Please review lab content for accuracy and provide detailed calculation
solution to 5, 6, & 7.
Titration Lab Background:
The minimum daily requirement for vitamin C is 30 mg, and the
recommended daily allowance is 60-70 mg. You can find the vitamin C
content on the labels of most commercially available juices. How close
is their reported number to the actual amount of the vitamin in the
juice? You will test that in this lab by titrating the juice with an
iodine solution, using starch as an indicator.
The formula for ascorbic acid is C6H8O6 and the structures for the
reduced form and for the oxidized form (dehydroascorbic acid) are shown
below (as well as possible given limitation of typing characters), which
illustrates the first half of the redox reaction.
OH OH O
| | / \
H - C - C - C C = O
| | | \ /
H H H C=C
| |
HO OH
Ascorbic Acid
C6H8O6
(Reduced Form)
---> 2H+ plus 2e- plus
OH OH O
| | / \
H - C - C - C C = O
| | | \ /
H H H C-C
|| ||
O O
Dehydroascorbic Acid
C6H6O6
(Oxidized Form)
The amount of ascorbic acid can be determined by carrying out a redox
titration with a known concentration of iodine in solution. In this
procedure the iodine is reduced by the ascorbic acid to form iodide, as
shown in the other half of this redox equation.
I2 + 2e- --> 2I-
The titration end point is reached when slightly more moles of iodine
have been added to the sample being tested than the number of moles of
ascorbic acid in the sample (leaving some "free" iodine molecules
present in the sample). To highlight when the excess iodine is present,
thyodene "starch" indicator is added to the titrant. Thyodene reacts
with the excess iodine to form a bright blue complex. Thyodene does not
form this complex with iodide.
Before doing a titration on a solution of unknown concentration (such as
orange juice), you must first prepare a standardized iodine solution. A
standardized solution is one whose concentration is known approximately
from its preparation and is then determined precisely by making a
controlled titration against a known standard.
The iodine stock solution on the Chemicals shelf has been prepared
according to a calculation that gives an approximate concentration of
0.0015M when originally prepared (about five weeks ago). Your task in
lab part 1 is to determine the exact current concentration of the iodine
solution. In lab part 2, you will use this now-known concentration of
iodine solution to determine the concentration of vitamin C in samples
of orange juice.
Per the redox equation pair that describes the reaction occurring in
this lab, one mole of ascorbic acid will react with one mole of iodine.
Thus the following titration formula will apply:
M1*V1 = M2*V2
where M is the concentration of each component (in mol/L) in the
solution being added together, and V is the volume (in liters) of the
particular solution being added.
When standardizing the iodine, M1 and V1 for the ascorbic acid solution
are known, as is the volume of iodine delivered from the burette, and so
you can solve for the current concentration of the iodine solution.
Later, the current iodine concentration will be known along with the
volume of orange juice and the volume of iodine added from the burette,
and so you can solve for the concentration of ascorbic acid in the
orange juice sample.
Lab Part 1 Procedure
The prepared iodine solution on the Chemicals shelf with a stated
(anticipated) concentration of 0.0015M is standardized (to determine the
current actual concentration) by performing the following titration:
1. Prepare a solution of pure ascorbic acid of known strength.
1a. Weight out and add exactly 0.05 g ascorbic acid to the
volumetric flask.
1b. Fill the volumetric flask with water making a 100 mL solution.
Record the amount of ascorbic acid used and the total volume prepared.
2. Do a titration of the iodine stock solution.
2a. Pour 10.0 mL of the ascorbic acid solution into a 150
Erlenmeyer flask.
2b. Add 10.0 mL of distilled water to the 150 Erlenmeyer flask.
2c. Add 1.0 mL of the starch indicator to the 150 Erlenmeyer
flask.
2d. Take a burette and fill it with 50.00 mL of the iodine stock
solution (with nominal strength of 0.0015M.)
2e. Place flask under the lower half of the burette.
2f. Titrate the ascorbic acid sample by adding iodine solution
until the endpoint is reached where solution in the flask turns dark
blue when adding only one more drop (the "last 0.05 mL") of iodine
solution.
Lab Part 1 assignment:
1. Calculate the molarity of the freshly-prepared ascorbic acid
solution:
(a) Mass of ascorbic acid used: 0.05 gram
(b) Moles of ascorbic acid (MW=176.1 g/mol): 0.05 g / (176.1
g/mol) = 0.000284 mol
(c) Volume of solution (mL): 100mL
(d) Ascorbic acid concentration (mol/L): 0.000284 mol / 0.1 Liter
= 0.00284 mol/L
2.Do a titration, record and calculate the following:
Volume of iodine solution added (mL):
(b) Concentration of the iodine solution, using the formula M1*V1 =
M2*V2:
Lab results of Fine Titration
Burette
Volume
(mL) Volume of iodine
added to solution
(mL)
50 0
34 16
33.95 16.05
Note: Fine Titration, one increment of 16mL of iodine added and followed
by one 0.05 mL drop when solution turns blue.
Volume of iodine solution added (mL):
Volume added = Volume initial – Volume final
Volume added = 50 mL- 33.95 mL = 16.05mL
Concentration of the iodine solution, using the formula M1*V1 = M2*V2:
M1 ascorbic acid concentration (mol/L) = 0.00284 mol/L
V1 ascorbic acid volume = 21 mL ( 10mL ascorbic acid + 10 ml H2O + 1 mL
starch)
M2 concentration of iodine solution (unknown)
V2 Volume of iodine solution added = 16.05mL
M1(acid) * V1 (acid) = M2 (base) * V2 (base)
M2 = (M1 * V1) / V2
M2 = (0.00284 mol/L * 0.021 L) / 0.01605 L = 0.00372 mol
3. From your two or more best "fine" titrations, calculate the average
iodine concentration, thus determining today's standard value for the
iodine solution.
Note: 2nd fine titration has same results as the first.
Average titration = (1st titration + 2nd titration) / 2
Average titration = (0.00372 mol + 0.00372mol) /2
Average titration = 0.00372mol
Today’s standard value for iodine solution = 0.00372 mol
Lab Part 2 Procedure
1. Determine the ascorbic acid concentration in commercial orange juice
from a freshly opened container.
a. Prepare a sample of orange juice from the new container by
adding 10.0 mL of the juice to an Erlenmeyer flask.
b. Add 10.0 mL of water to flask
c. Add 1.0 mL of starch indicator to flask
2. Titrate the orange juice with the recently-standardized iodine
solution. Record the volume of iodine delivered in titration.
3. Repeat steps 1 and 2 above using the week-old orange juice.
Lab Part 2 Assignment:
1. Do a titration of the fresh orange juice(OJ), record and calculate
the following:
(a) Volume of iodine solution added (mL):
(b) Concentration of the ascorbic acid in the juice, using the
formula M1*V1 = M2*V2:
Lab results of Fine Titration
Burette
Volume
(mL) Volume of iodine
added to solution
(mL)
50 0
40 10
39.95 10.05
Note: Fine Titration, one increment of 10 mL of iodine added and
followed by one 0.05 mL drop when solution turns blue.
Volume of iodine solution added (mL):
Volume added = Volume initial – Volume final
Volume added = 50 mL - 39.05 mL = 10.05mL
Concentration of the iodine solution, using the formula M1*V1 = M2*V2:
M1 ascorbic acid concentration (mol/L) = (unknown)
V1 ascorbic acid volume = 21 mL (10mL new OJ ascorbic acid + 10 ml H2O
+ 1 mL starch)
M2 concentration of iodine solution = 0.00372 mol
V2 Volume of iodine solution added = 10.05mL
M1(acid) * V1 (acid) = M2 (base) * V2 (base)
M1 = (M2 * V2) / V1
M1 = (0.00372 mol * 0.01005 L) / 0.021 L = 0.0018mol
2. Calculate the average ascorbic acid concentration for the fresh
orange juice.
Note: 2nd fine titration has same results as the first.
Average titration = (1st titration + 2nd titration) / 2
Average titration = (0.0018 mol + 0.0018 mol) / 2
Average ascorbic acid concentration for the fresh orange juice = 0.0018
mol
3. Do a titration of the week-old orange juice, record and calculate the
following:
(a) Volume of iodine solution added (mL):
(b) Concentration of the ascorbic acid in the juice, using the
formula M1*V1 = M2*V2:
Lab results of Fine Titration
Burette
Volume
(mL) Volume of iodine
added to solution
(mL)
50 0
48 2
47.95 2.05
Note: Fine Titration, one increment of 2 mL of iodine added and followed
by one 0.05 mL drop when solution turns blue.
Volume of iodine solution added (mL):
Volume added = Volume initial – Volume final
Volume added = 50 mL - 47.95 mL = 2.05mL
Concentration of the iodine solution, using the formula M1*V1 = M2*V2:
M1 ascorbic acid concentration (mol/L) = (unknown)
V1 ascorbic acid volume = 21 mL (10mL old ascorbic acid + 10 ml H2O + 1
mL starch)
M2 concentration of iodine solution = 0.00372 mol
V2 Volume of iodine solution added = 2.05mL
M1(acid) * V1 (acid) = M2 (base) * V2 (base)
M1 = (M2 * V2) / V1
M1 = (0.00372 mol * 0.00205 L) / 0.021 L = 0.00036 mol
4. Calculate the average ascorbic acid concentration for the week-old
orange juice.
Note: 2nd fine titration has same results as the first.
Average titration = (1st titration + 2nd titration) / 2
Average titration = (0.00036 mol + 0.00036 mol) / 2
Average ascorbic acid concentration for the week old orange juice =
0.00036 mol
This is where help is needed. Not sure how to calculate mg per ml of
juice without density. Please show each step of calculation.
5. Report the average amount of ascorbic acid in each of the 2
containers of commercial orange juice in units of mg per mL of juice.
(The molecular weight of ascorbic acid is 176.12).
6. The recommended minimum daily requirement for vitamin C is 60 mg per
day.
6a. What percentage of this requirement is in one cup (200 mL) of
fresh orange juice?
6b. What percentage of this requirement is in one cup (200 mL) of
week-old orange juice?
7. What happens to the ascorbic acid in orange juice over time? (hint:
oxygen makes up 20% of our air.)
Chemistry Homework Solutions
#102510
Tritraton conversion
Please review lab content for accuracy and provide detailed calculation solution to 5, 6, & 7.
Titration Lab Background:
The minimum daily requirement for vitamin C is 30 mg, and the recommended daily allowance is 60-70 mg. You can find the vitamin C content on the labels of most commercially available juices. How close is their reported number to the actual amount of the vitamin in the juice? You will test that in this lab by titrating the juice with an iodine solution, using starch as an indicator.
The formula for ascorbic acid is C6H8O6 and the structures for the reduced form and for the oxidized form (dehydroascorbic acid) are shown below (as well as possible given limitation of typing characters), which illustrates the first half of the redox reaction.
OH OH O
| | /
H - C - C - C C = O
| | | /
H H H C=C
| |
HO OH
Ascorbic Acid
C6H8O6
(Reduced Form)
---> 2H+ plus 2e- plus
OH OH O
| | /
H - C - C - C C = O
| | | /
H H H C-C
|| ||
O O
Dehydroascorbic Acid
C6H6O6
(Oxidized Form)
The amount of ascorbic acid can be determined by carrying out a redox titration with a known concentration of iodine in solution. In this procedure the iodine is reduced by the ascorbic acid to form iodide, as shown in the other half of this redox equation.
I2 + 2e- --> 2I-
The titration end point is reached when slightly more moles of iodine have been added to the sample being tested than the number of moles of ascorbic acid in the sample (leaving some "free" iodine molecules present in the sample). To highlight when the excess iodine is present, thyodene "starch" indicator is added to the titrant. Thyodene reacts with the excess iodine to form a bright blue complex. Thyodene does not form this complex with iodide.
Before doing a titration on a solution of unknown concentration (such as orange juice), you must first prepare a standardized iodine solution. A standardized solution is one whose concentration is known approximately from its preparation and is then determined precisely by making a controlled titration against a known standard.
The iodine stock solution on the Chemicals shelf has been prepared according to a calculation that gives an approximate concentration of 0.0015M when originally prepared (about five weeks ago). Your task in lab part 1 is to determine the exact current concentration of the iodine solution. In lab part 2, you will use this now-known concentration of iodine solution to determine the concentration of vitamin C in samples of orange juice.
Per the redox equation pair that describes the reaction occurring in this lab, one mole of ascorbic acid will react with one mole of iodine. Thus the following titration formula will apply:
M1*V1 = M2*V2
where M is the concentration of each component (in mol/L) in the solution being added together, and V is the volume (in liters) of the particular solution being added.
When standardizing the iodine, M1 and V1 for the ascorbic acid solution are known, as is the volume of iodine delivered from the burette, and so you can solve for the current concentration of the iodine solution. Later, the current iodine concentration will be known along with the volume of orange juice and the volume of iodine added from the burette, and so you can solve for the concentration of ascorbic acid in the orange juice sample.
Lab Part 1 Procedure
The prepared iodine solution on the Chemicals shelf with a stated (anticipated) concentration of 0.0015M is standardized (to determine the current actual concentration) by performing the following titration:
1. Prepare a solution of pure ascorbic acid of known strength.
1a. Weight out and add exactly 0.05 g ascorbic acid to the volumetric flask.
1b. Fill the volumetric flask with water making a 100 mL solution.
Record the amount of ascorbic acid used and the total volume prepared.
2. Do a titration of the iodine stock solution.
2a. Pour 10.0 mL of the ascorbic acid solution into a 150 Erlenmeyer flask.
2b. Add 10.0 mL of distilled water to the 150 Erlenmeyer flask.
2c. Add 1.0 mL of the starch indicator to the 150 Erlenmeyer flask.
2d. Take a burette and fill it with 50.00 mL of the iodine stock solution (with nominal strength of 0.0015M.)
2e. Place flask under the lower half of the burette.
2f. Titrate the ascorbic acid sample by adding iodine solution until the endpoint is reached where solution in the flask turns dark blue when adding only one more drop (the "last 0.05 mL") of iodine solution.
Lab Part 1 assignment:
1. Calculate the molarity of the freshly-prepared ascorbic acid solution:
(a) Mass of ascorbic acid used: 0.05 gram
(b) Moles of ascorbic acid (MW=176.1 g/mol): 0.05 g / (176.1 g/mol) = 0.000284 mol
(c) Volume of solution (mL): 100mL
(d) Ascorbic acid concentration (mol/L): 0.000284 mol / 0.1 Liter = 0.00284 mol/L
2.Do a titration, record and calculate the following:
(a) Volume of iodine solution added (mL):
(b) Concentration of the iodine solution, using the formula M1*V1 = M2*V2:
Lab results of Fine Titration
Burette
Volume
(mL) Volume of iodine
added to solution
(mL)
50 0
34 16
33.95 16.05
Note: Fine Titration, one increment of 16mL of iodine added and followed by one 0.05 mL drop when solution turns blue.
(a) Volume of iodine solution added (mL):
Volume added = Volume initial - Volume final
Volume added = 50 mL- 33.95 mL = 16.05mL
(b) Concentration of the iodine solution, using the formula M1*V1 = M2*V2:
M1 ascorbic acid concentration (mol/L) = 0.00284 mol/L
V1 ascorbic acid volume = 21 mL ( 10mL ascorbic acid + 10 ml H2O + 1 mL starch)
M2 concentration of iodine solution (unknown)
V2 Volume of iodine solution added = 16.05mL
M1(acid) * V1 (acid) = M2 (base) * V2 (base)
M2 = (M1 * V1) / V2
M2 = (0.00284 mol/L * 0.021 L) / 0.01605 L = 0.00372 mol
3. From your two or more best "fine" titrations, calculate the average iodine concentration, thus determining today's standard value for the iodine solution.
Note: 2nd fine titration has same results as the first.
Average titration = (1st titration + 2nd titration) / 2
Average titration = (0.00372 mol + 0.00372mol) /2
Average titration = 0.00372mol
Today's standard value for iodine solution = 0.00372 mol
Lab Part 2 Procedure
1. Determine the ascorbic acid concentration in commercial orange juice from a freshly opened container.
a. Prepare a sample of orange juice from the new container by adding 10.0 mL of the juice to an Erlenmeyer flask.
b. Add 10.0 mL of water to flask
c. Add 1.0 mL of starch indicator to flask
2. Titrate the orange juice with the recently-standardized iodine solution. Record the volume of iodine delivered in titration.
3. Repeat steps 1 and 2 above using the week-old orange juice.
Lab Part 2 Assignment:
1. Do a titration of the fresh orange juice(OJ), record and calculate the following:
(a) Volume of iodine solution added (mL):
(b) Concentration of the ascorbic acid in the juice, using the formula M1*V1 = M2*V2:
Lab results of Fine Titration
Burette
Volume
(mL) Volume of iodine
added to solution
(mL)
50 0
40 10
39.95 10.05
Note: Fine Titration, one increment of 10 mL of iodine added and followed by one 0.05 mL drop when solution turns blue.
(a) Volume of iodine solution added (mL):
Volume added = Volume initial - Volume final
Volume added = 50 mL - 39.05 mL = 10.05mL
(b) Concentration of the iodine solution, using the formula M1*V1 = M2*V2:
M1 ascorbic acid concentration (mol/L) = (unknown)
V1 ascorbic acid volume = 21 mL (10mL new OJ ascorbic acid + 10 ml H2O + 1 mL starch)
M2 concentration of iodine solution = 0.00372 mol
V2 Volume of iodine solution added = 10.05mL
M1(acid) * V1 (acid) = M2 (base) * V2 (base)
M1 = (M2 * V2) / V1
M1 = (0.00372 mol * 0.01005 L) / 0.021 L = 0.0018mol
2. Calculate the average ascorbic acid concentration for the fresh orange juice.
Note: 2nd fine titration has same results as the first.
Average titration = (1st titration + 2nd titration) / 2
Average titration = (0.0018 mol + 0.0018 mol) / 2
Average ascorbic acid concentration for the fresh orange juice = 0.0018 mol
3. Do a titration of the week-old orange juice, record and calculate the following:
(a) Volume of iodine solution added (mL):
(b) Concentration of the ascorbic acid in the juice, using the formula M1*V1 = M2*V2:
Lab results of Fine Titration
Burette
Volume
(mL) Volume of iodine
added to solution
(mL)
50 0
48 2
47.95 2.05
Note: Fine Titration, one increment of 2 mL of iodine added and followed by one 0.05 mL drop when solution turns blue.
(a) Volume of iodine solution added (mL):
Volume added = Volume initial - Volume final
Volume added = 50 mL - 47.95 mL = 2.05mL
(b) Concentration of the iodine solution, using the formula M1*V1 = M2*V2:
M1 ascorbic acid concentration (mol/L) = (unknown)
V1 ascorbic acid volume = 21 mL (10mL old ascorbic acid + 10 ml H2O + 1 mL starch)
M2 concentration of iodine solution = 0.00372 mol
V2 Volume of iodine solution added = 2.05mL
M1(acid) * V1 (acid) = M2 (base) * V2 (base)
M1 = (M2 * V2) / V1
M1 = (0.00372 mol * 0.00205 L) / 0.021 L = 0.00036 mol
4. Calculate the average ascorbic acid concentration for the week-old orange juice.
Note: 2nd fine titration has same results as the first.
Average titration = (1st titration + 2nd titration) / 2
Average titration = (0.00036 mol + 0.00036 mol) / 2
Average ascorbic acid concentration for the week old orange juice = 0.00036 mol
This is where help is needed. Not sure how to calculate mg per ml of juice without density. Please show each step of calculation.
5. Report the average amount of ascorbic acid in each of the 2 containers of commercial orange juice in units of mg per mL of juice. (The molecular weight of ascorbic acid is 176.12).
6. The recommended minimum daily requirement for vitamin C is 60 mg per day.
6a. What percentage of this requirement is in one cup (200 mL) of fresh orange juice?
6b. What percentage of this requirement is in one cup (200 mL) of week-old orange juice?
7. What happens to the ascorbic acid in orange juice over time? (hint: oxygen makes up 20% of our air.)
See attached file for full problem description.
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