CEAC 104
GENERAL CHEMISTRY
Experiment 1
Preparation and Analysis of Potassium
Trisoxalatoferrate(III) Trihydrate, K3[Fe(C2O4)3].3H2O
Purpose: To prepare the complex trisoxalatoferrate(III), Fe(C2O4)3-3 anion and
isolate it as its hydrated potassium salt, K3[Fe(C2O4)3].3H2O. Also, to study the
photochemical reduction of the sample.
APPARATUS AND CHEMICALS:
K2C2O4.H2O (Potassium oxalate monohydrate)
filter paper
FeCl3.6H2O (Iron (III) chloride hexahydrate)
distilled water
K3Fe(CN)6 solution (Potassium hexacyanoferrate(III))
funnel
H2SO4 solution (Sulfuric acid)
100-mL beaker
test tubes
THEORY:
Potassium trisoxalatoferrate(III) trihydrate, K 3[Fe(C2O4)3].H2O is a green crystalline salt,
soluble in hot water but rather insoluble when cold. It can be prepared by the reaction of
K2C2O4.H2O with FeCl3.6H2O.
3K2C2O4.H2O(aq) + FeCl3.6H2O(aq)
K3Fe(C2O4)3].3H2O(aq) + 3KCl(aq)
The complex anion is photo-sensitive. This means that upon exposure to light of an
appropriate wavelength (<450 nm in this case) the Fe(C 2O4)3-3 undergoes an intramolecular
redox reaction in which the Fe(III) anion is reduced to Fe(II) while one of the oxalate groups
is oxidized to CO2.
[Fe(C2O4)3]3-
Fe2+ + 5/2 C2O42- +CO2(g)
As mentioned above, light causes an internal electron-transfer reaction to occur in
[Fe(C2O4)2]3- ion, producing CO2 and Fe2+ ions. The Fe2+ that is produced can readily be
detected by adding a solution of potassium ferricyanide K3Fe(CN)6. A deep blue colored
ferroferri cyanide complex is formed.
Fe2+ + Fe(CN)63-
Fe[Fe(CN)6]ferroferricyanide deep blue.
PROCEDURE:
A. Preparation of K3[Fe(C2O4)3].3H2O
1.
Weigh approximately 9.0 g of hydrated potassium oxalate, K2C2O4.H2O into a 250 mL
beaker.
2.
Add 30 mL of distilled water and heat to dissolve (do not boil).
3.
In a second small beaker dissolve 4.4 g of FeCl3.6H2O in a minimum amount of
cold water (10-15 mL). Add the FeCl 3.6H2O solution to the warm oxalate
solution and stir with a glass rod. Allow the product to crystallize (away from
strong sunlight) by cooling the solution in an ice-water mixture.
4.
Collect the crystalline product by filtration. The product is K3[Fe(C2O4)3].3H2O.
B. Blueprinting
1. Wet a piece of filter paper with [Fe(C2O4)2]3- solution.
2. Leave it to dry. (Meanwhile you can follow part C)
3. Place small opaque objects (coins, keys, etc.) on the paper.
4. Irradiate for few minutes using a light source (If not available you may use bright
sunlight)
5. Dip the paper into potassium ferricyanide solution (CAUTION potassium
ferricyanide is poisonous. Avoid contact with your skin. If it happens immediately
wash your skin with plenty of water.)
6. Remove the developed blueprint and dip in a beaker of distilled water to wash off excess
ferricyanide solution. Explain your observations.
C.Photochemical Reaction of [Fe(C2O4)2]31. Dissolve 0.7 g of your complex in 100 mL of distilled water in a flask. Add 3 mL of 2 M
H2SO4 and swirl the mixture. To each three labeled test tubes add 10 mL of this solution.
2. Keep one tube away from the light source as the control and irradiate the
remaining two tubes with the light source for 1 and 5 minutes respectively.
3. To all three tubes add 1 mL of 0. 1 M potassium ferricyanide solution
K3Fe(CN)6.
4. Record and explain your observations.
DATA SHEET
Preparation and Analysis of Potassium Trisoxalatoferrate(III) Trihydrate,
K3[Fe(C2O4)3].3H2O
Student's Name
:
Laboratory Section/Group No :
Assistant's Name and Signature :
A .Blueprinting
Observations:
Explain:
B.Photochemical Reaction of [Fe(C2O4)2]-3
Observations:
1st sample:
2nd sample:
3rd sample:
Explain:
Date:
CEAC 104
GENERAL CHEMISTRY
Experiment 2
Titration of Antacids
Purpose: To standardize a solution of base using the analytical technique known
as titration. Using this standardized solution, you will determine the acid
neutralizing power of two commercially available antacid tablets.
APPARATUS AND CHEMICALS:
500-mL Erlenmeyer flask
Wash bottle
250-mL Beaker
Ring stand
Sodium hydroxide (NaOH)
Hydrochloric acid (HCl) (primary
standard)
50-mL buret
Watch glass
Phenolphthalein solution
Buret clamp
Potassium hydrogen phthalate
Antacid
(KHPh)
THEORY:
An understanding of the properties of acids and bases is an essential part of
understanding chemical reactions. In aqueous solutions, a compound that produces H+ ions
upon dissolution is termed an acid. A compound that produces OH– ions when dissolved in
water is called a base. The reaction of an acid and base is a neutralization reaction, the
products of which are a salt and water. In an aqueous solution, virtually all of the OH– ions
present will react with all of the H+ ions that are present:
H+ (aq) + OH– (aq) → H2O (l)
Because this reaction is essentially quantitative, it is possible to determine the concentration
of an acid or base in an aqueous solution with high accuracy.
When a solution of hydrochloric acid, HCl, is exactly neutralized with a solution of
sodium hydroxide, NaOH, the number of moles of NaOH used will equal the number of
moles of HCl originally present. The following relationship then holds true:
moles NaOH = moles HCl
(MNaOH)(VNaOH in liters) = (MHCl)(VHCl in liters)
where M = concentration in molarity and V= volume. If three of the above quantities are
known, the fourth can be calculated.
In order to determine when a solution has been exactly neutralized, an acid-base
indicator is used that changes color in a certain pH range (pH is a scale used to measure
acidity). This color change is termed the endpoint of the titration. Because the pH of a neutral
solution is 7, an indicator that changes color near this pH should be used for an acid-base
titration. Phenolphthalein indicator changes color in the range pH = 8.3 – 10.0 and can be
used to determine when the correct amount of base has been added to an acidic solution to
exactly neutralize it.
Standardization of a Sodium Hydroxide Solution
In order to determine the concentration of an acidic or basic solution, it is necessary to
know the number of moles of acid or base that are required to neutralize it. This quantity can
be calculated by accurately weighing a solid sample of an acid or a base, dissolving it in water
and titrating this solution; that is, adding the solution of unknown concentration to it until the
endpoint has been reached.
It is difficult to accurately weigh sodium hydroxide since it is hygroscopic (absorbs
water readily from air). A solution of NaOH is usually standardized using an acid known as a
primary standard.
A primary standard must satisfy the four following criteria:
1. Solid compound that is not hygroscopic and can be easily handled
2. Is available in very pure form
3. Stable
4. Has a medium to high molecular weight.
For this experiment, a solution of NaOH, which has an approximate concentration of
0.1 M, will be standardized using potassium acid phthalate, KHPh. The molecular weight of
KHPh is 204.23 g/mole, and it has one acidic proton, which will react quantitatively with
OH–:
OH– (aq) + KHPh (aq) → H2O (l) + KPh– (aq)
For the highest accuracy, a sample size is chosen such that it will consume as large a
volume of the base as possible without exceeding the capacity of the buret. If a 25 mL buret is
used, the amount of KHPh is chosen such that it will require approximately 20 mL of 0.1 M
NaOH solution to reach the endpoint. Thus, about 0.002 moles, or 0.4 g, of KHPh is needed.
At the endpoint, the number of moles of NaOH equals the number of moles of KHPh
used:
or
Once the NaOH solution has been standardized, it can be used to determine the acid
neutralizing capacity of an antacid tablet.
Determination of the Acid Neutralizing Capacity of an Antacid Tablet
The stomach has an acidic interior generated by dilute HCl, "stomach acid", which
insures proper digestion. When the acidity of the stomach becomes high enough to cause
discomfort, brought about by the ingestion of certain types of food, an antacid preparation can
be taken to neutralize the excess stomach acid. The active ingredient in every antacid is a
base, the most common being metal hydroxides, metal carbonates or a mixture of the two.
Table 2.1 lists the active ingredients in several commercial brands of antacid.
Table 2.1. Brands of antacid tablets and their major ingredients
The acid neutralizing capacity of a tablet is the amount of hydrochloric acid that it can
neutralize. It is the quantity that is referred to in some advertisements when it is stated that the
tablet “neutralizes x times its weight in stomach acid”. This capacity can be determined by a
technique called back-titration. A known amount of antacid is dissolved in an excess of HCl,
and then the excess acid is back-titrated with standardized NaOH solution. When the endpoint
is reached, the number of moles of acid that was added to the antacid sample is equal to the
number of moles of base present, NaOH plus the antacid. Therefore, the number of moles of
HCl that was neutralized by the antacid is equal to the total number of moles of HCl added
minus the number of moles that were neutralized by the NaOH:
moles acid neutralized = (moles of HCl added) – (moles of NaOH required for backtitration)
= (MHCl x VHCl) – (MNaOH x VNaOH)
where M = molarity and V = volume in liters.
REVIEW QUESTIONS:
Before beginning this experiment in the laboratory. you should be able to answer the
following questions:
1. Define standardization and state how you would go about doing it.
2. Define titration and molarity.
3. What are equivalence points and end points and how do they differ?
4. What is the molarity of a solution that contains 1.89 g of H2C2O4. 2H2O in 100 mL of
solution?
PROCEDURE:
Standardization of NaOH solution
1. Accurately weigh out a sample of approximately 0.3-0.4 g of primary standard potassium
hydrogen phthalate, KHPh, which has been previously dried at 120°C. Do not use more
than 0.4 g.
2. Dissolve the KHPh sample in about 50 mL of CO2-free water and add 2-3 drops of
phenolphthalein indicator solution.
3. Begin adding the approximately 0.1 M NaOH solution from the buret while continuously
swirling the flask contents. Do not open the stopcock completely. As the endpoint nears,
a pink color will appear at the point where the NaOH mixes with the flask contents. This
color will disappear with subsequent swirling. When the color persists for 30 seconds
after swirling, the endpoint has been reached. The color will fade after some time due to
absorption of CO2 from the air.If a deep pink color results, the endpoint has been
overrun.
4. When the endpoint is reached, record the final buret reading to the nearest 0.01 mL.
Refill the buret so that you do not run out of NaOH solution in the middle of the next
titration.
5. Weigh two more 0.3-0.4 g samples of KHPh into two separate 250 mL Erlenmeyer
flasks. Dissolve one of these samples in about 50 mL of CO2-free water and repeat the
titration procedure. Dissolve the last sample, and titrate it as well. Your three
determinations should not differ by more than 0.002 M. In the case of poor precision, an
additional sample may be run if there is time. Use an average of these molarities for
analyzing the antacid in the next part of the experiment.
Back-titration of an antacid
1. Choose the brands and obtain 2 antacid tablets. Avoid touching them with your fingers as
much as possible. Record the brand name. Weigh each tablet separately on weighing
paper to the nearest 0.001 g. Transfer each tablet to a 250 mL Erlenmeyer flask and label
the flasks with the corresponding masses.
2. Add 20 mL of standardized HCl from the beaker slowly to a flask containing one of the
antacid tablets. Record precisely how much acid was added using the initial and final
readings of the buret. Also, be sure to record the exact molarity of the HCl solution
written on the label of the bottle. Hydrochloric acid is corrosive!
3. Rinse down the inner walls of the flask with 50 mL of distilled water, and swirl the flask.
You may need to crush the tablet with a glass stirring rod to allow the reaction to go to
completion, but be sure to rinse all the solution and particles off the rod before removing
it. The tablet may contain an insoluble binder and filler that will not dissolve; however,
be certain that no large chunks or chips of the tablet remain.
Carbonate ions in the sample will react with the HCl and produce CO2. The CO2 can be
driven off by heating the contents of the flask just to boiling:
2 H+ (aq) + CO32- (aq) ↔ H2CO3 (aq) → CO2 (g) + H2O (l)
heat
4. Adjust the hot plate to maintain a gentle boil for exactly 5 minutes. Remove from the heat
and allow the flask to cool until it is comfortable to hold.
5. Add 8 drops of phenolphthalein to the flask and swirl to mix; then rinse down the sides of
the flask with your wash bottle. If the solution is pink at this point, more acid must be
added from the buret, 1 mL at a time, until the color disappears.
6. Titrate the solution, as in the first part of the experiment, to the pale pink endpoint using
the standardized NaOH solution. Refill your buret and repeat the entire procedure with
the other antacid tablet.
*When you have finished all the titrations, drain the NaOH from your buret into the sink,
rinse it thoroughly with distilled water and return it to your instructor. Clean up your bench
area and sink.
*You will be asked to compare the weight and cost effectiveness of the two brands
of antacid tablets. Before you leave the laboratory, find a student who tested the brand
that you did not test and exchange data with him/her.
CALCULATIONS:
1. Calculate the average molarity of the NaOH solution from your three standardization
trials.
2. Calculations to determine the effectiveness of the antacid;
• Calculate the average number of moles of HCl neutralized per antacid tablet.
• Calculate the number of moles of HCl neutralized per gram of antacid
(weight
effectiveness).
• Calculate the average cost of the antacid tablet that would be needed to neutralize 1.00
mole of HCl (cost effectiveness).
3. The HCl concentration in a hyperacidic stomach is 0.03 M. The volume of liquid in the
stomach is 300 mL. How many tablets of the antacid that you analyzed would have to be
taken to bring the concentration of HCl in the stomach to a more normal level of 0.0003
M? Show your work.
QUESTIONS:
1. Define analyte, indicator solution, standard solution and titrant.
2. Titration reveals that 11.6 mL of 3.0 M sulfuric acid are required to neutralize the
sodium hydroxide in 25.00 mL of NaOH solution. What is the molarity of the NaOH
solution?
DATA SHEET
Titration of an Antacids
Student’s Name
:
Laboratory Section/Group No
:
Date:
Assistant’s Name and Signature :
A. Standardization of NaOH:
1. Weight of KHPh:………......................g
2. Molarity of KHPh: ………….............M
3. Volume of NaOH used:1st………….., 2nd……………, 3rd………........
4. Average volume of NaOH: ………….mL
5. Molarity of NaOH:……………..........M
B. Back Titration:
Sample
CALCULATIONS:
Volume of HCl
Volume of NaOH
CEAC 104
GENERAL CHEMISTRY
Experiment 3
The Thermodynamics of Solubility of Potassium Chlorate
Purpose: To observe the effect of temperature on solubility product constant
(Ksp) and to draw the graph of lnKsp vs 1/T in order to calculate the standard
enthalpy change (ΔH°) and standard entropy change (ΔS°) via this graph.
APPARATUS AND CHEMICALS:
5 pieces of test tubes
Glass rod
Thermometer
Bunsen burner
Pipet bulb
Distilled water
10-mL pipet
Ring stand
Potassium Chlorate(KCIO3)
THEORY:
The solubility of KCIO3 will be measured at various temperatures. From the molar
solubility, s: moles of solute per liter of solvent, a value of the solubility product constant,
Ksp, can be calculated at various temperatures for the following reaction:
KCIO3(s) ⇄ K+(aq) + CIO3-(aq)
Ksp = [K+][CIO3-] = s2
The solubility of a solid is exponential with respect with temperature. A plot of the solubility
product constant, Ksp, vs temperature (K) will give an exponential curve.
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The relationship between ΔHo and the Ksp comes from the free energy equations
(1) ΔGo = - RT ln K, where R = the gas constant, 8.3143 Joule/mol K
and T = absolute temperature in degrees Kelvin.
(2) ΔGo = ΔHo - TΔSo , where ΔHo is the standard enthalpy change
and ΔSo is the standard entropy change.
From (1) and (2) comes the equation
(3) - RT ln K = ΔHo - TΔSo,
Both the molal heat of solution ΔHo and the entropy ΔSo values can be obtained for the
reaction by solving the equation relating the slope of the graph with ΔHo and the y- intercept
with ΔSo. The plot of ln K vs 1/T in K-1 will give the following equation of a straight line:
Each group will be assigned a different composition of solution to measure. The data from
each group will be combined and plotted.
SAFETY PRECAUTIONS


KCIO3 is flammable and strong oxidizing agent. Therefore, dispose of all wastes by
flushing them down the sink with plenty of running water.
Avoid any contact between the KCIO3 and paper.
REVIEW QUESTIONS:
Before beginning this experiment in the laboratory, you should be able to answer the
following questions:
1. Suppose a reaction has negative ΔH° and negative ΔS°. Will more or less product be
present at equilibrium as the temperature is raised?
2. Fisherman know that on hot summer days, the largest fish will be found in deep sinks
in lake bottoms, where the water is coolest. What is the reason of it, think about the
temperature dependence of oxygen solubility in water?
3. Nitrogen dioxide reacts with water to produce nitric acid according to following
reaction:
Predict the sign of ΔS° for this reaction
4. Suggest a reason why the value of ΔHsoln for a gas such as CO2 , dissolving in water, is
negative.
PROCEDURE:
1.In a test tube, weigh a sample of KCIO3 to the nearest 0.01g. The amount of sample will
range from 1 to 5 grams in 1 gram increments.
2. Pipet 10 mL of distilled water into the test tubes. Stir the mixture with a glass rod.
3. Gently warm the mixture until the crystals dissolve. For high sample amounts, an
evaporating dish can be used as a water bath to transfer heat more evenly. Insert thermomer
into test tube and try to note the temperature at which the last crystals go into solution. Then
allow the solution to cool until crystals reappear and again note the temperature.
4. Record the saturation temperature you have just determined and enter your values in your
data sheet.
CALCULATIONS:
A. Ksp values.
1. Write the equation for the equilibrium between solid KCIO3 and its ions in solution and the
expression for the Ksp of KCIO3.
2. From the weight-volume data, calculate the molar concentration of KCIO3 for each
composition. This is the molar solubility at the saturation temperature measured.
3. Calculate the solubility product of KCIO3 for each trial. Use the concentrations found in
step 1.
4. Tabulate the resulting values of solubility, Ksp, and temperature in your lab datasheet.
5. Using Excel, graph temperature (Kelvin) on the horizontal axis (x-axis) versus Ksp on the
vertical axis (y-axis).
B. Molar heat of solution of KCIO3.
1. Using excel calculate values of 1/T, the reciprocal of the absolute temperature, and ln Ksp
for each of the points graphed in part A.
2. Make a new graph with ln Ksp on the vertical axis and 1/T (K-1) on the horizontal axis. Use
the linear curve fit option in excel to obtain the best fit linear fit.
3. From the slope of your graph calculate ΔHo, the heat of solution for KCIO3
Also determine ΔS° from the y-intercept.
Tabulate data for the written report and include the curve fit equation with the graph of lnKsp
vs 1/T.
QUESTIONS:
1. When a certain solid dissolves in water, the solution becomes cold. Is ΔH soln for this
solute positive or negative? Is the solubility of this substance likely to increase or
decrease with increasing temperature?
2. What is the value of equilibrium constant for a reaction for which ΔG°= 0? What will
happen to the composition of the system if we begin the reaction with the products?
3.The solubility of ZnS at 25oC is 3.5 x 10-12 M. What is the Ksp for ZnS? What is ΔG° at
25°C?
DATA SHEET
The Thermodynamics of Solubility of Potassium Chlorate
Student’s Name
:
Laboratory Section/Group No
:
Date:
Assistant’s Name and Signature :
A. Ksp Values
1. Write the reaction equation and the solubility product expression for KCIO3.
2.
Amount
of Molar
KCIO3 (g)
Concentration of
KCIO3 (mol/L)
Ksp
3. Draw Ksp vs Temperature (K) graph using Excel.
T (°C)
T (K)
B. Molal Heat of Solution of KIO3
1.
1/T (K-1)
ln Ksp
2. Draw ln Ksp vs 1/T graph using Excel and calculate ΔH° from the slope and ΔS° from
y-intercept.
ΔH° =
ΔS° =