Sanford Medical LAM AND TAN Division of Loss Worksheet See the attached questions. Answers available in chegg and course heroPut it in excel Using calibrat

Sanford Medical LAM AND TAN Division of Loss Worksheet See the attached questions. Answers available in chegg and course heroPut it in excel Using calibration curves to determine concentration
This lab introduces you to the use of spectrophotometry to measure the concentration of a
species in a solution using Beer’s Law. In the lab you will be determining the concentration
of potassium permanganate in a solution based upon its absorption. The following video
introduces the key ideas associated with the technique

When white light shines on a colored substance, some of the wavelengths that make up white
light (colors) are absorbed. The remaining wavelengths, or colors, are reflected from the
colored substance or transmitted through it. Our eyes interact with the reflected or transmitted
rays and produce our perception of the color of the substance. Consequently, the
characteristic color of a material is not the color of light that it absorbs; it is the mixture of the
reflected or transmitted colors. For example, if an object absorbs all visible wavelengths but
red, red is reflected or transmitted, and we perceive the object as red.
Transmission and absorption of light are measured with an instrument called a
spectrophotometer. When light passes through the sample cell, light with a certain
wavelength will be absorbed, while other wavelengths may be transmitted or reflected. If a
particular color is absorbed, we see the complementary color in the transmitted or reflected
light. The color we see is the complementary color of that which was absorbed.
Complementary colors are best shown on a color wheel, which is well known to art students.
A simple color wheel is shown in Figure 1.
Complementary colors are opposite each other on the color wheel. For example, if red is
absorbed and all other colors reflected, we may see the complementary color of red, which is
green.
Figure 1: Color wheel
A spectrophotometer separates white light into its component wavelengths and
measures the intensity of light at a particular wavelengths before and after it passes through a
solution.
If a solution contains a species that is colored this implies that it will absorb at a particular
wavelength According to Beer’s law, A = abc, where the absorbance (A) at a given
wavelength is proportional to the concentration of the colored material in the solution (c). In
this equation, a is the molar absorptivity, and b refers to the path length of the sample cell.
We can use a calibration curve to determine the molar absorptivity of a species. To make a
calibration curve, you will measure the absorbance in a series of samples with known
concentration. You will plot these values in Excel, to create a calibration plot.
Part 1. Calculate maximum absorbance.
It is first necessary to determine the wavelength at which the absorption is a maximum. The
absorption at this wavelength will then be used to construct a calibration plot. Go to the
simulation at
https://phet.colorado.edu/sims/html/beers-law-lab/latest/beers-law-lab_en.html
select ‘Beer’s Law Lab’ to launch the simulation
1. select potassium permanganate.
2. make sure the cell is about 1 cm wide (this is the typical width of cell)
3. use a solution that is 450 ?M, measure the absorbance across the entire wavelength range
in increments of 10 nm between 380 nm and 650 nm, recording your data in an excel
spreadsheet: column A wavelength, column B absorbance
4. Draw a ‘scatter with smooth lines and markers’ plot (see ‘intro excel’ section C) and
estimate the maximum absorbance wavelength. Appropriately format the graph and label the
graph and axis (‘intro_excel’ section D and section F)
Maximum absorbance wavelength = _________________________
Note that in the simulation it is assumed that the spectrophotometer has already been
calibrated with a blank
Part 2. Create a calibration plot
In order to determine the concentration of an unknown solution, we first construct a
calibration curve. This is done by creating five standard solutions of known concentration
and measuring the absorbance of each.
In practice the standard solutions are made by pipetting a known amount of a stock solution
with a certain amount of distilled water. Assuming the stock solution is 900. ?M, determine
the concentrations of the standards A through E (give your values to three sig figs).
Soln
ml stock
A
5.00
ml distilled
water
5.00
B
5.00
9.00
C
5.00
13.00
D
5.00
17.00
E
5.00
21.00
Show your calculations with working
standard conc
(?M)
absorbance
Measure the absorbance at each of the above at the maximum wavelength you determined in
Part 1 and record the values in Excel: column A= concentration of standard, column B =
absorbance. Then plot a scatter plot of absorbance (y) vs conc (x) (see intro_excel’ sect C).
Fit a linear line to your data and use excel to get the equation of the line (see ‘intro_excel
section E’). Appropriately label the axis and the graph.
Equation of the line
_________________________________________
Measurement of unknown concentration
An unknown is then measured to have absorption of 0.62. Use your graph/equation to
determine its concentration (give your value to three sig figs)
Unknown concentration = ______________________________________
.
Graphs to include
1. Absorption vs wavelength (Part 1)
2. Absorption vs concentration (Part II)
Acid-Base Classification of Salts – Part I
In this assignment you will be asked to classify aqueous solutions of salts as to whether they are
acidic, basic, or neutral. This is most easily done by first identifying how both the cation and anion
affect the pH of the solution and then by combining the effects. After predicting the acid-base
properties of these salts, you will then test your predictions in the laboratory.
1. State whether 0.1 M solutions of each of the following salts are acidic, basic, or neutral. Explain
your reasoning for each by writing ionic equations to describe the behavior of each salt in water:
NaCN, KNO3, NH4Cl, NaHCO3, and Na3PO4.
NaCN:
KNO3:
NH4Cl:
NaHCO3:
Na3PO4:
Once you have predicted the nature of each salt solution, you will use Virtual ChemLab to confirm
your prediction. Each solution must be approximately 0.1 M for your comparisons to be valid. Most
of the solutions in the Stockroom are approximately 0.1 M already. Three solutions must be prepared
from solid salts. One of these salt solutions is already prepared and on the lab bench ready for you to
measure the pH.
2. Start Virtual ChemLab, select Acid-Base Chemistry, and then select Acid-Base Classification of
Salts from the list of assignments. The lab will open in the Titrations laboratory.
3. On the stir plate, there will be a beaker of 0.10 M ammonium chloride (NH4Cl) that has already
been prepared. The pH meter has been calibrated and is in the beaker. Record the pH of the
NH4Cl solution in the data table on the following page. When finished, drag the beaker to the red
disposal bucket, and drag the bottle of NH4Cl to the stockroom counter.
4. Click in the Stockroom to enter. Double-click on the NH4Cl bottle to return it to the shelf and then
double-click on the NaHCO3 and KNO3 bottles to move them to the Stockroom counter. Return to
the laboratory.
5. Open the beaker drawer (click on it) and drag a beaker to the spotlight next to the Balance. Click
and drag the bottle of NaHCO3 and place it on the spot light near the balance. Click in the Balance
area to zoom in. Place a weigh paper on the balance and tare the balance. Open the bottle by
clicking on the lid (Remove Lid). Pick up the Scoop and scoop up some salt by dragging the Scoop
to the mouth of the bottle and then down the face of the bottle. Each scoop position on the face of
the bottle represents a different size scoop. Pull the scoop down from the top to the first position
(approximately 0.20 g) and drag it to the weigh paper in the balance until it snaps into place.
Releasing the scoop places the sample on the weigh paper. Record the mass in the table below.
Now drag the weigh paper from the balance to the beaker until it snaps into place and then empty
the salt into the beaker. Return to the laboratory and drag the beaker to the stir plate.
6. Drag the 25 mL graduated cylinder to the sink under the tap until it fills. When filled, it will
return to the lab bench and will indicate that it is full when you place the cursor over the cylinder.
Drag the 25 mL cylinder to the beaker on the stir plate and empty it into the beaker. Place the pH
probe in the beaker and record the pH in the data table. Drag the beaker to the red disposal bucket.
Double-click the bottle of NaHCO3 to move it to the Stockroom counter. Repeat steps 5 and 6 for
KNO3.
7. Click in the Stockroom. The stock solutions of NaCN and Na3PO4 are already approximately 0.1
M. Double-click each bottle to move them to the counter and return to the laboratory. With these
solutions you can pour a small amount into a beaker that you have placed on the stir plate and
place the pH probe in the solution. Record each pH in the data table. Drag each beaker to the red
disposal bucket when you have finished. Were your predictions correct?
Data Table
solution
NH4Cl
mass used
(g)

concentration
(M)
pH
acidic, basic or
neutral
NaHCO3
KNO3
NaCN

Na3PO4

8. For some of the salts, the concentrations were not quite 0.1 M. If you had used exactly 0.1 M
solution would this change whether they are acidic basic or neutral? Discuss
9. Use Le Chatelier’s Principle to qualitatively determine what the pH of a 0.2 M NaCN would be
compared to the pH of a 0.1 M solution of NaCN. Explain your reasoning.
10. . Use Le Chatelier’s Principle to qualitatively determine what the pH of a 0.2 M NH4Cl would be
compared to the pH of a 0.1 M solution of NH4Cl. Explain your reasoning
Ranking solutions by pH – Part II
In this assignment you will be asked to rank aqueous solutions of acids, bases and their salts in order
of increasing pH. This is most easily done by first identifying the strong acids that have the lowest
pH, the strong bases that have the highest pH, and the neutral solutions that have a pH near 7. The
weak acids will have a pH between 1 and 6 and the weak bases between 8 and 14. The exact order of
weak acids and weak bases is determined by comparing the ionization constants (Ka for the weak
acids and Kb for the weak bases. Remember: the larger the Ka of acid, the further the equilibrium
(acid dissociation) is to the right and hence the lower the pH). After ranking the pH of these solutions,
you will then test your predictions in the laboratory.
1. Arrange the following 0.1 M solutions in order of increasing pH NaCH3COO, HCl, HCN, NaOH,
NH3, NaCN, KNO3, NH4Cl, H2SO3, NaHCO3, Na3PO4 and CH3COOH. (give a brief explanation
as to why it is the position, e.g. strong acid/base, larger Ka or Kb etc):
Once you have predicted the nature of each solution, you will use Virtual ChemLab to confirm your
prediction. Each solution must be approximately 0.1 M for your comparisons to be valid. Most of the
solutions in the Stockroom are approximately 0.1 M already. Two solutions will need to be diluted
and three solutions will need to be prepared from solid salts. One of these salt solutions is already
prepared and on the lab bench ready for you to measure the pH.
2. Start Virtual ChemLab, select Acid-Base Chemistry, and then select Ranking Salt Solutions by pH
from the list of assignments. The lab will open in the Titrations laboratory.
3. On the stir plate, there will be a beaker of 0.10 M ammonium chloride (NH4Cl) that has already
been prepared. The pH meter has been calibrated and is in the beaker. Record the pH of the
NH4Cl solution in the data table on the following page. When finished, drag the beaker to the red
disposal bucket, and drag the bottle of NH4Cl to the stockroom counter.
4. Click in the Stockroom to enter. Double-click on the NH4Cl bottle to return it to the shelf and then
double-click on the NaHCO3 and KNO3 bottles to move them to the Stockroom counter. Return to
the laboratory
5. Open the beaker drawer (click on it) and drag a beaker to the spotlight next to the Balance. Click
and drag the bottle of NaHCO3 and place on the spot light near the balance. Click in the Balance
area to zoom in. Place a weigh paper on the balance and tare the balance. Open the bottle by
clicking on the lid (Remove Lid). Pick up the Scoop and scoop up some salt by dragging the Scoop
to the bottle and then down the face of the bottle. Each scoop position on the face of the bottle
represents a different size scoop. Pull the scoop down from the top to the second position
(approximately 0.20 g) and drag it to the weigh paper in the balance until it snaps into place.
Releasing the scoop places the sample on the weigh paper. Now drag the weigh paper from the
balance to the beaker until it snaps into place and then empty the salt into the beaker. Return to
the laboratory and drag the beaker to the stir plate.
6. Drag the 25 mL graduated cylinder to the sink under the tap until it fills. When filled, it will
return to the lab bench and will indicate that it is full when you place the cursor over the cylinder.
Drag the 25 mL cylinder to the beaker on the stir plate and empty it into the beaker. Place the pH
probe in the beaker and record the pH in the data table. Drag the beaker to the red disposal bucket.
Double-click the bottle of NaHCO3 to move it to the Stockroom counter. Repeat steps 5 and 6 for
KNO3.
7. Click in the Stockroom. Double-click on the bottle of NH3 to move it from the shelf to the counter
and return to the laboratory. Drag the bottle of NH3 to one of the three spotlights on the lab bench.
Place a beaker from the drawer on the stir plate. Drag the bottle of NH3 to the 5 mL graduated
cylinder (the smallest one) by the sink and fill the cylinder by dropping the bottle on the cylinder.
Now drag the 5 mL graduated cylinder to the beaker on the stir plate and add the 5 mL of NH3.
Add 20 mL water to the beaker by filling and emptying the 10 mL cylinder into the beaker twice.
Place the pH probe in the beaker and record the pH in the data table. Drag the beaker to the red
disposal bucket. Double-click on the NH3 bottle to move it back to the counter.
9. Each of the other solutions is already approximately 0.1 M. With these solutions you can pour a
small amount into the beaker that you have placed on the stir plate and place the pH probe in the
solution to measure the pH. Record the pH of each solution in the data table. Drag each beaker to
the red disposal bucket when you have finished. You must determine the pH for HCl, H2SO3,
CH3COOH (HAc), HCN, NaOH, NaCN, Na3PO4, and NaCH3COO (NaAc). You may take two
bottles at a time from the stockroom.
Data Table
Solution
NH4Cl
pH
Solution
CH3COOH (HAc)
NaHCO3
HCN
KNO3
NaOH
NH3
NaCN
H2SO3
NaCH3COO (NaAc)
HCl
Na3PO4
pH
Useful information
In the case of the diprotic species (e.g. H2SO3 the major contribution to acidity comes from the first
proton loss. The second typically a makes a much smaller contribution)
H2SO3: This is a weak acid with Ka1 = 1.7 x 10-2. (This is the dissociation constant associated
with the loss of the first proton)
CH3COOH: This is a weak acid with Ka = 1.8 x 10-5.
HCN: This is a weak acid with Ka = 5.8 x 10-10.
Hydrogen carbonate is a weak base with Kb = 2.3 x 10-8.
NH3: This is a weak base with Kb = 1.8 x 10-5.
Na3PO4: The phosphate ion is a weak base with Kb = 2.2 x 10-2.
Graphs
You are not required to include any graphs with this worksheet
Question 1
A 0.16 M solution of Naf (assume the Ka of HF is 8.6 x 10-4) has a higher pH than a 0.16 M solution of NaNO2 (assume the Ka of HNO2 is 7.2 x 10-4)
True
O False
Question 1
An absorption vs wave length plot shows that a given species has a maximum absorption at wavelength of 450 nm.
If a transmission vs wave length plot was made for the same species, the maximum transmission would also be at 450 nm (true or false?)
True
False

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