Download Lab #3 – A Microscale Study of Chemical Changes

Lab #1 – Scientific Observation and Description
In order to become a good scientist, you must learn how to become a good
observer. An observation is the process of using your senses to gather information. There
are five senses that a good scientist must use. These senses are: sight, hearing, smell, taste
and touch. In this lab you will learn how to use your senses in order to observe a variety
of science “experiments”.
We tend to think of ourselves as good observers. Yet there is much more to
observation than meets the eye. It takes concentration, alertness to detail, ingenuity and
patience. It also takes practice. Try it yourself. See how complete a description you can
write about a familiar object – a burning candle. Be “scientific” about this and start with
an experiment. This means you will observe a burning candle in a laboratory, because
the laboratory is a place where conditions can be controlled.
Most of our observations tend to be qualitative, like color, texture, sound, etc. In
science, while these observations are very important, we need to realize that quantitative
observations, those that can be measured or quantified, are often even more important.
After all, these can be measured!
How can we know what conditions in an experiment should be controlled?
Sometimes the important conditions are difficult to discover, but an experiment can be
meaningless unless the conditions that matter are controlled. Here are some conditions
that might be important but are not important in this one:
1.
The experiment is done on the second floor.
2.
The experiment is done in the daytime
3.
The room lights are on.
Here are some conditions that might be important in your experiment:
1.
The lab bench is near the door.
2.
The windows are open.
3.
You are standing close enough to the candle to breathe on it.
Why is the second set of conditions important? It is important because all three
conditions relate to a common factor: a candle does not burn well in a draft.
Important conditions are often not as easily recognized as these. A good
experimentalist pays much attention to the discovery and control of conditions that are
important.
Purpose:
What is your reason for doing this lab? What are you trying to discover/ learn?
(2-3 sentences)
Pre-lab:
1.
What is an observation?
2.
What is the difference between a qualitative and quantitative observation?
3.
Why are controls important in experiments?
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Procedure:
1.
Obtain a candle, beaker, piece of foil and match
2.
Place candle on foil and begin observations.
3.
Record your observations
4.
After you have observed your candle unlit, light your candle.
5.
Observe your lit candle, record data.
6.
Cover candle with beaker and observe.
7.
Record your observations
Post-lab/ analysis:
1.
How many observations did you make, total?
2.
Label your observations as either qualitative or quantitative.
3.
Which type of observation did you make most often? Why?
4.
What are some things that you could do to improve your observational skills?
Conclusion:
What did you learn from doing this lab activity? What would you have changed
to make this a better experience? Do you have any questions? (2-3 sentences)
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Lab #2 – Laboratory Techniques
Read lab 1A, pp 702-705 in your textbook, Holt Chemistry, Visualizing Matter.
Purpose:
What is your reason for doing this lab? What are you trying to discover/ learn?
(2-3 sentences)
Pre-lab:
1.
What tools will you be using in this lab?
2.
Where are the following items, and how is each item used? Eyewash station, fire
blanket, fire extinguisher.
3.
Copy the data tables below into your notebook. (note that there are columns for
finding masses with a triple or quad beam balance as well as with a digital balance)
Data Table 1
Material
Mass (g) on triple-beam
Mass (g) on digital
balance
Weighing paper
Weighing paper + NaCl
Data Table 2
Material
Mass (g) [tb]
Step 11
Mass (g) [tb]
Step 12
Mass (g) [dig]
Step 11
Mass (g) [dig]
Step 12
Empty beaker
Beaker + 50 mL
water
50 mL of water
Beaker + 100
mL of water
100 mL of
water
Beaker + 150
mL of water
150 mL of
water
Procedure:
Follow the procedure as written in your book with one exception: when you are asked to
find the mass of water in steps 11 and 12, you will not only find the mass on a triple or
quad beam balance, but on a digital balance as well.
Post-lab:
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1.
Based on your observations, which type of flame is hotter, the flame formed when
the air ports are open or the flame formed when they are closed? What is the hottest part
of the flame?
2.
Which of the following measurements could have been made by your balance:
3.42 g of glass, 5.66672 g of aspirin, or 0.000017 g of paper? Why?
3.
Make a graph of mass versus volume for data from steps 11 and 12. The mass of
water (g) should be graphed along the y-axis as the depended variable, and the volume of
water (mL) should be graphed along the x-axis as the independent variable. Use as much
of your graph paper as possible, and use a separate color for the best-fit line for each of
the sets of data.
4.
When methane (natural gas) is burned, it usually produces carbon dioxide and
water. If there is a shortage of oxygen, the flame is not as hot, and black carbon is
formed. Which steps in the lab demonstrate these flames?
5.
Which is the most accurate method for measuring volumes of liquids, a beaker or
a graduated cylinder? Explain why.
6.
The density of water is equal to its mass divided by its volume. Calculate the
density of water using your data from step 11 (use your data from the triple-beam
balance) Then calculate the density of water using your data from step 12. How do they
compare?
Conclusion:
What did you learn from doing this lab activity? What would you have changed
to make this a better experience? Do you have any questions? (2-3 sentences)
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Lab #3 – Accuracy and Precision
Read lab 2A, pp 708-711 of your textbook, Holt Chemistry, Visualizing Matter.
Purpose:
What is your reason for doing this lab? What are you trying to discover/ learn?
(2-3 sentences)
Pre-lab:
1.
Define accuracy and precision.
2.
How is the uncertainty of a measurement expressed?
3.
Copy data tables 1 and 2 from p 709 into your lab notebook.
Procedure:
Follow the procedure as written in the textbook.
Post-lab:
1.
What is the uncertainty in a measurement made by each of these devices: the
ruler, the 100 mL graduated cylinder, the 24 mL graduated cylinder, and the balance?
2.
Calculate the volume of the 100 mL graduated cylinder up to its 50 mL mark.
(Hint:
V = r2h for a cylinder)
3.
Assume the accepted value for the volume of a graduated cylinder up to the 50
mL mark is 50.00 mL. Calculate the percentage error in your calculations from question
2.
4.
Subtract the volume of the water from the combined volume of the metal and the
water to calculate the volume of the metal alone. Subtract the mass of the water and
cylinder from the mass of the metal, water, and cylinder to determine the mass for the
metal, then calculate its density. Show all your work!
5.
Record the metal-density calculations made by the other teams, and calculate the
average density of the unknown metal.
6.
Compare your value and the class average of the unknown metal’s density with
the densities for metals given in the handbook. Determine the identity of the metal.
7.
A store has a balance with a scale marked in gram units, and each unit is divided
in half by a smaller mark. A student working in the store after school measured a sample
mass as
5.367 g using this balance. What is wrong with this measurement?
Conclusion:
What did you learn from doing this lab activity? What would you have changed
to make this a better experience? Do you have any questions? (2-3 sentences)
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Lab #4 – Separation of a Mixture (percent composition) from Flinn Scientific
ChemTopic Labs
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7
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Lab #5 – Using the Scientific Method with a “Black Box” (adapted from Lab-aids #100
Ob-scertainer kit)
One of the first topics one focuses on in any science class is that of the scientific
method. There is not just “one formula” for the scientific method, but there are general
guidelines. In many ways, the scientific method is a cycle, as pictured in the diagram
below:
For our purposes, in chemistry class, we will focus on the following steps:
1.
Question – “what do I want to learn more about?”
2.
Hypothesize – research to help you make an educated guess, and then
answer your question
3.
Experiment – test your hypothesis
4.
Observe and record – make careful observations (both direct and
indirect) and write down what happens.
5.
Analyze – use your information to draw conclusions about your
experiment. Was your hypothesis correct?
6.
Share results – explain your results by presenting your experiment,
observations, and conclusions.
In our “black box” activity, we will be using careful observations to check
the accuracy of our hypotheses. These observations will be indirect, meaning
that we will use senses other than sight to determine the configuration of the black
box. In science, many things cannot be seen directly, therefore learning how to
make indirect observations is very important.
Our hypothesis will be an intelligent guess as to what is “going on” inside
the container, and how it relates to the inside partitions. There is a steel ball
rolling around in the container, and our goal is to determine what the partitions
inside look like without actually seeing them.
Purpose:
What is your reason for doing this lab? What are you trying to discover/ learn?
(2-3 sentences)
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Pre-lab:
1.
What is the scientific method? (general description)
2.
What are the steps to the scientific method as described in the lab?
3.
What is the difference between direct and indirect observations?
4.
Which type of observation will you be using in this lab and why?
5.
Set up your data collection as shown below. Make sure there are enough spaces
for 5 different ob-scertainers.
Ob-scertainer # ______
Hypothesis
Re-test (final hypothesis)
Actual model
Procedure:
1.
Obtain an ob-scertainer and record the # in your data section.
2.
Carefully shake and tilt your ob-scertainer.
3.
From the sound and path of the steel ball, determine a hypothesis for the shape
and location of the partitions within the ob-scertainer and record this in your first data
blank.
4.
Test this hypothesis by moving the ball along the partitions according to your
hypothesis. If you wish to make changes, do so in the second circle. This will be your
final hypothesis.
5.
Save the third circle to fill in after your teacher has revealed the actual
configuration.
6.
Continue testing at least four more ob-scertainers. Some of them may be more
difficult than others, but do not spend more than 5 minutes on each. DO NOT OPEN
THE OB-SCERTAINERS!!!!!
Post-lab questions/ analysis:
1.
Were any of your hypotheses correct?
2.
Which one(s) were not correct?
3.
Why do you think some were harder to guess than others?
4.
How did you use the scientific method in this experiment?
Conclusion: (should be at least 2-3 sentences)
What did you learn from doing this lab activity? What would you have changed
to make this a better experience? Do you have any questions?
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Lab # 6 - A Microscale Study of Chemical Changes (Adapted from Addison-Wesley’s
Microscale Chemistry)
Chemistry is a science that investigates changes in matter. Chemical reactions are
the changes chemicals undergo. The changes you observe are called “macroscopic
changes.” Often these changes, such as color change, the formation of a solid, or the
formation of gas bubbles, are visible. Thus, though we cannot see the atoms and
molecules reacting, we can see indications that chemical changes have taken place.
Different atoms and molecules often react in different ways. Chemistry attempts
to explain macroscopic changes in terms of the behavior of atoms and molecules, that is,
on the submicroscopic level. You can use these different reactions to detect the presence
of specific kinds of chemicals in mixtures.
In this lab, you will study some reactions of common chemicals contained in
consumer products. You will observe the notable macroscopic changes these chemicals
undergo. In later labs, you will interpret these macroscopic changes in terms of
submicroscopic changes, the behavior of atoms and molecules. As the name implies,
submicroscopic changes are changes we cannot see, even with a microscope. The
essence of understanding chemistry is to infer from macroscopic changes the
submicroscopic behavior of atoms and molecules.
Pre-lab:
1.
What is a macroscopic change?
2.
What are three macroscopic changes that can indicate a chemical reaction is
taking place?
3.
What is a submicroscopic change?
4.
Create a data table in your notebook in which you can record any macroscopic
changes you might view when mixing chemicals as outlined on the experimental page.
Procedure:
See experimental page for the actual procedure
Clean-up
Avoid contamination by cleaning up in a way that protects you and your
environment. Carefully clean the small-scale reaction surface by absorbing the contents
onto a paper towel, wipe it with a damp paper towel, and dry it. Dispose of the paper
towels in the waste bin. Wash your hands thoroughly with soap and water.
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Experimental page:
Use small-scale pipets to put 2 drops of each chemical on the X’s in the indicated
spaces below. For the background contrast, view the drops on the black and white
backgrounds provided by the X’s. Stir each mixture by blowing air through an empty
pipet. Record what you see in your data table.
a.
b.
c.
d.
e.
f.
g.
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X
X
X
X
X
X
X
NaHCO3
+
HCl
h.
HCl
+
Blue dye
i.
Blue dye
+
NaOCl
Now add
1 drop of
HCl
j.
NaOCl
+
KI
Now add
1 drop of
starch
k.
KI
+
Pb(NO3)2
l.
Pb(NO3)2
+
CaCl2
m.
CaCl2
+
NaHSO4
Be
patient!
n.
X
X
X
X
X
X
X
NaHSO4
+
Na2CO3
Na2CO3
+
Phen.
Phen.
+
NaOH
NaOH
+
AgNO3
AgNO3
+
NH3
NH3
+
CuSO4
CuSO4
+
NaHCO3
Place a drop
of this in
your lab NB
Post-lab:
1.
Sodium hydrogen carbonate (NaHCO3) is baking soda. When HCl is added to
NaHCO3, carbon dioxide bubbles are formed. What is the chemical formula for carbon
dioxide, and in what consumer product is this gas commonly found?
2.
Which of the other mixings gave a similar result?
3.
What gas do you think results from the mixing in Question 2?
4.
Sodium hypochlorite, NaOCl, is a common ingredient in household bleaches and
cleansers. What happens to the color of blue dye when both HCl and NaOCl are added?
5.
A precipitate is a solid that separates upon mixing solutions. Which reaction
produced a very bright yellow precipitate?
6.
Observe the scrap of paper you used to absorb the AgNO3 + NH3 mixture.
What evidence do you see that indicates that silver compounds are light-sensitive?
Conclusion:
Two – three complete sentences to “wrap-up” your lab.
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Lab #7 - Penny Density
Introduction:
Today’s penny is quite different from pennies minted 3 decades ago. Before
1982, pennies were made of an alloy of copper. An alloy is a solid solution consisting of
atoms of different metals. After 1983, pennies have been made with copper coating a
core composed of another metal. These differences in the composition of older and more
recently minted pennies have resulted in differences in the penny’s characteristics,
including its density, or mass per unit volume. In this experiment, you will determine
and compare the densities of pennies minted before 1982 and after 1983. We will use
density data to try to identify the metal used in the core of pennies minted after 1983
Pre-lab:
1.
Define the terms meniscus, alloy and density.
2.
What are some possible reasons for using a metal other than copper as the core of
the penny now?
3.
Copy the data tables below into your lab notebook. Make sure to leave enough
room to enter the correct information.
Table 1 (pre-1982 pennies)
# of pennies
Mass (g)
Total volume (mL)
Net vol. of pennies
(mL)
5
10
15
20
25
Table 2 (post-1983 pennies)
# of pennies
Mass (g)
Total volume (mL)
Net vol. of pennies
(mL)
5
10
15
20
25
Procedure:
1.
Work with one set of pennies at a time. Find the mass of 5 pennies from one set.
Record the mass in your data table.
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2.
Add 5 more pennies to the first group. Obtain the mass of these 10 pennies and
record their mass in the appropriate space in your data table. Repeat this procedure for
the remaining pennies (5 at a time)
3.
Fill a 50 mL graduated cylinder to the 20 mL mark with tap water. Be sure to use
the bottom of the meniscus (the curve of water in a container due to surface tension) to
measure the water level.
4.
Still working with the same set of 25 pennies, gently drop 5 of the pennies into
the graduated cylinder. Record the new water level in the appropriate space of your data
table.
5.
Add 5 more pennies to the graduated cylinder, making a total of 10 pennies.
Record the water level in your data table. Repeat this procedure for the remaining
pennies (5 at a time).
6.
Discard the water. Dry the pennies with a paper towel and either pass them to
another group to use or return them to the teacher.
7.
Repeat steps 1-6 using the 25 pennies in the other set of coins. Record your data
in the other data table.
8. Complete your data tables. Find the net volume of each group of pennies by
subtracting 20 mL (your starting volume) from the total volume recorded for each group.
Enter the net volume in the appropriate column.
Post-lab:
1.
Construct a graph of your results. Let the y-axis reflect the mass of the pennies,
and the x-axis represent the net volume. Plot the data for the pre-1982 pennies first.
Then draw the best-fitting straight line (the straight line that connects as many points as
possible).
2.
On the same graph, plot the data for the pennies minted after 1983. Draw the
best-fitting straight line. Use a different color for this graph.
3.
Find the slope of each line. Recall that slope = y/x (also recall that you should
use points on the line, not necessarily data points)
4.
Find the density of copper from a resource book.
5.
Is the density of copper similar to the slope of the line you determined for the pre1982 pennies?
6.
What is the density of the pennies minted since 1983? (this is the slope from your
post-1983 graph)
7.
Compare the slope of the line for these pennies with the density listed for each
metal in your resource. Which element(s) have a density similar to your results? (name
at least two)
8.
Be sure to attach your graph to your notebook!
Conclusion: Write a conclusion statement in your notebook summarizing what you
learned in this lab.
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Lab #8 – Rutherford Experiment
Pre-lab: (you may need your textbook for some of this - pp 84-85)
1.
Describe the Rutherford gold-foil experiment.
2.
What is an alpha particle?
3.
Which atomic model did Rutherford disprove with his experiment?
4.
What piece of the atom did Rutherford discover, and how did he describe it?
5.
The fact that the vast majority of alpha particles passed undeflected through the
gold foil indicates what about atoms?
Introduction:
In Rutherford’s day, Thomson’s “plum-pudding” model of the atom was used to
describe the atom. In this model, Thomson proposed that the electrons of an atom were
embedded in a positively charged ball of matter. One of Thomson’s students, Ernest
Rutherford quickly disproved this model.
In Rutherford’s experiment, he shot alpha particles, or positively charged helium
ions, at a thin piece of gold foil. Had Thomson’s model been correct, many of the alpha
particles would have been deflected. What Rutherford found, however, was that most of
the alpha particles shot at the foil passed straight through undeflected. A very small
number of particles were deflected, some even backward. He reasoned that only a very
concentrated positive charge, localized somewhere within the gold atom, could possibly
repel the fast-moving, positively charged alpha particles sufficiently to reverse their
direction of travel.
Rutherford hypothesized that this localized region, called the nucleus, must have
a large mass compared with the alpha particle, or else the incoming particle would have
knocked the positive charge out of the way. He concluded that most of the gold foil was
“empty space”.
Rutherford was able to determine the diameter of the nucleus using indirect
methods. After all, the nucleus is a very small piece of an atom, and an atom itself is too
small to be measured with ordinary devices.
In this activity, you will also be using indirect methods to determine the diameter
of a “nucleus”. In our case, the “nucleus” will be a marble. You will be rolling “alpha
particles” (other marbles) at the “nucleus” (target marble). From the ratio of collisions to
trials, you will determine the size of your nuclear marble. It’s a little bit like throwing
snowballs at a car while you’re blindfolded. If you have very few hits per number of
trials, the car will “feel” small.
We’re going to develop two expressions for the probability of a hit between the
rolling marble (RM) and the nuclear marble (NM). One expression of the probability (P)
of a hit would be the ratio of the path width within which RM could roll and collide with
the NM and the overall width (w) of the target area.
P = path width/ overall width
P = 2(R+r)/w (where R is the radius of the
RM
NM, r is the radius of the RM and R + r
NM
is the distance between the centers of
RM
the RM and the NM)
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If the number of target marbles is increased to N, the probability of a hit is
increased by a factor of N, thus P = 2(R + r)N/w
The probability of a hit can also be determined experimentally by counting the
number of trials and the number of hits such that
P = H/T
You now have two expressions for the probability of a hit. Assume that both of
the expressions are equivalent. Also, if the radius of the RM and NM are equal, then
R + r = d where d is the diameter of either one of the marbles.
Procedure:
1.
Place anywhere from 1-5 marbles in a 60 cm square target area. (this is
approximately an area of 3 tiles by 3 tiles)
2.
Roll marbles randomly toward the whole target area. (Don’t look at the NM
since you “can’t really see it anyway”.) If a RM hits more than one NM, it counts as just
one hit. If a RM goes outside the target area, it does not count as a trial.
3.
Continue rolling marbles randomly until your group has completed at least 200
trials. Make sure you keep a record of trials/ hits in your lab notebook.
Post-lab:
1.
Determine the radius of your NM using equations from the introduction.
2.
Actually measure the diameter of your NM (If you used more than one marble,
measure the diameter of the whole “nucleus”)
3.
Calculate your percent variation.
(actual-calculated) x 100
actual
Conclusion:
Don’t forget a conclusion
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Lab #9 - Conservation of Mass Lab; Reaction in a Bag
Introduction:
“Plop, plop, fizz, fizz oh, what a relief it is,” claims an old television ad for a
popular antacid. Just what is in the tablet that is relieving the upset stomach? What
reaction is causing the fizzing? Can you write a chemical equation for this process? With
a bit of investigating, you will be able to discover answers to all these questions.
As you learned in this chapter, Antoine Lavoisier, in the 18th century, formulated
the “LAW OF CONSERVATION” of mass, which states that matter can neither be
created nor destroyed. During a chemical reaction, the bonds of the reactants are broken
and rearranged to form new substances. Because matter must be conserved, these new
substances, or products, must contain the same number and type of atoms as the
reactants.
In this investigation, you will first verify the law of conservation of mass. Then in
the second part you will be given some known compounds to react.
Pre-Lab:
1.
Define reactants:
2.
Define products:
3.
What are 6 indicators a chemical reaction has happened? (see pp 18-19)
4.
Why is it necessary to use a zip-lock bag?
5.
If the density of water is 1.00 g/ml and you measure a volume of 25.00 mL. What
is the mass?
6.
What is the common name for sodium bicarbonate?
7.
What is the Law of Conservation of Mass?
8.
Copy the data tables in your lab notebook.
Materials
� Goggles
� Graduated Cylinder
� Electronic Balance
� Scoopula or spoon
� Sodium Carbonate
� Apron
� Zip – lock plastic bag
� Antacid tablet
� Calcium Chloride
� Bromthymol blue indicator
Procedure:
PART A~ Antacid tablet
1.
Put your goggles and apron on. Measure 25 mL of tap water into a re-sealable
plastic bag. Flatten the air out of the bag and seal it. Record its mass in data table 1.
2.
Record the mass of the antacid tablet in data table 1.
3.
Tip the bag sideways, and while holding the bag, pour 25 ml of water into one
corner.
4.
Place the antacid tablet in the bag, but DO NOT LET IT TOUCH THE WATER.
5.
Seal the plastic bag, while still holding onto the tablet and the water in the corner
of one of the bags.
6. Let the tablet drop into the water.
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7. Observe the reaction until it comes to a complete stop. You will know this when the
bubbling and fizzing stop.
8. Mass the bag and all the reactants and products and record the mass in data table #1.
PART B – CaCl2, NaHCO3 and water
1.
Calculate the total mass of the bag and reactants in each reaction and record these
values in the appropriate data table.
2.
Using the formula for the density of water, calculate the mass of the water.
Record the results in data table #2.
3.
Add 1 scoop of calcium chloride, CaCl2, to the second plastic bag.
4.
Add 1 scoop of sodium bicarbonate, NaHCO3 to the bag, and shake gently to mix.
5.
Determine the mass of the bag and its contents. Record this value in data table 2.
6.
Measure 25 mL of water into the graduated cylinder. Add 5 drops of bromthymol
blue indicator to the water.
7.
Tip the bag sideways, and while holding the solids in the bottom corner, pour the
liquids in the other corner. Twist the bag if you like so the solids will not get wet from the
water/bromthymol blue.
8.
Keeping the trapped air to a minimum, reseal the bag. Hold the bag and let the
liquid move from one end of the bag to the other until the content are mixed.
9.
Observe the reaction until it comes to a complete stop. Record your observations.
10.
Record the mass of the unopened bag in data table #2. Clean up your work area
and wash your hands before leaving the lab.
Observations
Data Table #1 - Do not forget units!!!!
Mass of bag and water
Mass of antacid tablet
Mass of all reactants (mass of bag and
water + mass of antacid tablet)
Mass of bag and products
Write observations here
Data Table #2 – Do not forget units!!!!!
Mass of bag and dry reactants
Volume of water
Mass of water
(hint: D = M/V ~ pre-lab question)
Mass of all reactants (mass of bag + mass
of water + dry reactants)
Mass of bag and products
Write observations here
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Table 3 – Physical properties of reactants
Substance
Property 1
Property 2
antacid tablet
calcium chloride
baking soda
Property 3
Post-lab:
1.
How do the values for total mass before and after each reaction demonstrate the
law of conservation of mass?
If your values do not match, what is a possibility for this discrepancy?
2.
In part A did you observe a physical or chemical change to the antacid tablet?
What is your evidence for this observation?
3.
In part B did you observe a physical or chemical change to the reactants (Calcium
chloride and sodium bicarbonate)? What is your evidence for this observation?
7.
Bromthymol blue is an indicator that turns yellow when present in an acidic
solution. Was your bag acidic or basic, and how do you know?
Conclusion: Don’t forget to write at least 2-3 sentences to conclude your lab!
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Lab # 10 – Flame Tests (see p 720-723 of textbook)
Pre-lab:
1.
Explain how and why different chemicals emit different colors of light.
2.
How will you determine the identity of the salts in your unknown solution?
3.
What might be a reason for using cobalt blue glass?
4..
Set up an appropriate data table (see p 721)
Procedure:
Follow procedure from the textbook except sketch the line spectrum you see rather than
writing a wavelength in nm.
Post-lab:
1.
For 3 of your samples, explain how the color of the flame and the line spectrum
relate.
2.
What ions were in your unknown solution, and how did you make that decision?
3.
How can a spectroscope help you to identify the components of a solution having
several different metal ions?
4.
A student observes three unknowns, all of which are some shade of red. How
could he correctly identify the metal ions present in each?
Conclusion:
Don’t forget to write at least 2-3 sentences to conclude your lab!
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Lab #11 – The Mendeleev Lab of 1869 – p 726 of textbook
Pre-lab:
1.
Who is credited with creating the periodic table?
2.
How was his arrangement of elements unique?
3.
Copy the data table from p 727 into your lab notebook. Be sure to leave enough
space for all nine unknown elements.
Procedure:
Follow the procedure from the textbook.
Post-lab:
1.
Summarize your reasoning for the identity you assigned each unknown.
2.
Describe any trends you saw in two of the families of elements.
Conclusion:
Don’t forget to write at least 2-3 sentences to conclude your lab!
24
Lab #12 – Periodic Table Trends
In this lab, you will investigate the properties of several elements on the periodic
table and classify them as metals, non-metals or metalloids.
The periodic table can be divided into three rather distinct groups of elements.
Metals are to the left side of the periodic table. They tend to be good conductors of heat
and electricity. They tend to have a high luster (are very shiny), are malleable (able to be
bent or hammered into sheets), and ductile (able to be pulled into thin wires). Not all
metals are malleable and ductile, but most are. Nonmetals are to the right side of the
periodic table, and tend to have properties opposite those of metals. Nonmetals tend to
be dull, brittle (break easily when hit) and non-conductive. Metalloids fall somewhere
between metals and non-metals in spatial relation on the periodic table and in their
properties.
Pre-lab Questions:
1.
In your lab notebook, draw the following table. Write each of the following
properties under the appropriate heading. Write the whole property, not just the letter of
the property.
METALS
a.
b.
c.
d.
e.
f.
g.
h.
i.
j.
25
NON-METALS METALLOIDS
good conductor of heat
poor conductor of heat (insulator)
semi conductor
shiny, high luster
solids tend to be dull
malleable
brittle
ductile
good conductor of electricity
poor conductor of electricity
Procedure:
1. In your lab notebook, draw a table like the one shown below.
2. Observe the appearance of each of the elements. Record physical state, color,
luster and other observable characteristics.
3. Using the micro-conductivity tester, determine whether the elements conduct
electricity. If you observe carefully, you might see that some are semiconductors.
4. To determine which elements are malleable, place a single piece of the element on
a paper towel, and gently tap it with a hammer. An element is brittle if it shatters
when it is hit. An element is malleable if it flattens when it is tapped.
5. To test the reactivity with 1 M HCl, label 9 wells with the symbols for each
element. Add 5 drops of the acid to each well. Then add a small sample (approx
0.1 gram) of each element to the labeled wells. Formation of bubbles of hydrogen
is evidence that a reaction is occurring. (Note: not all reactions are vigorous, so
watch closely)
Element
Copper
Silicon
Magnesium
Carbon
Nickel
Aluminum
Zinc
Sulfur
Tin
26
Appearance
Conductivity
Malleability
Reactivity
with HCl
Non-metal
Metal or
Metalloid
6.
For the following elements, try to decide if they are a metal, non-metal or
metalloid based on their appearance.
Element
Appearance Non-metal
Metal or
Metalloid
Oxygen
Lead
Bismuth
Silver
Nitrogen
Antimony
Hydrogen
7.
On the blank periodic table, label the elements that we tested in lab. From your
observations label them as metal, non-metal or metalloid. Color each group (metals,
non-metals and metalloids) a different color.
Post-lab:
Answer the following questions in your lab notebook. Use complete sentences.
1.
2.
3.
4.
5.
Which elements displayed characteristics of metals?
Where are the metals located on the periodic table?
Which elements displayed characteristics of non-metals?
Where are the non-metals located on the periodic table?
Which elements displayed some characteristics of metals and some of nonmetals?
6. Do metallic characteristics of elements seem to increase from left to right or right
to left?
7. Do metallic characteristics seem to increase from top to bottom or bottom to top?
Conclusion:
Don’t forget to write at least 2-3 sentences to conclude your lab!
27
Lab #13 - Creating ionic compounds
Ionic compounds are created when positively charged ions (cations) and
negatively charged ions (anions) are attracted to each other. While the ions themselves
are charged particles, compounds composed of ions are electrically neutral. The ionic
compounds are electrically neutral because the total number of positive charges in the
compound are equal to the total number of negative charges. Every ionic compound is
composed of both cations and anions. You will notice that most cations are metal atoms
which have lost at least one electron while most anions are non-metals which have gained
at least one electron.
Ions from which these compounds can be composed are either monatomic,
containing only one atom, or polyatomic, containing more than one atom. The charge of
an ion is the charge on the whole ion, regardless of whether the ion is monatomic or
polyatomic. In this activity, we will be using models of ions to create and name ionic
compounds.
Writing formulas for ionic compounds is not difficult, but it does require practice.
The number of cations needed to make the compound neutral is written as a subscript
after the symbol for that cation. (for example, if you need 2 sodium ions, you would
write Na2) The number of anions needed to make the compound neutral is written as a
subscript after the symbol for that anion (2 sulfur = S2 or 2 phosphate = (PO4)2).
Naming ionic compounds is fairly simple. The cation is always written first in the
formula, and is therefore always named first. The cation name is the same as the name of
the element (if it’s monatomic), or the name of the polyatomic ion (if it’s a polyatomic
ion). If the cation can have more than one charge, as with iron, a roman numeral is
written after the cation name to show the charge of the cation. For example, Fe2+ would
be written iron (II), and Fe3+ would be written iron (III).
The anion is named by dropping the end of the element name and changing it to
“-ide”. For example, when lithium is combined with fluorine, you get the compound LiF.
The name for this compound would be lithium fluoride. (Li has a +1 charge, and F has a
-1 charge)
Pre-lab:
1.
2.
3.
4.
5.
6.
7.
28
What is an ion?
What is a cation?
What is an anion?
How do monatomic and polyatomic ions differ?
How do you change the name of an element to make it an anion name?
Copy the data table from the procedure into your lab notebook.
List the names of the polyatomic ions found in the table below (you may use
your book)
Procedure:
Use the ion models to determine the formulas and names for compounds composed of the
ions listed in the table below:
#
1
2
3
4
5
6
7
Cation
Mg2+
Na+
Fe2+
Al3+
NH4+
Fe3+
K+
Anion
ClO2BrPO43NO3SO42S2-
Formula
Compound Name
Post-lab:
1.
Which of the ions above have more than one possible charge?
2.
Using your knowledge from the lab, what would the formula of a compound
composed of copper (I) and sulfur look like? What would the name of this
compound be?
3.
Combine iron (II) and phosphate. Write the formula and name.
4.
Why do we use a roman numeral in some names of ionic compounds?
Conclusion:
Don’t forget to write at least 2-3 sentences to conclude your lab!
29
Lab #14 – Formula of an Ionic Compound (Flinn lab)
30
Lab #15 – Molecular structure and VSEPR model
Pre-lab:
1.
2.
3.
4.
5.
What is a Lewis structure, and how is one determined?
What is the VSEPR model?
What is meant by polarity and how do polar and nonpolar bonds differ?
What are the basic shapes which a molecule can have?
Set up a table for your observations in your lab notebook (see the one in the
lab procedure for a guide)
Introduction:
A single covalent bond is formed when tow atoms share a pair of electrons. Each
atom provides one of the electrons of the pair. If the two atoms are alike, the bond is said
to be nonpolar covalent. If the atoms are not alike, one exerts a greater attractive force on
the electrons, and the bond is polar covalent. More than one pair of electrons may be
shared, resulting in a multiple bond.
A group of atoms held together by covalent bonds is called a molecule.
Molecules can be either polar or nonpolar. If the bonds are nonpolar, the molecule is
nonpolar. If bonds are polar, molecules can still be nonpolar if the charge distribution
throughout the molecule is symmetrical. A molecule’s symmetry depends on its shape,
that is, the positions in space of the atoms making up the molecule. Some possible
shapes are linear, angular (bent), pyramidal, and tetrahedral. Other possible shapes can
be found in chapter 6.
Although we represent molecules on paper as being two-dimensional for
convenience, they are actually three-dimensional. By building molecular models,
chemists come to understand the bonding, shapes, and polarity of even the most complex
molecules.
Procedure:
1.
Obtain a molecular model building set. Study the color code identifying the
different kinds of atoms.
2.
Observe that the following atoms have one hole (bonding site): hydrogen,
fluorine, chlorine, bromine and iodine. The atoms with two holes are oxygen
and sulfur. A nitrogen atom has three holes, and a carbon atom has four holes.
The shorter “bonds” are for single bonds while the longer, more flexible
“bonds” are for multiple bonds.
3.
Construct models of the molecules listed in your data table, checking with
your teacher along the way.
4.
Record your observations.
31
Post-lab:
1.
Which molecules were non-polar because all bonds were non-polar?
2.
Which molecules had polar covalent bonds but were non-polar because of
symmetry?
3.
Which two shapes appeared to produce polar molecules?
4.
Name two types of substances that do not contain molecules with covalent
bonds.
Conclusion:
Data table example:
Name
Formula
Hydrogen
Water
Methane
Chlorine
Ammonia
Hydrogen
fluoride
Ethyne
Dichloromethane
Nitrogen
Carbon dioxide
Methanol
Hydrogen
peroxide
Oxygen
Hydrogen
sulfide
Ethanol
32
Structure
Shape
Polarity
Lab # 16 - Paper Chromatography
Introduction:
Chromatography is used to separate mixtures of substances into their components.
All forms of chromatography work on the same principle. They all have a stationary
phase (a solid, or a liquid supported on a solid) and a mobile phase (a liquid or a gas).
The mobile phase flows through the stationary phase and carries the components
of the mixture with it. Different components travel at different rates. We'll look at the
reasons for this further down the page.
In paper chromatography, the stationary phase is a very uniform absorbent paper.
The mobile phase is a suitable liquid solvent or mixture of solvents.
Rf values: Some compounds in a mixture travel almost as far as the solvent does;
some stay much closer to the base line. The distance traveled relative to the solvent is a
constant for a particular compound as long as you keep everything else constant - the
type of paper and the exact composition of the solvent, for example.
The distance traveled relative to the solvent is called the Rf value. For each
compound it can be worked out using the formula:
Rf = distance traveled by compound
distance traveled by solvent
For example, if one component of a mixture traveled 9.6 cm from the base line while the
solvent had traveled 12.0 cm, then the Rf value for that component is: 9.6/ 12 = 0.80
Pre-lab:
1.
What is chromatography?
2.
What is a stationary phase? (define and give an example)
3.
What is a mobile phase? (define and give an example)
4.
What is meant by an Rf value and how is it calculated?
5.
When drawing lines on your coffee filter to show where your dyes began, why
should you use pencil rather than ink?
Procedure:
1.
Coffee filters usually are round, but it's easier to compare your results if the paper
is square. So, your first task is to cut the coffee filter into a square. Measure and cut an
8x8 cm square from a coffee filter.
2.
Using a pencil (ink from a pen would run, so pencil is better), draw a line 1 cm
from the edge of one side of the paper.
3.
Make six pencil dots (or however many colors of candy/markers you have) along
this line, about 0.5 cm apart.
4.
Underneath each dot, label the color of the candy/ marker you will test on that
spot. You won't have space to write the whole color name. Try BMM for blue m & m,
GS for green skittle, or something equally easy.
5.
Space 6 drops of water (or however many colors you are testing) equally distant
on a plate or piece of foil.
33
6.
Position one candy of each color on the drops. Give the color about a minute to
come off into the water. Pick up the candy and throw it away.
7.
Dip a toothpick into a color and dab the color onto the pencil dot for that color.
Use a clean toothpick for each color. Try to keep each dot as small as possible.
8.
Allow the filter paper to dry, then go back and add more color to each dot, a total
of three times, so you have lots of pigment in each sample.
9.
When the paper is dry, fold it in half with the color sample dots on the bottom.
Ultimately, you are going to stand this paper up in a salt solution (with the liquid level
lower than the dots) and capillary action is going to draw the liquid up the paper, through
the dots, and toward the upper edge of the paper. The pigments will become separated as
the liquid moves.
10.
Pour salt solution into a clean tall glass so that the liquid level is 0.5 cm. (You
want the level to be below the sample dots.) Check this by holding the paper up against
the outside of the glass. Pour out a little salt solution if the level is too high.
11.
Once the level is correct, stand the filter paper inside the glass, with the dot side
down and the edge of the paper wetted by the salt solution. [Capillary action will draw
the salt solution up the paper. As it passes through the dots, it will begin to separate the
dyes. You will notice some candy colors contain more than one dye. The dyes separate
because some dyes are more likely to stick to the paper, while other dyes have a higher
affinity for the salt water. In paper chromatography, the paper is called the 'stationary
phase' and the liquid (salt water) is called the 'mobile phase'.]
12.
When the salt water is 0.5 cm from the top edge of the paper, remove it from the
glass and place it on a clean, flat surface to dry. (before it dries, record the distance that
both the dye and the solvent moved for at least one of your candies/ markers so that you
can complete post-lab question number 6!)
13.
Prepare a second filter using food dyes. Follow the procedure steps 1-4.
14.
Place one drop of each food dye on the corresponding dot and allow the filter
paper to dry before following steps 9-12.
Post-lab:
1.
When the coffee filters are dry, compare the results of chromatography for the
different colors.
a.
Which candies/ markers contained the same dyes? (These are the candies
that have corresponding bands of color.)
b.
Which candies contained multiple dyes? (These are the candies that had
more than one band of color.)
2.
Can you match any of the colors with the names of the dyes listed on the
ingredients for the candies? (compare your first chromatogram with that of the food
dyes)
3.
Were you surprised at any of the separations you saw? Explain.
4.
Do you think this same procedure would work for permanent markers? Why or
why not?
5.
Tape the chromatograms into your notebook.
6.
Determine the Rf value for one of the dyes. Show your work!
Conclusion:
34
Lab # 17 – Mole Relationships in a Decomposition Reaction
Introduction:
One of the most important applications of the mole concept is for expressing mole
relationships between substances in a chemical reaction. To illustrate this, let us consider
the reaction between hydrogen and oxygen forming water:
2H2(g) + O2(g)  2H2O(l)
In this equation, the formulas of all participants represent molecules. The coefficients
represent the relative numbers of each kind of molecule involved in the reaction. The
number of molecules is directly proportional to the number of moles. Therefore, we can
say that 2 moles of hydrogen molecules react with 1 mole of oxygen molecules to form 2
moles of water molecules. Each mole may be expressed in terms of a gram-molecular
mass. Converting the mole ratio into gram ratios shows that 4.0 g of hydrogen react with
32.0 g of oxygen and form 36.0 g of water. This discussion is summarized in the
equations below:
Equation:
2H2(g) +
O2(g) 
2H2O(l)
Moles:
2
1
2
23
23
Molecules:
2(6.02 x 10 )
1(6.02 x 10 )
2(6.02 x 1023)
Grams:
4.0
32.0
36.0
23
23
Atoms
4(6.02 x 10 )
2(6.02 x 10 )
6(6.02 x 1023)
It can be seen that the total mass of product is the same as the sum of the mass of
the reactants. Also, the total number of all reactant atoms is equal to the total number of
product atoms, in agreement with the Law of Conservation of Matter. Note that the total
number of moles or total number of molecules of product do not necessarily equal those
of the reactants. This is because atoms rearrange themselves into different combinations
during a reaction.
This experiment is designed to illustrate mole relationships in chemical reactions.
You will thermally decompose sodium bicarbonate (baking soda). The products are
sodium carbonate, water and carbon dioxide. Decomposition refers to the breaking down
of a compound into smaller compounds or even individual elements. Thermal
decomposition refers to the breaking down of a compound by using heat. From the
experimentally determined masses of the sodium bicarbonate, you will (or should) be
able to determine the mole ratio and the coefficients for a balanced equation.
Pre-lab:
1.
2.
3.
4.
5.
6.
35
What is a mole?
What is a coefficient?
Write formulas for the following compounds: sodium bicarbonate, sodium
carbonate, water and carbon dioxide.
Calculate the molar masses of each of the above compounds.
Define thermal decomposition.
Set up an appropriate data table. (hint: anywhere in the procedure where you
see the word “record” will need a spot in a data table)
Procedure:
1.
Weigh a clean, dry crucible to the nearest 0.01 g and record
2.
Place 2 to 3 g of sodium bicarbonate in the crucible and weigh the crucible
and contents to the nearest 0.01 g. Record.
3.
Pace the crucible on a clay triangle which rests on an iron ring. (see figure A).
Heat gently for 5 to 6 minutes. Increase the intensity of the flame and heat
strongly for another 3 to 4 minutes.
4.
5.
6.
Cool Crucible to room temperature and weigh. Record.
Reheat the sample strongly for another 5 minutes.
Cool, reweigh and record. If this mass is within +/- 0.03 g, the sample has
completely decomposed and you may begin your calculations. If not, repeat
step 4 until the masses are precise.
Follow-up:
Carbonates, bicarbonates, sulfites and bisulfites are examples of thermally
unstable compounds. Carbonates decompose to yield carbon dioxide and a simpler
compound as seen in the equation below:
(heat)
CaCO3(s)

CaO(s) + CO2(g)
Post-lab:
1.
Write an unbalanced equation to represent the decomposition of sodium
bicarbonate to sodium carbonate, carbon dioxide, and water vapor.
2.
Determine the number of moles of sodium bicarbonate that were used in your
experiment.
3.
a. Calculate the number of moles of sodium carbonate formed in your
experiment. (hint: you’ll first need to determine the mass of sodium carbonate
formed)
b. Calculate the mole ratio of sodium bicarbonate to sodium carbonate. (you
will need to use your answers from #2 and #3a to determine this ratio)
36
4.
5.
6.
c. Write an equation for the decomposition reaction which is balanced
according to your experiment (remember that the coefficients represent the
mole ratio). It is okay if your equation is not “correct”.
a. Write the correctly balanced decomposition equation for sodium
bicarbonate.
b. Use the correctly balanced equation to calculate the mass of sodium
carbonate that should have been obtained by decomposing the original sodium
bicarbonate.
c. Calculate the percent error between the experimental and calculated
values. ( |calculated – experimental| / calculated = percent error)
List the sources of error and indicate how the error affects your results.
Write an equation for the thermal decomposition of BaCO3.
Conclusion:
Don’t forget to write at least 2-3 sentences to conclude your lab!
37
Lab #18 – Isotopic Pennies/ Relative abundance of “pennium” isotopes
Isotopes are atoms of the same element which differ in their numbers of neutrons. The neutrons
are contained in the nucleus of the atom along with protons. Because the two isotopes are the
same element, by definition, their nuclei must contain the same number of protons. Only the
neutrons differ. Protons and neutrons, each with a mass of approximately 1 amu, are much more
massive than electrons. The mass number of the atom is defined as the number of protons + the
number of neutrons. Electrons are not accounted for because they have so little mass.
mass number = number of protons + number of neutrons
Because each proton or neutron has a mass of 1 amu, the mass of the atom, in amu, is
approximately equal to the mass number. Isotopes must have different mass numbers and
therefore also different masses because of the different numbers of neutrons. For example, there
are three different isotopes of carbon atoms: carbon-12, carbon-13, and carbon-14. Each type
has 6 protons, but the numbers of neutrons are 6, 7, and 8 respectively. Their masses also differ,
and are approximately 12 amu, 13 amu, and 14 amu. Pennies from before 1982 and after 1982
are like isotopes... they are both worth 1 cent, but they have different masses (do you know
why?).
It is important to understand that despite the difference in mass, all isotopes of a single element
are chemically alike! The number of protons and electrons are the same, and it is the number of
electrons that primarily determine the chemical behavior of the elements.
When the atomic mass of an element is reported, it is actually an average of all the isotopes which
exist for that element. Similarly, a mixture of pre-1982 and post-1982 pennies must have an
average mass in between the mass of either type of penny. Exactly what the mass is will depend
on the relative amounts of the isotopes. Some isotopes are more common in nature and are said
to have a higher “natural abundance”. When the amount of one isotope is expressed as a percent
of the total atoms of that type, it is known as a “percent abundance.” In this lab you will work
“backwards.” You will measure the mass of a sample of a mixture of ten of the pennies, but you
will not be able to look at them. From this mass, you will estimate and/or calculate the
abundance (how many) of each penny is in your sample.
You will use this lab to become familiar with the ideas of isotopes and relative abundance and
to understand how an average atomic mass can be used to determine relative abundance.
Pre-lab: Read the sections in your text on these topics and use them to help you in the lab.
1. Isotopes of the same element have the same what? They have different what?
2. The atomic mass of an isotope is measured in amu and is nearly equal to the mass number.
Look up the three isotopes of hydrogen and sketch each of them in your lab notebook. Also fill in
the isotope symbol name, the “word” name, and the approximate mass (with units) of each of
these isotopes.
3. The atomic mass of an element can be found in the periodic table under the letter symbol for
each element. Look at some atomic masses. Record the atomic mass of: Boron, Carbon and
Nitrogen.
38
Notice that these are not whole numbers, as mass numbers are. The atomic mass of an element is
the weighted average of the masses of the isotopes of that element. A weighted average reflects
both the mass and the abundance of the isotopes. For example, the atomic mass of H is 1.0079
AMU. This means that on average, the mass of each H atom is 1.0079.
4. Note that the average mass of hydrogen (1.0079 amu) is not close to 2 amu, even though the
three isotopes hve masses of about 1,2, and 3 amu respectively. Based on this information, which
isotope must be most abundant (plentiful)? Explain.
5. Set up an appropriate data table.
Procedure: Begin by massing individual samples of both pre-1982 and post-1982 pennies as
directed. Return used pennies to the bin and randomly choose new pennies for each trial or
measurement.
1. Find the mass of one pre-1982 penny
2. Find the mass of one post-1982 penny
3. Find the mass of six (6) pre-1982 _______; divide this by six (6)________________
4. Find the mass of six (6) post-1982 _______; divide this by six (6)________________
5. Using the model that pennies are atoms of the element pennium, which isotope must have more
neutrons (pre- or post-1982)? Explain.
6. Which isotope must have more protons? Explain.
7. Obtain an empty container. Find its mass:
Bring the container to your teacher to get a 10 atom sample of the element pennium. Pennium is,
of course a mixture of two different isotopes of pennies. Without looking, you will...
- determine which isotope is more abundant in questions 8-10. (Do not answer those
questions here!)
- calculate the exact percent abundance of each isotope in questions 11-15.
8. Do not open the container. Find its mass with the pennies inside: ______________
Return the unopened container to your teacher
9. Find the mass of the 10 pennies by difference and record.
Post-lab:
1. What is the average mass of a penny in your ten-penny sample? (show your calculation) This
represents the “atomic mass” of pennium:
2. Looking at your “atomic mass”, do you have more pre-1982 isotopes, more post-1982, or about
the same? Make a guess at the number of post-1982 pennies in your sample. Explain fully.
The exact percent abundance of each isotope can be calculated using the following equation:
(Note that in this equation, the percent must be expressed as a decimal or fraction, not a percent,
and is therefore sometimes called a fraction abundance... ex. 65% (percent) = 0.65 (decimal
equivalent of fraction))
Equation for processing (use in step 13): [(Mass, isotope 1)•(fraction abundance, isotope 1) +
(mass, isotope 2)• (fraction abundance, isotope 2) = Atomic Mass]
39
3. If the % abundance is unknown (as it is here), “x” may be substituted for the fraction of one
isotope. We can also substitute (1 - x) for the fraction of the second isotope. Do NOT
substitute here! Explain (a) why we can use (1 - x) for the second isotope, and (b) can you use
x-1 instead of 1 - x ?
4. Based on the masses of pennies you took in steps 1-4, what would be the best value to plug in
for the mass of each isotope? Explain and give a specific number to use in #5 for each isotope.
5. (a) Now substitute into the equation above to algebraically calculate the percent abundances
of the two isotopes of pennium. You may not use trial and error.
6. This question is NOT about YOUR sample. (a) Noting that the various samples I handed out
in class contained exactly 10 pennies, what percent abundances are the only ones possible? (b)
What allows real samples of matter to have much more specific abundances... such as 69.17%?
7. (a) Now what is your best guess for the real percent and number of post-1982 pennies in your
sample? Explain briefly. (b) Why is it unlikely that your calculated answer in step 13 came
out “exact” to sig figs? (Hint: is there a difference between pennies and atoms?)
Conclusion:
Extra credit: (No trial and error allowed.. show algebra) Chromium exists in the following
isotopes: 49.941 amu (4.41% abundant), 51.9405 amu (83.46% abundant), 52.9407 amu, and
53.9389 amu. What are the abundances of the last two isotopes?
40
Lab #19 - Jelly Bean Equations
The Law of Conservation of Matter states that matter cannot be created
nor destroyed. This applies to chemical equations. The number of atoms of
each type of element must be equal on both sides of an equation.
Materials Needed:
Reaction Grid
4 different colors of jelly beans (16,12,8,& 4 of the respective colors)
Pre-lab:
1. What is the law of conservation of matter, and how does it apply to
balancing chemical equations?
2. What is the difference between a coefficient and a subscript?
Directions…Balance the equations on your lab sheet by doing the
following:
1. Write down a color of jelly bean to represent an atom of each element in
the equation. Use the color that you have the most of to represent
oxygen.
2. Create the molecules in the equation using the jelly beans for each
element in the compound.
*Be sure to put the correct number of atoms of each
element on! For example: Ca(OH)2 equals 1 Ca, 2 O’s, and 2
H’s.
3. Place each molecule in the correct box on the “Reaction Grid.” You
may not have to use all four boxes on the grid! Once you have the
original equation laid out on the grid, you can only add more molecules
like you have with the original. Adding single jelly beans, or taking
them away, is the same as changing a subscript, and you CANNOT
EVER change a subscript.
4. Balance each equation by adding additional molecules until all types of
atoms are equal. For example: #of orange reactants must equal #
of orange products, etc.
5. Write down the balanced equation in your lab notebook. Use the
table provided as an example. The number of molecules in each box
(number of toothpicks of each molecule) is equal to the coefficient
for that compound.
41
EQUATIONS:
1. Al + O2  Al2O3
Balanced
Equation:______________________________
Al: ____________
Write in
O: ____________
the color
of the jelly
bean here!
2. HgO  Hg + O2
Balanced
Equation:______________________________
Hg:____________
O:_____________
3. NaOH + H2SO4  Na2SO4 + H2O
______________________________
Na:__________
H:_____________
O:___________
S: ____________
4. Fe + O2  Fe2O3
______________________________________
Fe:_____________
O: _____________
5. H2 + O2  H2O
______________________________________
H:____________
O:____________
6. Fe + CuCl2  FeCl2 + Cu
_________________________________
Fe:____________
Cu:____________
Cl:_____________
7. Mg + HCl  H2 + MgCl2
_________________________________
Mg:_____________
H: ______________
Cl:______________
8. H2O + Fe  Fe2O3 + H2
_________________________________
H:____________
O:____________
Fe:____________
42
9. HgO + Cl2  HgCl + O2
_________________________________
Hg:____________
O:_____________
Cl:_____________
10. Ca(OH)2 + HNO3  Ca(NO3)2 + H20
____________________________
Ca:_____________
H:_______________
O:_____________
N:_______________
Post-lab:
1.
What is meant by a balanced equation?
2.
How does the Law of Conservation of matter apply to chemical
equations?
3.
How do you know that the equations you balanced are truly
balanced?
4.
Why weren’t you allowed to just add or remove single jelly beans?
5.
On your reaction grid, you had jelly bean “molecules” to represent
your reactants and separate jelly bean “molecules” to represent your
products. In a real chemical reaction, would the atoms (jelly beans) be
different on each side, or would they be the same? Explain.
6.
The number of “molecules” in each box of your reaction grid was
equal to the _______________________ for that compound.
Conclusion: 2-3 sentences explaining what you learned through this
activity.
43
Lab #20 – Precipitation Reactions and Net ionic Equations
The majority of ionic solids like salts (KCl and NaCl) are soluble in water
because the polar water molecules surround the individual ions of the salt. Those that do
not dissolve and go into solution form solid products called precipitates. These
precipitates have many colors and often help scientists identify what the precipitate is
present. In meteorology, the term precipitation refers to meteorological phenomena
such as rain or snow. Precipitation in chemical solutions occurs when two chemicals react
to form a product that is insoluble in water and falls out of solution like rain or snow. A
precipitate is a solid substance that separates from solution during a chemical reaction. A
precipitate can be identified by the cloudy, milky, gelatinous, or grainy appearance it
gives to the mixture.
A barium sulfate precipitate can be produced by the reaction of barium chloride and
sodium sulfate. A chemical equation to describe the reaction is written and balanced like
this:
BaCl2 (aq) + Na2SO4 (aq)  2 NaCl (aq) + BaSO4(s)
Barium sulfate, BaSO4 (s), is a common precipitate used as an X-ray contrast medium
because it is insoluble in water an opaque to x-rays. Typically a patient drinks an aqueous
slurry of barium sulfate just before he is x-rayed. The precipitate coats his stomach and
intestines. These organs then show up on the x-ray film in vivid contrast, aiding the
doctor’s diagnosis. Notice the reaction that forms BaSO4 is a double-replacement reaction
in which the cations and anions of the reactants trade partners to form the products.
To write a total ionic equation, rewrite all aqueous substances as their component ions
and keep all solid substances unchanged.
Ba+2(aq) + 2Cl- (aq) + 2Na+(aq) + SO4-2(aq)  2Na+(aq) + 2Cl-(aq) +
BaSO4(s)
Cancel all components of the reaction that are identical on both the reactant and product
side of the reaction. These are called spectator ions. Write the net ionic equation by
rewriting the reaction equation including only those reactants and products that change in
the reaction.
Ba+2(aq) + SO4-2(aq)  BaSO4(s)
44
Always check to see that your net ionic equation is balanced.
When doing these reactions involving ionic compounds remember that they are made of
positive and negative ions held together by the attractive, electrostatic forces that occur
between oppositely charged particles. In water, soluble ionic compounds break apart
completely into their respective ions. Example: NaCl (s) when put into water yields Na+
Ag+(aq)
(aq) and Cl (aq), and AgNO3 also dissociates in water to form these respective ions
and N03-(aq). It turns out that when these two solutions of sodium chloride and silver
nitrate are mixed a solid falls out and precipitation occurs. The new mixture still contains
Na+(aq) and N03- (aq), and the newly formed precipitate is AgCl (s). How do we know that
AgCl forms a solid precipitate? Remember to reference the solubility pyramid I use that
helps you decide what will and won’t go into solution (fig 1) .
Fig 1
In this experiment, you will mix 7 different ionic solutions in all possible combinations
to determine which combinations result in precipitate formation. Based on your results,
you will write complete ionic equations and net ionic equations for each reaction that has
takes place. Use your solubility chart to determine which compounds yield
precipitates. While the lab focuses on the solubility of different compounds, my purpose
in having you do this activity is to give you more practice looking at chemical reactions,
writing formulas, naming inorganic compounds, and describing chemical changes with
equations.
Pre-lab:
1.
2.
3.
4.
5.
45
In chemistry, how is precipitate defined?
What characteristics can help you identify a precipitate?
What is a spectator ion?
What is a net ionic equation? How does it differ from a complete ionic equation?
Set up an appropriate data table. (see figure 2 below)
Procedure: In this experiment, you will mix 7 different ionic solutions in a well plate
in all 49 possible combinations to determine which combinations result in precipitate
formation. Based on your results, you will write complete and net ionic equations for
each reaction that creates a precipitate. Use the solubility chart that is represented by Fig
1 above to compare your lab results with the expected results. You will be combining
solutions containing the dissolved ionic substances below. In each test, you will need to
add 2 drops (Don’t make a Mess) of the first solution to a well plate. Then add 2
drops (Don’t make a Mess) of the second solution a drop at a time. When/if you see
the formation of a precipitate, you have witnessed a reaction and will need to write a
balanced chemical equation for that reaction. You will be required to (a) describe the
precipitate (b) find the correct formula for the ionic products, (c) write and balance the
entire equation for each reaction you witness and (d) write and balance the net ionic
equation.
Chart out your chemical tests in an organized table similar to the fig 2 and then write the
reactions underneath.
Solutions you will have available:
sodium sulfate
Na2SO4
sodium chloride
NaCl
silver nitrate
AgNO3
sodium iodide
NaI
lead nitrate
Pb(NO3)2
ammonium carbonate (NH4)2CO3
calcium chloride
CaCl2
Presenting your data: Make a table which has columns and rows headed by the
formulas of each of the solutions to be tested. As you mix chemicals, you will report the
results from each of those combinations in the “box” on the table which lies at the
intersection of the column and row labeled with those two substances. If a precipitate
forms, describe its physical characteristics in the box. If no precipitate forms, mark
NR in the box.
Na2SO4
NaCl
AgNO3
NaI
Pb(NO3)2
(NH4)2CO3
CaCl2
Na2SO4 NaCl AgNO3 NaI Pb(NO3)2 (NH4)2CO3 CaCl2
x
x
x
x
x
x
x
Fig 2
46
Make sure you include in your lab a description of any precipitates and the chemical
compounds that make up the reactants, products, the full ionic equation and the net ionic
equation with (aq) and (s) designations
Post-lab:
1. For each of the wells in which a precipitate occurred;
a. Write the complete balanced equation for the reaction, being sure to
include (aq) or (s).
b. Write the complete ionic equation for the reaction.
c. Write the net ionic equation for the reaction.
2. What conclusions can be drawn from your observations? List any similarities
and differences in the physical behavior (ie color, texture) of the precipitates as
compared to their formula.
3. Some heavy ions that precipitate out of solutions are used for specific purposes
because of this insolubility. Barium sulfate is an example. It is used as a
darkening compound for x-rays, since it shows as a dark outlined area that might
help locate abnormalities in the upper and lower gastrointestinal tracts. Explain
how silver’s ability to make a precipitate when activated by light is used in our
every day life.
Conclusion:
47
Lab #21 - Stoichiometry and Gravimetric Analysis – lab 10B on p 744
Read your textbook (pp 744-747) for an introduction to stoichiometry and gravimetric
analysis.
Pre-lab:
1. Explain gravimetric analysis.
2. What are the reactants being used in this lab?
3. What are the formulas for these reactants?
4. Write a balanced double-displacement reaction for these substances.
5. How is percentage error calculated?
6. Copy the data table from p. 745 into your lab notebook.
Procedure:
Follow the procedure as outlined on pages 746-747 of the text. Instead of using a
drying oven, we will be leaving our samples on the lab table to dry overnight.
Post-lab:
1. From pre-lab question #4, determine the precipitate from the reaction.
2. Calculate the mass of the dry precipitate.
3. Calculate the moles of the dry precipitate.
4. How many moles of sodium carbonate were present in your original 15 mL
sample? (hint: use the mole-ratio from the equation)
5. There was 0.30 mol of strontium chloride in every liter of solution. Calculate the
number of moles of strontium chloride that were added. Determine whether the
strontium chloride or sodium carbonate was the limiting reactant.
6. How many grams of sodium carbonate were present in the 15 mL sample? (use
question 4)
7. Assuming you were given a 0.5 M solution of sodium carbonate, the theoretical
yield would equal 0.795 g for every 15 mL. Calculate your percentage error.
Conclusion: Don’t forget a 2-3 sentence conclusion!
48
Lab #22 - Heat of Fusion of Ice
Introduction:
When a physical or a chemical change takes place, energy is either given off or
absorbed. This energy is usually in the form of heat. When heat is given off the change
is exothermic. If heat is absorbed the change is endothermic.
It is important for chemists to be able to measure these changes in heat. Most
measurements are made with a device called a calorimeter. The technique used in
measuring heat changes is called calorimetry.
A simple calorimeter can be constructed by placing a styrofoam cup in a beaker.
Styrofoam is an excellent insulator that absorbs heat very slowly. More sophisticated
calorimeters have been developed where very precise measurements are required.
When two bodies at different temperatures come in contact with each other, heat
flows from the warmer body to the colder body. This exchange of heat will continue
until the two bodies are at the same temperature. In the calorimeter heat is exchanged
between the substance to be measured and water. By making direct measurements of the
temperature of the water it is possible to compute the heat gained by the water and
therefore the heat lost by the body being measured. The specific heat capacity of water,
the amount of energy required to raise the temperature of 1 g of water by 1oC, is equal to
4.18 J/goC and is represented by the letter “c” in the equation q = mcT.
In this experiment you will measure the amount of heat required to melt ice. The
amount of heat required to melt 1 gram of ice is called the heat of fusion. In this
experiment the heat lost by the hot water will be equal to the heat gained by the ice in the
process of melting. The following relationships will be used in this experiment:
Heat lost by hot water = heat gained by ice
Materials:
Styrofoam Cup
Ring Stand
Thermometer
Pre-lab:
1.
2.
3.
4.
5.
6.
49
Ice
Ring
Crucible Tongs
Bunsen Burner and Striker
Wire Gauze
400ml Beaker
What is the difference between an endothermic and an exothermic reaction?
What device is used to measure changes in heat?
Of what will our simple calorimeter be composed and why?
What is heat of fusion?
What is specific heat capacity, and what is that value for water?
Copy the data table into your lab notebook
Procedure:
1. Heat 300ml of water to a temperature of 50˚C. Note: This is below boiling.
2. Transfer 150ml of this heated water to the Styrofoam cup. Record this volume
(V1) on your data table.
3. Measure and record the temperature (T1) of the heated water that is in the
Styrofoam cup. Immediately add 2 to 3 ice cubes to the heated water.
4. Gently stir the ice water mixture with the thermometer. The cup should
contain ice at all times. Therefore, if the last of the ice is about to melt, add
another ice cube.
5. Monitor the temperature of the ice water mixture as you stir. Continue stirring
(and adding ice, if necessary) until the temperature evens off (no longer
drops). Record this final temperature (T2).
6. Using tongs, slowly remove the un-melted ice. Allow any water from the
tongs and ice to drip back into the cup. Measure and record the final volume
of water in the Styrofoam cup (V2).
Data Table: Remember Units!
Volume
Initial
Final
V1
V2
Temperature
T1
T2
Post-lab: SHOW WORK!!
1. Using the known density of water (1g/ml), find the mass in grams (m1) of the
original volume of water (V1).
2. Find the volume, in milliliters, of the water resulting from the melted ice.
(Vice = V2 - V1)
3. Find the mass in grams (m2) of the volume of water resulting from the melted ice.
4. Find the change in temperature of the water. (ΔT = T1 - T2 )
5. Calculate the heat lost by the original mass of water using q = mcΔT. Note: The
specific heat of water is 4.18 J/g˚C.
6. Find the heat of fusion of the ice used in this experiment. (ΔHfus = q/m2)
7. Convert the ΔHfus from J/g into kJ/mol using the molar mass of H2O.
8. Find your percent error by comparing your ΔHfus in kJ/mol to the true value of
6.01kJ/mol. %error = (true value – experimental value) / true value X 100.
9. List a minimum of three possible sources of error in this experiment.
10. How might the use of a “real” calorimeter reduce your percent error?
11. One source of error is the flow of heat between the water in the cup and the
surrounding. Explain how this error is reduced by starting with the water at about
50˚C.
12. In what way does a calorimeter make use of the law of conservation of energy?
13. Is the process of fusion exothermic or endothermic? Explain using evidence from
your experiment.
14. What is the difference between heat and temperature?
50
Lab #23 - Graphing a Heating Curve for Water
Materials:
safety glasses
stopwatch or timer
small beaker
thermometer
graph paper
lab coat or apron
hot plate
ice
plastic thermometer clamp
Introduction:
The first law of thermodynamics basically states that energy can be transformed (changed
from one form to another), but cannot be created or destroyed. This leads into the concept
of how different substances can change from one phase to another by absorbing or
releasing energy. When the system is heated, energy is transferred into it. In response to
the energy it receives, the system changes, for example by increasing its temperature. A
plot of the temperature versus time is called the heating curve.
Water is a common substance. Ice is the stable phase below 0oC. Both solids and liquids
coexist at 0oC. When heat is put into the system, more solid will melt. Thus, the
temperature does not change. The normal boiling point is 100oC. As heat is absorbed,
some water will boil off but the temperature is kept at 100oC. This change in temperature
may be observed and measured against time in an effort to visualize the heat curve for
water.
Pre-lab:
1. What is a heating curve?
2. Explain the first law of thermodynamics.
3. Copy the data table into your notebook.
4. Why do you think you should keep the thermometer bulb from touching the
bottom of the beaker?
Procedure:
1. Put on safety goggles and lab coat or apron.
2. Use the data chart provided to record time and temperature. The time column
starts with 0. The temperature column is blank. You will record the temperatures
in the temperature column during the investigation.
3. Fill the small beaker with ice. Insert the thermometer. Wait 2 minutes. Observe
and record the starting temperature (0 time) in the data table.
4. Place the beaker of ice on the hot plate. Position the thermometer in the clamp so
that the bulb of the thermometer does not touch the bottom of the beaker.
5. Turn the hot plate on high and start the timer. After 30 seconds, record the
thermometer reading without removing the thermometer from the beaker. (DO
NOT TOUCH THE HOT PLATE !!!!!!)
6. Continue to record the temperature on the chart every 30 seconds.
7. Make a note when the ice has melted and when the water begins boiling.
51
Analysis:
Write your data in the following table:
Time
Temperature
Post-lab:
1.
Prepare a graph from your data that includes the following information:
a.
Label the x-axis with the time (in minutes). This is your independent
variable. Label the y-axis as temperature (in degrees Celsius). This is your
dependent variable.
b.
Plot your points using your recorded data.
c.
Label the 5 areas on your graph: solid (S), liquid (L), gas (G), freezing
point/melting point FP/MP and condensation/boiling point (CP/BP).
d.
Trace, with colored pencils, the following parts of the line on your graph:
slowest molecular motion (in red), fastest molecular motion (in green).
e.
DON’T FORGET TO TITLE YOUR GRAPH!!!!!!!!
f.
Your graph should look like stair steps not a straight line.
2.
Explain what is happening to the water molecules in the flat areas of the line on
your graph during the phase changes from solid to liquid and liquid to gas.
3.
When the ice is melting is it releasing heat or absorbing heat? Explain your
answer.
4.
If you put the liquid water into the freezer and recorded its temperature as it
refroze, would it be absorbing heat or releasing heat? Explain your answer.
Explore Further (for extra credit).
Research what is occurring when you have a fever. What part does water play in
regulating your body temperature? What happens to the chemical bonds of enzymes
when exposed to too much heat?
52
Lab #24 – Hess’s Law and Enthalpy of Formation
Introduction:
Most chemical reactions either use heat from the environment or release heat into
the environment. The amount of heat released or absorbed when 1 mole of compound is
made from its elements under standard conditions is known as standard enthalpy of
formation, and is symbolized Hf. When heat is released, the reaction is said to be
exothermic. The heat produced in an exothermic reaction appears as one of the products,
and the standard enthalpy is negative. When heat is absorbed, the reaction is said to be
endothermic. The heat absorbed in an endothermic reaction appears as one of the
reactants, and the standard enthalpy is positive.
In this experiment, you will measure and compare the quantity of heat involved in
three chemical reactions. You will be using this information to prove Hess’s law, which
states that the overall enthalpy change in a reaction is equal to the sum of the enthalpy
changes of the individual steps in the process. Essentially, two of the reactions, when
added together, should equal the other reaction. The hope is that the sum of the heats of
formation of the first two reactions will be equal to the heat of formation of the third
reaction.
In this experiment, a coffee cup will again serve as our calorimeter. We will
assume that the enthalpy of formation is used to change the temperature of the aqueous
solution only. Any small heat losses to the surroundings will be neglected. We will also
assume that the aqueous solution has the same heat capacity as water, 1 cal/goC or
4.18 J/goC.
The reactions we will be comparing in this lab are as follows:
1.
Solid sodium hydroxide dissolving in water to form an aqueous solution of ions.
NaOH(s)  Na+(aq) + OH-(aq) + x1cal
2.
An aqueous solution of sodium hydroxide reacts with an aqueous solution of
hydrogen chloride to form water and an aqueous solution of sodium chloride.
Na+(aq) + OH-(aq) + H+(aq) + Cl-(aq)  H2O(l) + Na+(aq) + Cl-(aq) + x2cal
3.
Solid sodium hydroxide reacts with an aqueous solution of hydrogen chloride to
form water and an aqueous solution of sodium chloride.
NaOH(s) + H+(aq) + Cl-(aq)  H2O(l) + Na+(aq) + Cl- (aq) + x3cal
53
Pre-lab:
1.
Define enthalpy of formation
2.
Define Hess’s law
3.
What are the differences between an exothermic and an endothermic reaction?
4.
What specific heat capacity, c, will be used to calculate the enthalpy of formations
for the aqueous solutions?
5.
Set up a data table. You will need room for three reactions. Mass/ volume of
NaOH, volume of HCl, T1 and T2 for each of the three reactions (you will not fill in all
spaces for all reactions)
Procedure:
CAUTION: SODIUM HYDROXIDE IS EXTREMELY CORROSIVE TO THE
SKIN AND MAY CAUSE BLINDNESS IF IT GETS INTO YOUR EYES
Part I
1.
Put 200 mL of cool tap water into a coffee cup. Stir CAREFULLY with a
thermometer until a constant temperature is reached (about room temperature). Measure
this temperature as accurately as possible and record it as T1.
2.
Obtain a sample of about 4 g of sodium hydroxide and determine its mass to the
nearest 0.01 g. Record this mass. Do not handle the sodium hydroxide with your
fingers! Since sodium hydroxide becomes moist when it is in contact with the open air,
be sure to keep the lid closed when you are not using it.
3.
Pour the solid sodium hydroxide into the water in the coffee cup. Place the
thermometer into the solution and stir gently but continuously until the sodium hydroxide
is dissolved. Record the extreme temperature reached as T2. Before proceeding to part
II, discard the solution and rinse the cup thoroughly with water.
Part II
1.
Measure 100 mL of 1.0 M HCl into the coffee cup. Record the exact volume.
Measure 100 m of 1.0 M NaOH into a 250 mL beaker and record its exact volume. Both
of these solutions should be at, or slightly below, room temperature. Check this with the
thermometer and record their temperatures as T1.
2.
Add the sodium hydroxide solution to the hydrochloric acid solution in the coffee
cup. Mix quickly and record the extreme temperature reached as T2.
3.
54
Discard the solution and rinse the cup before proceeding to Part III.
Part III
1.
Repeat steps 1, 2 and 3 of part I but substitute 100 mL of 1 M HCl and 100 mL of
tap water for the 200 mL of tap water in step 1.
2.
Discard the solution and rinse the cup thoroughly.
Calculations and post-lab questions:
1.
For each part, calculate the following:
a.
the change in temperature of the water or water solution
b.
the amount of heat absorbed by the solution in joules
c.
the number of moles of NaOH used (realize that 1M NaOH has 1 mol of
NaOH per 1 L of solution)
d.
the amount of heat evolved (part b) per mole of NaOH used (part c)
2.
Express the above results as enthalpy of formation, Hf, (recall that enthalpy of
formation should be expressed as J/mol with a sign designating whether the reaction is
endothermic or exothermic)
3.
Write net ionic equations for reactions 2 and 3.
4.
In part I, H represents the heat evolved as solid NaOH dissolves. Look at the net
ionic equations for parts II and III and make similar statements as to what the enthalpy of
formation for parts II and III represent.
5.
a.
Compare H3 with (H1 + H2) and explain in terms of your answer to
question 4.
b.
Calculate the percentage difference between H3 with (H1 + H2),
assuming H3 is the accepted value.
6.
Suppose you had used 8 g of NaOH(s) in part I.
a.
how would this have affected the change in temperature?
b.
what would have been the number of joules evolved in your experiment?
c.
what effect would this have had on your calculations of H1, the heat
evolved per mole?
Conclusion
55
Lab #22 -Virtual chemistry lab for acid-base titration (Modified from
http://lrs.ed.uiuc.edu/students/mihyewon/chemlab_instruction.html)
Have you seen an acid or base solution?
If your answer is no, you must be an alien because we earth people encounter acids and
bases in everyday life. Orange juice, which we drink everyday, is a citric acid
solution. Vinegar? Acetic acid. Commercial antacid? Magnesium hydroxide (base).
Household cleaning products? Ammonia (base)...Hmmm...
What is the common property of orange juice and vinegar? Sour taste, right? Acids
have a sour taste and acid solutions are capable of dissolving certain metals such as iron
and zinc. In contrast, bases have a bitter taste and slippery feel.
What is acid-base exactly?
Acids were defined by the Swedish chemist Arrhenius as substances that, when
dissolved in water, produce hydrogen ions (H+). For example, gaseous hydrogen
chloride reacts with water to give hydrochloric acid and increases the amount of
hydrogen ion in the water.
Bases are also defined as substances that when dissolved in water, yield hydroxide ions
(OH-). For example, sodium hydroxide (NaOH) dissolves in water and increases of the
amount of hydroxide ions in the water.
cf. The symbol (g) is used to designate the gaseous state, (s) for the solid, and (aq) for the aqueous state.
56
How can we know it's an acid or base?
Tasting a lab chemical is very dangerous and you don't want to taste a liquid to
determine whether it is an acid or not. Then what would you do?
One easy way is using indicators. One familiar example is a litmus indicator. Litmus
indicators change the color from blue to red when added to acid solutions. In contrast,
base solutions change litmus indicators in color from red to blue. Many natural dyes
found in fruits, vegetables, and flowers act as indicators, too. For example, red cabbage
extract is red in acidic solution and blue in base solution.
However, the indicators cannot show which acid solution is more acidic than the other
acid solutions or which base solution is more basic than the other base solutions. Then
what would you do to measure the acidity of solutions?
Scientists use the "pH" scale to determine how acidic or basic a solution is. pH is a
measure of the hydrogen ion concentration of a solution and is defined as the equation.
Usually the pH scale goes from 0 to 14 and acids are found between 0 to 7 and bases
are from 7 to 14. The middle point of the pH scale is 7 and distilled water is exactly 7.
In other words, water is neither acidic nor basic, but neutral. The lower pH value an
acid solution has, the more acidic it is. In contrast, the higher pH value a base solution
has, the more basic it is. According to the pH values, some common acids and bases
are arranged in the following figure. Now, can you tell which one is more acidic?
lemon juice or vinegar ?
57
What would happen if an acid and a base were mixed
up?
Let's think about it this way. What would happen if a hydrogen ion (H+) from an acid
reacted with a hydroxide ion (OH-) from a base? That's right. Water would be
produced.
If the spectator ions (other than H+ and OH- ions) from the acid and the base are put
back into the equation, it would read, for example,
showing that an acid reacts with a base to yield water plus an ionic compound called a
salt. This acid-base reaction is called a neutralization reaction.
58
What is titration?
Using the neutralization reaction, you can determine the concentration of an acid or a
base solution. As shown in the previous equation, one hydrogen ion from an acid reacts
with exactly one hydroxide ion from a base to produce a water molecule. If adding 100
hydroxide ions to an acid solution makes it a neutral solution (only water and salt),
there must be 100 hydrogen ions in the original acid solution and you can calculate the
concentration of the acid solution.
To titrate an unknown acid/base solution, take a certain amount of the unknown
solution and add a standard reagent of the known concentration carefully until the
neutralization reaction is completed. This point where the number (or mole) of
hydrogen ions and hydroxide ions are equal is defined as the equivalence point. To
determine the equivalence point, scientists use an indicator or a pH meter. With the
data of volume of the standard reagent used, the concentration of the unknown solution
can be calculated. This whole process is called "titration."
Titration procedure
First, you need to choose volume of an unknown acid or
base solution and put it in an Erlenmeyer flask.
Second, fill a burette with a standard reagent of known
concentration and read the initial volume of the solution.
Of course, if you put an acid in the Erlenmeyer flask, you
need to put a base in the burette and vice versa. A burette is
a good apparatus for the determination of an equivalence
point in acid-base titration because you can accurately read
the volume of solution used.
Third, add a couple of drops of an indicator in the flask for
titration. An indicator is a soluble dye that changes its
color noticeably over a fairly short range of pH. Different
indicators show color changes at different pH values and it
is important to determine an indicator to be used according
to the expected equivalence point.
If a pH meter is available, put a pH electrode in the flask.
Fourth, slightly open the cork of the burette and add the
standard reagent into the unknown solution. Around the
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expected equivalence point of the titration, you need to drop
the solution very slowly and mix the solutions very well
because, around the equivalence point, just one drop of
solution from the burette can make a radical pH change in
the mixed solution. If the color of the solution in the
Erlenmeyer flask changes, record the volume of the solution
in the burette and add a few drops of the solution to make
sure the equivalence point you found is correct.
Finally, using the data from your acid-base titration, you can calculate the concentration
of the unknown solution. The equation for this procedure is as follows. (Here, M means
molarity (concentration) of solution, and V means volume of solution.)
For example, if you choose 10.0 mL hydrochloric acid solution as a concentrationunknown solution and it takes 10.0 mL of 0.100 M sodium hydroxide solution to titrate,
the concentration of the hydrochloric acid is 0.100 M.
If you want to use a polyprotic acid or base (which can donate two or more hydrogen or
hydroxide ions per molecule) for titration, you need to multiply the molarity of the
solution by the number of hydrogen or hydroxide ions it can donate per molecule.
Pre-lab:
1.
Define the following terms: equivalence point, titration and standard reagent.
2.
What is a burette? A pipette?
3.
How can the concentration of an unknown acid or base be determined if you
know the volume of standard base or acid with which the unknown is titrated?
4.
Write equations for the reactions of solutions of hydrochloric acid with sodium
hydroxide and vinegar (acetic acid – CH3COOH) with sodium hydroxide.
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Procedure:
Go to http://lrs.ed.uiuc.edu/students/mihyewon/chemlab_experiment.html and read
through the procedure.
Part A: You will choose a strong acid (either HCl or HNO3) and titrate it with a strong
base (NaOH)
Part B: You will choose a weak acid (acetic acid) and titrate it with a strong base
(NaOH)
Post-lab:
1.
At what point in the titration did your mixture experience the greatest pH change
in the shortest time?
2.
If you titrate 30 mL of an unknown concentration of hydrochloric acid with 33
mL of 0.5 M NaOH, what is the concentration of your unknown acid?
3.
Why is phenolphthalein such a good choice for acid-base titrations?
Conclusion:
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Lab #26 – Elasticity of Gases, Boyle’s Law
Robert Boyle (1627-1691), an Irish chemist, made several contributions to
physics and chemistry; the most notable of which is Boyle’s Law, which details the
relationship between the pressure of a gas and the volume it inhabits.
Pressure is defined as an applied force per unit area, and is given by the formula:
P = F/A. The SI unit for pressure is the pascal (Pa), although you may be more familiar
with pounds per square inch (psi), torr, or millimeters of mercury (mm Hg). There are
several different devices which can be used for measuring pressure. For example, when
measuring tire pressure, you would use a pressure gauge. A barometer is a special form
of pressure gauge used to measure atmospheric pressure.
In the experiment we will be doing, a given quantity of gas will be trapped in a
syringe. The pressure on this gas will then be increased by placing weights (books) on
top of the plunger of the syringe. The total pressure acting on the gas consists of the
weight of the books plus the weight of the atmosphere: Ptotal = Pbooks + Patm
Pre-lab:
1.
Who was Robert Boyle?
2.
Boyle determined the relationship between what two aspects of gas at the same
temperature?
3.
What is pressure?
4.
What types of devices are used to measure pressure?
5.
What are several units used for measuring pressure?
6.
Leave space for Part I observations and create a data table similar to the one
below for Part II:
Pressure (#
books)
1
2
3
4
5
6
Vol. 1 (mL)
Vol. 2 (mL)
Vol. 3 (mL)
Avg. vol.
(mL)
Procedure:
Part I – Qualitative Observations
Record all observations under Part I:
1.
Take the barrel of the syringe in one hand and the plunger in the other. Place the
plunger in the syringe and push down on the plunger (be sure the cap is on, and make
sure you don’t press too hard!) Record your observations!
2.
Place a piece of string alongside the plunger and push the plunger to the bottom of
the syringe, again making sure the cap is on the syringe. Remove the string and then
attempt to remove the plunger from the syringe. Record your observations.
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Part II – Quantitative Observations
1.
Set the plunger at the highest mark of a dry syringe by inserting the piece of string
as before. Remove the string and place the sealed tip of the syringe into the block of
wood which is cut out for it. Clamp the syringe stopper arrangement in an upright
position, as shown in the diagram.
2.
Record the mass of one textbook.
3.
Carefully center a textbook on top of the plunger. Read as precisely as possible
the volume of gas trapped in the syringe when a load of one book is on the plunger. Be
sure to record the uncertainty that you estimate is associated with each reading. Record
the pressure in books and the volume of gas in milliliters. Repeat the procedure twice by
removing the book and replacing it. Read and record the volume each time the book is
replaced.
4.
Repeat step 3 for each of the 5 remaining books. Do not worry if the syringe does
not return to the original volume when the books are removed. This is not due to air loss,
but to friction between the plunger and the syringe. Do not attempt to correct for it.
Post-lab:
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