of complex ions. In this experiment we will synthesize a transition metal complex containing cobalt,
Co(III), and ethylenediamine.
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CHAPTER 10. TRANSITION METALS
96
10.1.3.1 Stereochemistry
The most common coordination numbers (the number of individual ligands bound) are two, four, and six,
with geometries illustrated in Fig 1:
Figure 10.1
Fig 1. Common geometries for complex ions. (A) linear, (B) square planar, (C) tetrahedral, and (D)
octahedral
−
Complexes of Cu(I), Ag(I), Au(I) and some of Hg(II) form linear structures (A) such as Cu (CN)2 ,
+
Ag (NH3 )2 , etc. Four-fold coordination (C) is not too common with transition metals, and the square
planar geometry (B) occurs in complexes of Pd(II), Pt(II), Ni(II), Cu(II), and Au(III). Six-fold coordination
(D) is the most common and in fact the one we will study in this laboratory exercise.
A ligand that is capable of occupying only one position in the inner coordination sphere by forming only
F − , Cl− , OH− ,
−
H2 O, NH3 and CN . If the ligand has two groups that are capable of bonding to the central atom, it is
one bond to the central atom is called a monodentate (one tooth) ligand. Examples are
called a bidentate ("two teeth") ligand, and so forth. An example of a bidentate ligand is ethylenediamine
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97
(CH2 NH2 CH2 NH2 ),
which is commonly abbreviated "en". Both nitrogen atoms in "en" can bond to the
central atom in a complex at the same time.
Complex ion salts with the same chemical formulas often behave dierently because the same number of
atoms can be arranged into dierent forms called isomers. Hydrate isomerism is illustrated by the following
example: There are three distinct compounds with the formula Cr (H2 O)6 Cl3 . One of these, violet in color,
reacts immediately with AgNO3 to precipitate all of the chlorines as AgCl. The second is light green but
only 2/3 of the chlorine is precipitated as AgCl.
The third compound is dark green and only 1/3 of the
chlorine is precipitated as AgCl. The last compound has only one reactive Cl, so apparently two chlorines
in this compound are bonded tightly to the Cr and are not available for reaction. We might thus write this
compound as [CrCl2
(H2 O)4 ] · (H2 O)2 ,
where the species within the brackets are regarded as ligands bonded
fairly strongly to the central chromium, and this species would behave as a single ion in solution. i.e., in
aqueous solution,
+
[CrCl2 (H2 O)4 ] Cl · (H2 O)2 → [CrCl2 (H2 O)4 ] + Cl− + water
The light green compound with two reactive chlorines is apparently [CrCl (H2 O)5 ] Cl2
· H2 O,
while the
violet compound with three reactive chlorines is Cr (H2 O)6 Cl3 .
Closely related to hydrate isomerism is ionization isomerism, where an ion takes the place of water.
Consider two dierent compounds with the formula Co (NH3 )5 SO4 Br. One of these, [Co (NH3 )5 (SO4 )] Br,
appears red, whereas the other, [Co (NH3 )5 Br] SO4 , appears violet.
In addition to these coordination sphere isomers there are geometrical isomers, which have coordination
spheres of the same composition but dierent geometrical arrangement.
Geometrical isomers are distinct
compounds and can have dierent physical properties (although often not too dierent) such as color, crystal
structure, melting point, and so on. For example, dichlorodiamine platinum (II) occurs in the square planar
geometry (B) so the chlorine ligands can be either next to one another (cis) or opposite from one another
(trans). The compound you will synthesize has an octahedral geometry with two (bidentate) "en" ligands,
and two nitro
(NO2 )
ligands. The geometrical isomer you will make is the trans form, in which the NO2
ligands are not adjacent to one another. This dierence between cis and trans octahedral isomers is shown
in Fig 2.
Figure 10.2
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CHAPTER 10. TRANSITION METALS
98
Fig 2.
The trans and cis geometrical isomers for octahedral complexes with two bidentate (en) and
monodentate
(NO2 )
ligands specically dinitrobis(ethylenediamine)Co(III). The two black balls represent
the NO2 ligands and the two pairs of linked white balls represent the two ethylenediamine ligands. Cis and
trans describe the relationship (relative position) between the two NO2 ligands.
In the procedure that follows we start with a cobalt solution made from the salt hexaquacobalt(II) nitrate,
[Co (H2 O)6 ] (NO3 )2 .
2+
−
When this salt dissolves it ionizes to form two ions of NO3 and one of Co (H2 O)6 . We
wish to prepare a Co(III) compound of ethylenediamine, so we must add ethylenediamine (en) and oxidize
the Co(II) to Co(III). Because Co(II) is more reactive than Co(III), we allow it to react with (en) rst, and
then oxidize the resulting complex ion.
In aqueous solution (en) reacts with water to produce OH
−
ions
which can also bind to Co(II), so the pH is adjusted close to 7 rst by adding HNO3 . (Other acids would
introduce new ligands to compete for the Co.) NaNO2 is added to provide the ligands that will be trans in
the nal compound. Lastly, Co(II) is oxidized to Co(III) by bubbling oxygen through the solution.
10.1.4 Experimental Procedure
1. Use your 10 mL graduated cylinder to measure out 20 mL of the 20% by weight solution of ethylenediamine in dilute HNO3 .
2. Pour it into a clean 125 mL Erlenmeyer ask. Rinse the graduated cylinder with about 5mL of deionised
water (DI water from white handle faucet) and add the rinse water to the ask. Set this aside for a
moment and prepare the second set of reactants as described below.
3. Weigh out 9.0 g of hexaquacobalt(II) nitrate and 6.0 g sodium nitrite ( NaNO2 ) using a rough balance
(Record mass on report form). Add these reactants to approximately 15 mL of DI water in an Erlenmeyer ask. After they have dissolved, add the neutralized ethylenediamine solution prepared in steps
1-2. Record your observations.
4. For the next set of instructions, refer to the diagram below. Fit a piece of rubber tubing over an inert
gas "IG" tap (on benchtop) and open the valve slowly to obtain a gentle ow of oxygen. Then insert a
Pasteur pipet into the other end of the rubber tubing. CAUTION: Too high a gas ow might blow the
pipet out of the tubing and cause serious injury. Always adjust the valve carefully while pointing your
pipet in a safe direction. Test the ow by immersing the pipet tip in a beaker of waterit should bubble
vigorously, but not enough to cause much splashing. When the ow is set to your satisfaction, immerse
the tip of the pipet in the Erlenmeyer ask containing the reaction mixture.
Secure the ask to a
stand with a clamp because the reaction mixture may need about 10 minutes of moderately vigorous
bubbling to reach completion. Record your observations.
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Figure 10.3
Figure 3. The bubbling apparatus.
1. After about 10 minutes of bubbling, turn o the gas ow and immerse the ask in ice water. This will
cause further crystallization. After approximately 5 minutes in the ice bath, pour the ask's contents
through the lter crucible while it has suction applied using the setup shown below.
observations.
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CHAPTER 10. TRANSITION METALS
100
Figure 10.4
Figure 4.
Schematic diagram showing sintered-glass lter crucible mounted on suction ask with rubber
lter adapter. Clamp the lter ask to a support post to prevent breakage.
1. The crystals will remain in the crucible while the solution passes through.
Wash your crystals by
slowly pouring approximately 5 mL of ethanol over them while suction is applied. Why do we wash
with ethanol? Answer on lab report form.
2. The next step is recrystallization to obtain a more puried product. Transfer the product crystals to
a 250 mL beaker. Add about 80 mL of DI water and stir to dissolve the crystals. Gently heat the
beaker over a Bunsen burner (or on high on a hotplate if available), gradually bringing it to a `slight'
boil. Allow the solution to boil gently until its volume has been reduced to about 50 mL. Then let the
solution cooled to near room temperature, place the beaker into an ice bath (DO NOT PLACE THE
BEAKER IN THE ICE BATH WHILE HOT. IT WILL CRACK AND YOU WILL LOOSE YOUR
PRODUCT). Crystal growth should be immediately apparent. After a few minutes in the ice bath,
transfer the crystals into the lter crucible. To help with this transfer you may use a rubber policeman
on the end of a stirring rod. Remember that your crystals are water-soluble so if you use water in the
transfer you will lose the product. Apply suction and rinse the crystals three times with separate 5
mL portions of ethanol. Scrape the crystals onto a watch glass and place in your drawer to dry.
In terms of the materials used, the overall reaction is:
4{[Co (H2 O)6 ] (NO3 )2 } + 8NaNO2 + 8C 2 H4 (NH2 )2 + 4HNO3 + O2 (g) → 4trans−
[Co (en)2 (NO2 )2 ] NO3 + 8NaNO3 + 26H2 O
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101
+
NH2 CH2 CH2 NH3 .
From the reaction and quantities used, calculate the theoretical yield and your
percentage yield.
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102
10.2 Pre-Lab: Transition Metals
10.3 (Total 10 Points)
2 for the Pre-Lab
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Note: In preparing this report you are free to use references and consult with others. However, you may
not copy from other students' work (including your laboratory partner) or misrepresent your own data (see
honor code).
Name(Print then sign): ___________________________________________________
Lab Day: ___________________Section: ________TA__________________________
1. List and draw the common geometries transition metal complexes:
2. What are the two types of structural isomers for complex ion salts?
3. What are the two types of geometrical isomers for complex ion salts?
4. Why do we use Co(II) and then convert to Co(III) when synthesizing
4 − trans [Co (en)2 (NO2 )2 ] NO3 ?
5. List two common monodentate ligands and two common bidentate ligands:
10.4 Report: Transition Metals
On my honor, in preparing this report, I know that I am free to use references and consult with others.
3 for the Report Form
Hopefully here
Note: In preparing this report you are free to use references and consult with others. However, you may
not copy from other students' work (including your laboratory partner) or misrepresent your own data (see
honor code).
Name(Print then sign): ___________________________________________________
Lab Day: ___________________Section: ________TA__________________________
Date ________________ Lab Section___________
Note: In preparing this report you are free to use references and consult with others. However, you may
not copy from other students' work (except to compile the group data set) or misrepresent your own data.
10.4.1 1. Synthesis
A. Volume of 20% ethylenediamine solution used ______ (r = 0.980 g/mL)
Compound
Weight
Moles (Molar weight and stoichiometric coecient)
ethylenediamine
[Co (H2 O)6 ] (NO3 )2
NaNO2
[Co (en)2 (NO2 )2 ] NO3
Table 10.1
2
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10.4.2
10.4.3 a. Observations
1. Record your observations after adding the neutralized ethylenediamine solution.
1. Record your observations after 10 minutes of moderately vigorous bubbling.
2. Record your observations after pouring the ask's contents through the lter crucible while suction is
applied.
10.4.4
10.4.5 b. Questions
1. Why do we wash the crystals with ethanol?
1. Give the net chemical equation for the reaction, writing dissociated reactants as ions, the solid product
as an undissociated salt, and including all other ionic and neutral species needed to balance charge
and mass. Omit any spectator ions that would appear in equally on both sides.
1. Which is the limiting reactant in your experiment?
1. Calculate the maximum weight of product you would have obtained if the limiting reactant had reacted
fully. This is the theoretical yield. What is your percent yield (the actual yield divided by theoretical
yield)?
Theoretical yield _______g Actual yield _______ g
1. Is the yield less, same, or more than the theoretical yield? Give reasons for why the actual yield is
dierent theoretical yield.
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CHAPTER 10. TRANSITION METALS
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Chapter 11
1
Physical Properties of Gases
11.1 Physical Properties of Gases
11.1.1 Objectives
•
Learn and understand physical properties of gases and explain observations in terms of the kinetic
molecular theory of gases.
•
Plot and calculate the root mean square speed of the Carvone molecules. (Comparison with speed in
vacuum).
•
Estimate volume and volume change of a balloon when it goes from room temperature (RT) to liquid
nitrogen temperature.
•
Observe and explain behavior of gas in: a soda can, a balloon in a ask, Cartesian diver, etc., when a
change in pressure or temperature is applied.
11.1.2 Grading
You grade will be determined according to the following:
•
•
•
Pre-lab (10%)
Lab Report Form (80%)
TA points (10%)
11.1.3 Introduction
Expanding and contracting balloons, imploding soda cans, exploding marshmallows are just some of the
demonstrations that are often used to illustrate the empirical gas laws and the kinetic molecular theory
of gases.
In this experiment, you will be performing these and other `demonstrations' and using your
understanding of the physical properties of gases to explain your observations.
There will be demonstrations laid out at seven dierent stations (2 sets at each) around the room and
you will go in 2 groups of 4 people (two sets of lab partners) to each station (you don't need to start with
#1). If your group is assigned or start with, for example 5, you should then follow the following order: 5, 6,
7, 1, 2, etc. Your group should spend no more than 15 minutes at each station. Perform the experiment by
following the instructions placed at each station. Then discuss your observations with your group. For each
of the activities, it is important to ask yourself what is going on, "how can our observations be explained
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CHAPTER 11. PHYSICAL PROPERTIES OF GASES
106
using the kinetic molecular theory of gases?" Remember that for some demonstrations calculations may be
required also. Be thorough and precise in your explanations.
CAUTION: Important Safety Notes:
Remember to use tongs, hot grips as appropriate when dealing with hot liquids, vapors and containers.
Liquid nitrogen is extremely cold, with a boiling point of -196[U+F0B0] C and if it comes into contact
with skin can result in severe frostbite.
The vacuum dessicator should be regarded as a potential implosion hazard when evacuated. Handle it
carefully.
When doing the egg experiment do not put hot ask immediately in the water bath (let it for at least 3
minutes sitting on the bench) it will crack and you may have to pay for it if it breaks.
Observe and record what happens in your laboratory report form.
You are encouraged to discuss among yourselves possible explanations to your observations.
11.1.4 Experimental Procedure
11.1.4.1 Diusion:
The goal of this experiment is to measure the rate of diusion of Carvone, a major component of spearmint
oil. Find an area where there are few drafts and the air does not already smell of spearmint. (You may go
to the hallway to perform the experiment)
Stand in a line, with the rst person in the group holding the bottle of Carvone and several paper towels.
All four people should be 1 meter apart. You will need to know the distance each person is from the bottle
of Carvone.
The fourth person should act as the timekeeper.
When the timekeeper gives the signal, the
rst person should place a few drops of Carvone on the paper towels. Record the time that it takes for each
person to smell the Carvone. Seal the paper towel in a plastic bag when you are nished.
After the odor has dissipated, repeat the experiment twice.
Using Excel plot the data in distance traveled versus time. Obtain a least squares t (R value) for this
data and determine from it the rate of diusion of Carvone in meters per second. Create a graph for each
trial. Calculate the average of the rates for the three trials. Calculate the root mean square speed of carvone
molecules at 25[U+F0B0]C. Your TA will help you with this equation. Compare the result with the diusion
rate you measured. If they are signicantly dierent, oer an explanation. Would the diusion take place
faster in a vacuum?
Note: You should spend no more than one-half hour preparing the plots. Please stagger yourselves so
that everyone has an opportunity to get to the computer stations.
11.1.4.2 Gas Laws in a Soda Can:
Pour 15 mL of water into an aluminum soda can. Set the can on a hot plate and turn on to a high temperature
setting. While the can water heats, ll a 1000-mL beaker with cold water (You may have a metal tin set out
for this purpose). Continue heating the can until the water inside boils vigorously and until steam escapes
from the mouth of the can for about 20 seconds.
Using the hot grips to grip the can near the bottom, quickly lift the can from the burner and invert (so
water covers the mouth of the can) it in the beaker of cold water. Describe what happens. Explain why it
happens. You may repeat this experiment using a second soda can if you wish. Why is it necessary to invert
the can in the water? What would happen if a rigid container were used?
11.1.4.3 Balloon in liquid nitrogen:
Review the safety notes above regarding the handling of liquid nitrogen.
Inate a balloon and tie the end (Several balloons may have already been inated and tied). Using tongs,
place the balloon in a Dewar ask containing liquid nitrogen. After the balloon stops changing size, remove
it from the Dewar and allow it to warm to room temperature. Observe and record the changes (you should
be able to measure the radius and estimate volume).
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107
Estimate the size of the balloon in liters. What is the pressure inside the balloon before it is placed in
the liquid nitrogen? What is the pressure inside the balloon after it is placed in the liquid nitrogen? Use the
ideal gas law to calculate the percent change in volume expected on going from room temperature to liquid
nitrogen temperature. Is the volume of the cold balloon consistent with what you calculated, or is it larger
or smaller? Suggest an explanation for your observation. Explain all of your observations in detail using the
kinetic molecular theory of gases. How does the liquid nitrogen cool the gas in the balloon?
11.1.4.4 Tygon tube in liquid nitrogen:
Review the safety notes above regarding the handling of liquid nitrogen.
Place a 2 foot long tygon clear
tube in a Dewar with liquid nitrogen. Observe what happens and explain.
11.1.4.5 Balloon in a ask:
Place about 5 mL of water in a 125-mL Erlenmeyer ask. Heat the ask on a hot plate until the water boils
down to a volume of about 1 mL. Meanwhile, inate a balloon and then let the air out (this may not be
necessary if balloons on table have been previously used). Remove the ask from the heat, hold it with a
towel, and immediately place the open end of the balloon over the mouth of the ask. Observe the eect as
the ask cools. Can you get the balloon back out again? If you can, How?
11.1.4.6 Cartesian diver:
The Cartesian diver is named for Rene Descartes (1596-1650), noted French scientist and philosopher. At
this station, you will nd a plastic soda bottle containing a medicine dropper, water, and air. Squeeze the
bottle.
What happens? Why?
11.1.4.7 The Egg:
Lightly grease the inside of the neck of a 1 L Erlenmeyer ask with stopcock grease. Clamp the ask onto
the stand. Place about 5 mL
H2 O
in the ask and gently warm it with a Bunsen burner until the water
vaporizes. Do not boil the water to dryness. Meanwhile, prepare an ice water bath in an evaporating dish.
While the ask is warm, seat the egg, narrow end down, in the mouth of the ask. Unclamp the ask, allow
to cool slightly sitting on the bench and then immerse it in the ice water. (Read the safety notes above to
avoid breaking the ask)
Can you get the egg back out again?
Assuming that the ask reaches the maximum vacuum (minimum pressure) possible before the egg is
drawn into the ask, calculate the minimum pressure reached in the ask.
11.1.4.8 Expanding balloon:
Partially inate a balloon. Place the balloon inside the vacuum chamber and close the chamber with the
black rubber circle and the top of the chamber carefully centered on the base (A partially inated balloon
may already be in the dessicator). Close the needle valve (at the bottom of the black rubber tubing) by
turning it clockwise. Turn the stopcock to the up position to connect the chamber to the vacuum pump.
What happens? To open the chamber, turn the stopcock to the left position and open the needle valve.
11.1.4.9 Bonus 2 points:
1pt to name a real life example of the physical properties of gases at work
1pt for a good explanation of how and why it works according to what you have learned in the lab.
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CHAPTER 11. PHYSICAL PROPERTIES OF GASES
108
11.2 Pre-Lab: Physical Properties of Gases
11.3 (Total 10 Points)
2 for the Pre-Lab
Hopefully here
Note: In preparing this report you are free to use references and consult with others. However, you may
not copy from other students' work (including your laboratory partner) or misrepresent your own data (see
honor code).
Name(Print then sign): ___________________________________________________
Lab Day: ___________________Section: ________TA__________________________
1. Dene diusion and write down equation for diusion rate:
2. Write equation for ideal gas law and describe each term.
3. Dene Charles', Boyle's, and Avogadro's law:
4. Fill the blanks (which law applies):
5. When temperature increases in a close balloon the volume ________.
_______Law
•
When pressure is applied to a close balloon the volume _________.
_______Law
•
When temperature decreases in a close balloon the pressure _________.
_______Law
11.4 Report: Physical Properties of Gases
11.5 (Total 80 Points)
3 for the Report
Hopefully here
Note: In preparing this report you are free to use references and consult with others. However, you may
not copy from other students' work (including your laboratory partner) or misrepresent your own data (see
honor code).
1. Diusion:
At the end of your report attach the graphs of each trial.
The average of the rates for the three trials isThe root mean square speed of carvone molecules at 25[U+F0B0]C is
Compare the result with the diusion rate you measured.
If they are signicantly dierent, oer an
explanation.
Would the diusion take place faster in a vacuum?
2. Gas Laws in a Soda Can:
2
3
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http://cnx.org/content/m15545/latest/ReportPP07.doc
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109
Describe what happens.
Explain why it happens. You may repeat this experiment using a second soda can if you wish.
Why is it necessary to invert the can in the water?
What would happen if a rigid container were used?
3. Balloon in liquid nitrogen:
The estimated size of the balloon in liters is
What is the pressure inside the balloon before it is placed in the liquid nitrogen?
What is the pressure inside the balloon after it is placed in the liquid nitrogen?
Use the ideal gas law to calculate the percent change in volume expected on going from room temperature
to liquid nitrogen temperature.
Is the volume of the cold balloon consistent with what you calculated, or is it larger or smaller?
Suggest an explanation for your observation. Explain all of your observations in detail using the kinetic
molecular theory of gases.
How does the liquid nitrogen cool the gas in the balloon?
4. Balloon in a ask:
What was the eect as the ask cools?
Can you get the balloon back out again?
5. Kissell's tygon tube:
What happens?
Why?
6. Cartesian diver:
What happens?
Why?
7. The Egg:
What happens?
Why?
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CHAPTER 11. PHYSICAL PROPERTIES OF GASES
110
Can you get the egg back out again?
The minimum pressure reached in the ask is -
8. Expanding balloon:
What happens?
Moore's bonus 2 points:
1pt to name a real life example of the physical properties of gases at work
1pt for a good explanation of how and why it works according to what you have learned in the lab.
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ATTRIBUTIONS
111
Attributions
Collection:
Honors Chemistry Lab Fall
Edited by: Mary McHale
URL: http://cnx.org/content/col10456/1.16/
License: http://creativecommons.org/licenses/by/2.0/
Module: "Initial Lab: Avogradro and All That"
By: Mary McHale
URL: http://cnx.org/content/m15093/1.1/
Pages: 1-8
Copyright: Mary McHale
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Module: "Stoichiometry: Laws to Moles to Molarity"
By: Mary McHale
URL: http://cnx.org/content/m15095/1.14/
Pages: 9-16
Copyright: Mary McHale
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Module: "VSEPR: Molecular Shapes and Isomerism"
By: Mary McHale
URL: http://cnx.org/content/m15100/1.8/
Pages: 17-21
Copyright: Mary McHale
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Module: "Beer's Law and Data Analysis"
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Module: "Hydrogen and Fuel Cells"
By: Mary McHale
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Pages: 33-47
Copyright: Mary McHale
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Module: "The Best Table in the World"
By: Mary McHale
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Pages: 49-55
Copyright: Mary McHale
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ATTRIBUTIONS
112
Module: "Bonding 07"
By: Mary McHale
URL: http://cnx.org/content/m15205/1.4/
Pages: 57-66
Copyright: Mary McHale
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Module: "Solid State and Superconductors"
By: Mary McHale
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Pages: 67-84
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Module: "Organic Reactions"
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Pages: 85-93
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Module: "Transition Metals"
By: Mary McHale
URL: http://cnx.org/content/m15503/1.4/
Pages: 95-103
Copyright: Mary McHale
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Module: "Physical Properties of Gases"
By: Mary McHale
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Pages: 105-110
Copyright: Mary McHale
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