DAWSON COLLEGE
CHEMISTRY OF SOLUTIONS
202-NYB-05
Laboratory Experiments
Last revision: Winter 2018
TABLE OF CONTENTS
Safety Rules
2
List of chemicals
7
Volumetric Analysis of an Acid Solution
9
Colligative Properties
13
Spectrophotometry
19
Chemical Equilibrium I: The Equilibrium Constant
25
Titration Curves
31
Chemical Equilibrium II: The Solubility Product
37
The Rate Law in Chemical Kinetics
41
Activation Energy in Chemical Kinetics
47
The Indirect Determination of a Thermodynamic Quantity
51
Common laboratory equipment
57
Periodic Table of Elements
59
1
SAFETY RULES
1. ALWAYS WEAR SAFETY GLASSES IN THE LABORATORY TO PROTECT YOUR EYES
2. DO NOT WEAR CONTACT LENSES IN THE LABORATORY
3. ALWAYS WEAR A LAB COAT TO PROTECT YOUR CLOTHES AND TO REDUCE THE CHANCE OF
ANY CHEMICAL CONTACT WITH YOUR SKIN
4. ALWAYS WEAR CLOSED SHOES OR BOOTS. OPEN TOED SHOES OR SANDALS DO NOT
PROTECT AGAINST CHEMICAL SPILLS. DO NOT WEAR HIGH HEEL OR SLIPPERY SOLED SHOES
5. ALWAYS TIE BACK LONG HAIR
6. IDENTIFY THE LOCATION OF FIRE EXTINGUISHERS, FIRE BLANKETS, EYE BATH AND
EMERGENCY SHOWER IN THE LABORATORY
7. DO NOT EAT DRINK OR CHEW GUM AT ANY TIME IN THE LABORATORY
8. NEVER TASTE CHEMICALS
9. KEEP YOUR COATS, TEXTBOOKS, CALCULATORS AND OTHER OBJECTS OF VALUE OFF THE
WORKBENCH AND AWAY FROM THE WORKING AREA — MANY REAGENTS ARE CORROSIVE OR
LEAVE STAINS
10. ALWAYS WASH YOUR HANDS WITH SOAP AND WATER BEFORE LEAVING THE LABORATORY
11. REPORT ANY ACCIDENT, HOWEVER MINOR, TO THE INSTRUCTOR AT ONCE
12. THINK ABOUT WHAT YOU ARE DOING AT ALL TIMES
SAFETY PROCEDURES
Never enter a laboratory without a supervisor being present. Under no conditions are unauthorised or
unsupervised experiments to be performed.
Make sure that laboratory equipment is securely fastened — wobbly and loose equipment lead to
accidents.
Never leave an experiment unattended.
Never aim the opening of a test tube or flask at yourself or at anyone else.
After each laboratory period, use your sponge to wash your benchtop before leaving the laboratory.
Always clean any water on the bench top or the floor immediately.
Wash your hands thoroughly after carrying out the experiment.
2
HANDLING CHEMICALS
Always add acids to water; never add water to acids.
Use fume hoods whenever poisonous or irritating fumes or volatile solvents are handled.
Be very cautious when testing for odours. Using your hands, waft the gas towards your nose.
Never introduce glassware, such as medicine droppers or pipettes, into a reagent bottle — this will
contaminate the source of pure chemical. Pour a small amount of the reagent into one of your clean
flasks and use this to obtain dropper or pipette amounts.
Use distilled water for any experiments calling for water; tap water is used only for washing glassware.
It is not usually necessary to conduct an experiment using dry glassware. Drying glassware with paper
towel or a rag contaminates the glass. If it is necessary to use clean, dry glassware, wash it and dry it in
the oven.
Excess chemicals should never be returned to the reagent bottle – this will contaminate the contents of
the bottle. This is one of the most serious offences that can be committed in the laboratory. The excess
chemical should be shared with your classmates or disposed of as described in the section below.
WASTE DISPOSAL
Waste disposal is an environment problem that causes much concern to the chemical industry. In
Montreal, anything that goes down a drain eventually ends up in the St. Lawrence River. The municipality
removes solid wastes. Organic chemical wastes are stored in special containers for eventual incineration
or recycling. You can help reduce the amount of chemicals that must be disposed of by using only the
amounts recommended in the experimental procedures.
Use sinks to dispose of aqueous solutions only, and flush with plenty of tap water. Organic solvents
should be discarded in the container marked Organic Liquid Wastes that is found in a fume hood in the
laboratory. Solid waste must be disposed of in solid waste containers.
If you are unsure of how to dispose of any waste chemicals, consult you instructor.
Discard broken glass in the marked containers.
At the end of every experiment, wash the glassware used with soap and tap water then rinse with small
quantities of distilled water. Do not dry the glassware; it is acceptable to store wet glassware in your
drawer.
3
COMMON SYMBOLS INDICATING CHEMICAL HAZARDS
The following symbols identify some chemical hazards. They were adopted by the Canadian government
in 1988 as part of the Workplace Hazardous Materials Information System (WHMIS). Precautionary
measures must be taken when these symbols appear on chemical labels.
Symbol
Meaning
Compressed Gas
Flammable and Combustible Material
Oxidizing Material
Poisonous and Infectious Material
Immediate and serious toxic effects
Poisonous and Infectious Material
Other toxic effects
Poisonous and Infectious Material
Biohazardous Infectious Material
Corrosive Material
Dangerously Reactive Material
4
DAWSON Material Safety Data Sheet (MSDS)
Section 1-2: Chemical Product and Composition
Product Name: Acetone
Chemical Formula: C3H6O
CAS#: 67-64-1
Synonym: 2-Propanone; Dimethyl Ketone; Dimethylformaldehyde; Pyroacetic Acid
WHMIS CLASS:
B-2 D-2B
Section 3-4: Hazards Identification and First Aid Measures
EMERGENCY OVERVIEW
Appearance: clear, colorless liquid. Flash Point: −20 deg C.
Danger! Extremely flammable liquid and vapor. Vapor may cause flash fire. Causes eye irritation.
Breathing vapors may cause drowsiness and dizziness. Causes respiratory tract irritation. Aspiration
hazard if swallowed. Can enter lungs and cause damage. Prolonged or repeated contact may dry the skin
and cause irritation. Target Organs: Central nervous system, respiratory system, eyes, skin.
Potential Acute Health Effects:
Hazardous in case of skin contact (irritant), of eye contact (irritant), of ingestion, of inhalation.
Potential Chronic Health Effects:
Prolonged or repeated skin contact may cause dermatitis. Chronic inhalation may cause effects similar to
those of acute inhalation. Matsushita et al. exposed human volunteers 6 hours/day for 6 days at 500 ppm
acetone and found hematologic changes including significantly increased leukocyte and eosinophil counts
and decreased neutrophil phagocytic activity.
FIRST AID MEASURES
Eye Contact: Immediately flush eyes with running water for at least 15 minutes, keeping eyelids open.
Skin Contact: immediately flush skin with plenty of water. Cover the irritated skin with an emollient.
Inhalation: If inhaled, remove to fresh air. If not breathing, give artificial respiration. If breathing is difficult,
give oxygen. Get medical attention if symptoms appear.
Ingestion: Do NOT induce vomiting unless directed to do so by medical personnel.
Section 5: Fire and Explosion Data
Flammability of the Product: Flammable.
Flash Points: CLOSED CUP: -20°C (-4°F). OPEN CUP: -9°C (15.8°F).
Products of Combustion: These products are carbon oxides (CO, CO2).
Fire Hazards in Presence of Various Substances: Highly flammable in presence of open flames and
sparks, of heat.
Explosion Hazards in Presence of Various Substances: Slightly explosive in presence of open flames
and sparks, of oxidizing materials, of acids.
Fire Fighting Media and Instructions: Flammable liquid, soluble or dispersed in water. SMALL FIRE:
Use DRY chemical powder. LARGE FIRE: Use alcohol foam, water spray or fog.
Special Remarks on Fire Hazards: Vapor may travel considerable distance to source of ignition and
flash back.
Section 6: Accidental Release Measures
Small Spill: Dilute with water and mop up, or absorb with an inert dry material and place in an
appropriate waste disposal container.
5
DAWSON Material Safety Data Sheet (MSDS) cont.
Section 8: Personal Protective Equipment
Eyes: Wear chemical splash goggles.
Skin: Wear butyl rubber gloves, apron, and/or clothing.
Clothing: Wear appropriate protective clothing to prevent skin exposure.
Section 9: Physical & Chemical Properties, Stability and Reactivity Data
Volatility: Not available.
Conditions of Instability: Excess heat, ignition sources, exposure to moisture, air, or water, incompatible
materials.
Corrosivity: Non-corrosive in presence of glass.
Section 10: Stability and Reactivity
Chemical Stability: Stable at room temperature in closed containers under normal storage and handling
conditions.
Conditions to Avoid: High temperatures, ignition sources, confined spaces.
Incompatibilities with Other Materials: Strong oxidizing agents, strong reducing agents, strong bases,
nitric acid, hexachloromelamine, sulfur dichloride, potassium tert-butoxide.
Hazardous Decomposition Products: Carbon monoxide, carbon dioxide.
Hazardous Polymerization: Will not occur.
6
List of chemicals used in NYB
Volumetric Analysis of an Acid Solution
Oxalic acid dehydrate, H2C2O4·H2O,
Hydrochloric acid, HCl
Sodium hydroxide, NaOH,
Phenolphthalein indicator, C20H14O4
Colligative Properties
Cyclohexane, C6H12
Benzophenone, C13H10O
Camphor, C10H16O
Biphenyl, C12H10
p-dichlorobenzene, C6H4Cl2
The Rate Law in Chemical Kinetics
Potassium iodide, KI
Sodium thiosulfate, Na2S2O3
Ammonium peroxydisulfate, (NH4)2S2O8
Potassium chloride, KCl
ammonium sulphate, (NH4)2SO4
starch solution.
Activation Energy in Chemical Kinetics
Same chemicals as Experiment 3
iron-copper catalyst solution
Spectrophotometry
Potassium permanganate, KMnO4
Chemical Equilibrium I: The Equilibrium Constant
Iron(III)nitrate, Fe(NO3)3
Nitric acid, HNO3
Potassium thiocyanate, KSCN
Titration Curves
Propionic acid, C3H6O2
Formic acid, CH2O2
Sodium hydroxide, NaOH,
Hydrochloric acid, HCl
Chemical Equilibrium II: The Solubility Product
Potassium hydrogen tartrate, KHC4H4O6
Potassium nitrate, KNO3
Sodium hydroxide, NaOH,
Phenolphthalein indicator, C20H14O4
The Indirect Determination of a Thermodynamic Quantity
Potassium nitrate, KNO3
7
8
Volumetric Analysis of an Acid Solution
Objectives
Part A: A primary standard acid solution is prepared and used to standardize a NaOH solution.
Part B: The standardized NaOH solution is used to determine the concentration of a monoprotic acid
solution.
Introduction
Titration is a process of mixing measured volumes of reacting solutions until equivalent amounts of
reactants are consumed. The purpose of a titration is to determine the concentration of a solution. The
concentration of one of the reacting solutions is known. The equivalence point of a titration is the point at
which equivalent amounts of the reactants have reacted. In order to determine the equivalence point, an
indicator is added to one of the reacting solutions. Such an indicator, if properly selected, undergoes a
sharp color change at (or very near) the equivalence point. The point at which the color change occurs is
termed the end-point of the titration.
In this experiment, a solution of oxalic acid dihydrate, H 2C2O42H2O, of known concentration is used to
determine the concentration of a NaOH solution according to:
H2C2O4 (aq) + 2 NaOH (aq) → Na2C2O4 (aq) + 2 H2O (l)
(1)
In turn the NaOH solution will be used to determine the concentration of an unknown acid.
Procedure
Part A
Calculate the mass of oxalic acid dihydrate, H2C2O42H2O, needed to prepare 250 mL of 0.050 M
solution. Accurately weigh an amount of the acid close to that calculated. Calculate the exact molarity
using the amount weighed.
Transfer the solid acid, using a funnel and a washing bottle, into a clean 250.00 mL volumetric flask.
Dissolve the acid in a small quantity of distilled water. After the solid acid has dissolved, add distilled
water to the mark. Shake well.
Rinse a clean burette with the NaOH solution of unknown concentration and discard the rinsate. Fill the
burette with NaOH to below the zero mark. Make sure that there are no air bubbles trapped in the tip of
the burette. Clamp the burette. Read carefully the level of the NaOH in the burette to two decimal
places.
Rinse a small, clean, beaker (beaker size: 150 mL or less) with few milliliter of the oxalic acid solution and
discard the rinsate. Fill the beaker with the oxalic acid solution. Rinse a 10.00 mL pipette with a small
amount of oxalic acid solution (from the beaker) and discard the rinsate. Fill the pipette with the acid
solution to the mark and transfer the acid aliquot into a clean 250 mL Erlenmeyer flask. Add water to the
50 mL mark and add five drops of phenolphthalein indicator.
Place the flask under the burette and start titrating very slowly while swirling the flask continuously. When
the pink color starts to linger add the titrant (NaOH) drop by drop. When, upon the addition of one drop,
9
the pink color remains in the solution stop titrating. Record the final burette reading. Repeat the process
twice with 10.00 mL aliquots of the acid.
Part B:
Rinse and fill a small beaker with the acid of unknown concentration. Rinse and fill a 10.00 mL pipette
with the acid and titrate with the NaOH solution as described above. (Fill the burette with NaOH, transfer
the acid aliquot into a clean 250 mL Erlenmeyer flask. Add water to the 50 mL mark and add five drops of
phenolphthalein indicator.)
Repeat the process twice with 10.00 mL aliquots of the acid.
Calculations
1.
Calculate the number of moles of oxalic acid used in each titration.
2.
Calculate the number of moles of NaOH that reacted with the oxalic acid in each titration.
3.
Calculate the molarity of the NaOH solution.
4.
Calculate the number of moles of NaOH used in each titration with the acid of unknown
concentration.
5.
Calculate the number of moles of the acid consumed in each titration.
6.
Calculate the concentration of the acid.
10
Name:
Section:
Partner’s Name:
Date:
VOLUMETRIC ANALYSIS OF AN ACID SOLUTION
DATA SHEET
Part A: Standardization of NaOH
Stoichiometric Equation for the Acid-Base Reaction:
Mass of oxalic acid dihydrate, H2C2O42H2O, g
__________________
Number of moles of H2C2O42H2O, mol
__________________
Volume of the volumetric flask, mL
__________________
Concentration of H2C2O42H2O solution, molL−1
__________________
Titration 1
Titration 2
Titration 3
Final burette reading, mL
_______
_______
_______
Initial burette reading, mL
_______
_______
_______
Volume of NaOH delivered, mL
_______
_______
_______
Volume of oxalic acid used, mL
_______
_______
_______
Moles of oxalic acid used, mol
_______
_______
_______
Moles of NaOH consumed, mol
_______
_______
_______
Concentration of NaOH , mol/L
_______
_______
_______
Average concentration of NaOH, molL−1
_______
11
Part B: Determination of the unknown concentration of an HCl Solution
Stoichiometric Equation for the Acid-Base Reaction:
Titration 1
Titration 2
Titration 3
Final burette reading, mL
_______
_______
_______
Initial burette reading, mL
_______
_______
_______
Volume of NaOH delivered, mL
_______
_______
_______
Moles of NaOH used, mol
_______
_______
_______
Moles of the acid used, mol
_______
_______
_______
Volume of the acid consumed, mL
_______
_______
_______
Concentration of the acid, molL−1
_______
_______
_______
Average concentration of the acid, molL−1
12
_______
Colligative Properties
Objective
Colligative properties of solutions depend on the quantity of solute dissolved in the solvent rather than the
identity of the solute. The phenomenon of freezing point lowering will be examined quantitatively as an
example of a colligative property.
Introduction
When a solute is dissolved in a solvent, the properties of the solvent are changed by the presence of the
solute. The magnitude of the change generally is proportional to the amount of solute added. Some
properties of the solvent are changed only by the number of solute particles present, without regard to the
particular chemical nature of the solute. Such properties are called colligative properties of the solution.
Colligative properties include the changes in vapour pressure, boiling point, freezing point and the
phenomenon of osmotic pressure.
If a nonvolatile solute is added to a volatile solvent (such as water), the amount of solvent molecules that
can escape from the surface of the liquid at a given temperature is lowered compared to the situation
where only pure solvent is present. The vapour pressure above such a solution will thus be lower than
the vapour pressure above a sample of the pure solvent under the same conditions. Molecules of the
nonvolatile solute physically block the surface of the solvent, thereby preventing as many molecules from
evaporating. This results in an increase in the boiling temperature of the solution as well as a decrease in
the freezing point.
In this experiment, the freezing points of three substances will be measured:
1. pure cyclohexane, C6H12;
2. a solution composed of cyclohexane and benzophenone, C13H10O; and
3. a solution of an unknown solid dissolved in cyclohexane.
The freezing point of the third solution will be used to determine the molar mass of the unknown solid.
The decrease in the freezing point, T, when a nonvolatile, nonionizing solute is dissolved in a solvent is
proportional to the molal concentration, m, of the solution,
∆𝑇 = 𝐾𝑓 ∙ 𝑚 .
(1)
∆𝑇 = 𝑇° − 𝑇
(2)
Here,
T • is the freezing temperature of pure cyclohexane, and T is the freezing temperature of the
benzophenone-cyclohexane solution. The freezing point depression constant, Kf, is constant for a given
solvent and establishes the number of degrees that the freezing point will be lowered when one mole of
solute is dissolved in one kilogram of solvent. The molal concentration is defined to be
𝑚𝑜𝑙𝑎𝑙𝑖𝑡𝑦, 𝑚 =
𝑛𝑢𝑚𝑏𝑒𝑟 𝑜𝑓 𝑚𝑜𝑙𝑒𝑠 𝑠𝑜𝑙𝑢𝑡𝑒
𝑘𝑖𝑙𝑜𝑔𝑟𝑎𝑚𝑠 𝑠𝑜𝑙𝑣𝑒𝑛𝑡
(3)
13
In Parts A and B the freezing points of pure cyclohexane and of a solution of benzophenone in
cyclohexane are measured respectively; the decrease in the freezing temperature, T, between the two
solutions can then be obtained. From known masses of benzophenone and cyclohexane used, the
molality of the solution can be calculated using equation (1). Together, these quantities provide the value
of Kf.
In Part C of this experiment, the freezing temperature of a solution of an unknown, non-ionizing solute
dissolved in cyclohexane will be measured. The known masses of solute and solvent, the measured
freezing temperature of pure cyclohexane and the previously determined value of Kf will then be used to
obtain the molar mass of the unknown solute using equations (1) and (3).
Procedure
Part A: Determination of the Freezing Point of Pure Solvent
Insert a solid rubber stopper into a clean, dry 200 mL test tube and place the stoppered test tube in
a clean 150 mL beaker. Weigh this apparatus on a balance. Put aside the beaker and the stopper
as they will be used for weighing again.
In the fume-hood, pour approximately 25 mL of solvent into the test tube.
Cap the test tube with a cork fitted with the temperature sensor and a copper wire stirrer. Make
sure the temperature sensor is connected to the CH1 input of the LabQuest mini interface and that
the LabQuest interface is active (green LED on).
Double click on the Logger Pro 3.8.3 icon to open the program.
Click on the File menu and choose Open. Double click on the Chemistry of Solutions folder then
double click on the 02_colligative file. If the sensor confirmation box opens, select the DirectConnect Temp sensor and press connect.
Fill a 400-mL beaker with approximately 300 mL of a tap water-ice mixture.
Immerse the test tube in the ice-water bath and begin stirring the solution immediately with the
copper wire stirrer. Click on the Logger Pro Collect button.. The contents of the test tube must be
stirred continuously for the duration of the experiment. After 10 min, data collection will
automatically stop.
Remove the test tube from its ice bath and replace the cork with the rubber stopper. Do not
discard this solvent. It will be used in Part B.
Select the portion of the graph containing points immediately at the start of the experiment where
the temperature drops quickly. In this region, the data should be reasonably linear. Click on the
regression line button on the Logger Pro Toolbar.
In a similar fashion, obtain the regression line for the linear portion of the graph towards the end of
the data range. Drag the regression line text boxes to convenient locations on the graph.
Click on the Analyze menu item and choose Interpolate; a third box will open. Move the mouse
cursor in the Graph Window so that the two circles appearing on the regression lines overlap.
Monitor the two temperature values. When the two values are identical (to within one decimal
place) record the temperature on the data sheet.
Save the Logger Pro file to a flash drive after providing it with an appropriate name. Close Logger
Pro.
14
Part B: Determination of the Freezing Point of a Benzophenone/Solvent Solution
Warm the solvent (from Part A) in your hands until it melts. Weigh the solvent using the same
150 mL beaker.
Weigh approximately 0.4 g benzophenone to three decimal places. Transfer the benzophenone to
the test tube containing the melted solvent and mix thoroughly until the solute is fully dissolved.
Cover with the rubber stopper and weigh the test tube using the same 150 mL beaker as before.
Replace the rubber stopper with the cork stopper, temperature probe and stirrer.
Double click on the Logger Pro icon to open the program.
Click on the File menu and choose Open. Double click on the Chemistry of Solutions folder then
double click on the 02_colligative file.
Monitor the temperature of the solution on the Logger Pro status bar. It should not differ from room
temperature by more than 1C.
Immerse the test tube into the ice-water bath and click on Collect. Stir the solution continuously
during data collection.
When data collection stops, analyze the data as in Part A to obtain the final solution temperature.
Record this temperature on the data sheet.
Save the Logger Pro file to a Flash drive after providing the file with a suitable name. Exit Logger
Pro.
Discard the contents of the test tube in the container labeled Organic Liquid Waste, wash it
thoroughly with soap and place it in the oven to dry.
Part C: Determination of the Freezing Point of a Solution of an Unknown Solid in Cyclohexane
Insert a solid rubber stopper into a clean, dry 200 mL test tube and place them in the 150 mL
beaker. Weigh this apparatus on a balance.
In the fume-hood, pour approximately 25 mL of cyclohexane into the test tube and recap.
Determine the freezing point of cyclohexane as in Part A. Weigh the test tube with the melted
cyclohexane.
Obtain a sample of an unknown solid. Measure out a mass of the solid within the range listed on
the bottle and transfer this quantity into the test tube containing cyclohexane.
Cap the solution and weigh the test tube in the 150 mL beaker as before.
Replace the solid rubber stopper with one of cork fitted with the temperature sensor and a copper
wire stirrer.
Stir the mixture, warming it if necessary, in order to dissolve all the solid.
Double click on the Logger Pro icon to open the program.
15
Click on the File menu and choose Open. Double click on the Chemistry of Solutions folder then
double click on the 02_colligative file.
Monitor the temperature of the solution on the Logger Pro status bar. It should not differ from room
temperature by more than 1C.
Immerse the test tube into the ice-water bath and click on Collect. Stir the solution continuously
during data collection.
When data collection stops, analyze the data as in Part A to obtain the final solution temperature.
Record this temperature on the data sheet.
Save the Logger Pro file to a Flash drive after providing the file with a suitable name. Exit Logger
Pro.
Discard the contents of the test tube in the container labeled Organic Liquid Waste, wash it
thoroughly with soap and place it in the oven to dry.
Log off the network.
Treatment of Data
1. Calculate the freezing point depressions for the benzophenone solution and the solution of the
unknown solute.
2. Calculate the molality of the benzophenone solution.
3. Calculate Kf for
20.2 kgCmol–1.
cyclohexane
and
compare
the
result
with
the
literature
value
of
4. Using the measured value of Kf, calculate the molar mass of the unknown solute use in Part C.
Pre-Laboratory Questions
1.
Define the term colligative property.
2.
Why does the presence of a solute lower the freezing temperature of a solvent?
3.
If the presence of 2.00 g of an unknown non-ionizing substance lowers the freezing temperature of
a 10.0 g sample of benzene by 6.33C, calculate the molar mass of the substance. Kf for benzene
may be obtained from your textbook.
4.
Why do ionic substances have a greater effect on the freezing and boiling temperatures of solvents
than do non-ionic substances for a given number of moles?
5.
Calculate the molality, molarity and percent by mass of a solution of ethanol, C 2H5OH, and water in
which the mol fraction of ethanol is 0.0820. The density of the solution is 0.978 g/mL.
16
Name
Section
Date
COLLIGATIVE PROPERTIES
DATA SHEET
Part A:
Data for Pure Cyclohexane
Mass of empty test tube, stopper & beaker, g
Freezing Temperature of cyclohexane, C
Part B:
Data for the Benzophenone/Cyclohexane Solution
Mass of test tube, stopper, beaker & cyclohexane, g
Mass of cyclohexane, g
Mass of test tube, stopper, beaker, benzophenone & cyclohexane, g
Mass of benzophenone, g
Freezing Temperature of benzophenone/cyclohexane solution, C
Molality of benzophenone/cyclohexane solution, molkg−1
Kf for cyclohexane, kgCmol−1
% error
17
Name
Section
Date
COLLIGATIVE PROPERTIES
DATA SHEET
Part A:
Data for Pure Cyclohexane
Mass of empty test tube, stopper & beaker, g
Freezing Temperature of cyclohexane, C
Kf (from Parts A and B), Ckgmol–1
Part C:
Data for the Unknown Solute/Cyclohexane Solution
Unknown Number:
Mass of test tube, stopper, beaker, & cyclohexane, g
Mass of test tube, stopper, beaker, & unknown solute/cyclohexane solution, g
Mass of cyclohexane, g
Mass of unknown solute, g
Freezing Temperature of unknown solute/cyclohexane solution, C
Molar mass of unknown solute, gmol−1
18
Spectrophotometry
Objective
The relation between absorbance and the concentration of KMnO 4 solutions is determined. This
relationship is used to measure the concentration of an unknown solution of KMnO 4.
Introduction
All molecules or ions absorb light. The quantity of light absorbed depends on both the wavelength of the
incoming light and the concentration of the absorbing species.
The absorption of light by a sample is measured by comparing the intensity of incident light on the
sample, I0, with the intensity of this light, I, after passage through the sample. Thus the percent
transmittance, T, is
T=
𝐼
𝐼𝑜
∙ 100%
(1)
The transmittance of a sample is related to its absorbance, A, defined by
A = log(100/T)
(2)
An instrument, which provides either T or A is known as a spectrophotometer.
The container (known as the cell), the solvent as well as the solute all normally absorb light. To correct
for this, the intensity of light transmitted by the solution is compared to the light transmitted by the solvent.
The solvent is known as the reference solution. Usually the spectrophotometer is adjusted so that the
transmittance of the reference solution reads 100%. In this case, the absorbance of the sample solution
is directly proportional to the molar concentration of the solute, c, according to Beer’s Law,
A = kc.
(3)
Various solutions of known KMnO4 concentrations are prepared and their absorbencies measured at a
pre-determined wavelength (565 nm or green position). The results are graphed in a calibration curve,
A vs. c. Finally the absorbance of an unknown sample of KMnO4 is measured and its concentration
obtained by interpolating in the calibration curve.
Procedure
Obtain 20 mL of the KMnO4 stock solution. Record its concentration.
Pipette 10.00 mL of this stock solution in a 250-mL volumetric flask. Dilute carefully, with distilled water,
to exactly to 250.0 mL (i.e. to the mark). Shake well to mix the solution. This solution is known as the
primary standard solution. Do not discard the primary solution until the end of the experiment.
19
Into each of five clean, dry beakers add the following quantities of distilled water using the 10 mL pipette
1.
2.
3.
4.
5.
10.00 mL distilled water
20.00 mL distilled water
30.00 mL distilled water
40.00 mL distilled water
60.00 mL distilled water
Rinse out a beaker with small quantities of the primary solution. Rinse the 10-mL pipette and deliver
10.00 mL volumes of the primary solution into each of the five beakers containing the distilled water
previously prepared. Do not pipette directly from the volumetric flask containing the primary (to avoid any
possibility of contaminating the primary solution). Stir to mix the solutions.
Calculate the concentrations of each of these solutions and record these values.
important as the experiment CANNOT be conducted without these values.
This is
Two special cells (cuvettes) have been provided. The following protocols for using cuvettes should be
observed at all times:
• Always hold cuvettes by the top edges of their ribbed sides
• Under no circumstances should they ever be more than three-quarters filled
• Make sure there are no bubbles in the solution
• Always cap cuvettes before insertion into the cell-well of the colorimeter
Fill one of the cuvettes with distilled water and cap it. This will be the reference solution.
Rinse the other cuvette with solution 5, discard the rinsate and fill the rinsed cuvette with the solution.
Make sure the LabQuest Mini interface is on (green led) and Vernier’s Colorimeter is connected to it at
the CH1 input. Double click on the Logger Pro 3.8.3 icon on the computer desktop
Open the Chemistry of Solutions folder and then double click on the file 05_spectrophotometry. This
experiment file is configured for prompted data acquisition, with the prompt being the concentration
values calculated earlier.
The colorimeter must now be calibrated. Wipe the cuvette containing the reference solution (water) with
tissue and place it in the cell well of the colorimeter. Make sure to position the cuvette so that one of its
clear faces is directed toward the white reference mark on the cell well. Close the cell well. On the
calorimeter, use the two arrows to select the light wavelength to 565 nm. Press the CAL button on the
calorimeter and a red light will blink 10 times. The colorimeter has been calibrated.
Remove the reference solution from the cell well and replace it with the cuvette containing solution 5 (the
most diluted one) after wiping its clear sides with tissue.
In the Logger Pro toolbar click on Collect. Monitor the transmittance on the status bar. When it has
stabilized, click on Keep; an entry window will appear. Enter the concentration of KMnO 4 (calculated
earlier for solution 5) in the box and press OK.
Remove the cuvette from the cell well and discard its contents. Rinse the cell with solution 4 and
measure its absorbance. Record its concentration by pressing Keep.
The button STOP should never be pressed throughout the experiment.
Repeat the previous step with each of the solutions 3, 2 and 1.
If an error has been made in entering one of the concentration values, it can be corrected at the end of
the acquisition process. Click on the cell containing the incorrect concentration in the Table Window.
The cell and its contents will be highlighted; make the correction and press Enter on the keyboard.
20
When all of the readings have been taken, click on Stop.
Obtain an unknown solution of KMNO4. Rinse a cuvette with the unknown solution, fill it, cap it and insert
it in the cell well. Read and record its absorbance on your data sheet.
Click in the Graph Window and then double click on the title “Absorbance vs. molar concentration”.
Select the Graph options tab and provide the graph with a suitable name in the dialog box that appears
then click OK.
Click on the regression line button to obtain the best straight line that fits the data.
The concentration of an unknown solution of KMnO 4 will be obtained by measuring its absorbance and
obtaining its concentration by interpolating in the regression line.
Click on the Analyze menu item and choose Interpolate; a new box will open. Move the vertical
interpolate line until it reaches the absorbance value of the unknown. Record the corresponding
concentration on your data sheet.
This information should be recorded on the graph. Click on the menu item Insert then choose Text. In
the dialogue box that appears, record the unknown number and its concentration then press Enter. Drag
the annotation text box to a suitable location on the graph window.
Save the graph on your flash drive, set up the graph for printing then print (press “Ctrl” and “P” together).
Remove the cuvette from the cell well and rinse both cuvettes thoroughly with distilled water. Place them
in the beaker beside the spectrophotometer. Do not put them in your drawer. Switch the colorimeter
wavelength selector button to the Off position.
Log off the network.
21
22
Name:
Section:
Partner’s Name:
Date:
SPECTROPHOTOMETRY
DATA SHEET
Concentration of the stock solution, M
0.020
Concentration of the primary solution, M
0.00080
Concentration of the solution 1, M
0.00040
Concentration of the solution 2, M
0.00027
Concentration of the solution 3, M
0.00020
Concentration of the solution 4, M
0.00016
Concentration of the solution 5, M
0.00011
Unknown Number:
11
Absorbance of the unknown
0.175
Concentration of the unknown (from graph), M 0.00023
Calculations of the Concentrations of the KMnO4 Solutions
1.
Concentration of the Primary
2.
Concentration of Solution 1
3.
Concentration of Solution 2
4.
Concentration of Solution 3
5.
Concentration of Solution 4
6.
Concentration of Solution 5
23
24
Chemical Equilibrium I: The Equilibrium Constant
Objective
This experiment uses the technique of spectrophotometry to determine the equilibrium constant for the
reaction
Fe3+ (aq) + SCN− (aq)
Fe(SCN)2+ (aq).
Introduction
The equilibrium constant, K, for the reaction
Fe3+ (aq) + SCN− (aq)
Fe(SCN)2+ (aq).
(1)
is given by
K=
[Fe(SCN)2+ ]
[Fe3+ ][SCN− ]
(2)
This equilibrium constant is sometimes referred to as the formation constant of the Fe(SCN) 2+ complex.
The Fe(SCN)2+ complex absorbs light in the blue and blue-green region of the visible spectrum (in the
range 430 nm – 490 nm) and transmits light in the remainder of the wavelength range, namely in the
orange and red end at 500 nm to 700 nm. Thus solutions of Fe(SCN) 2+ have a red-orange colour, having
an absorbance maximum at a wavelength of 470 nm (blue). Since Fe3+ and SCN– are both colourless
(or very nearly so) in aqueous solution, it is evident that the equilibrium concentration of Fe(SCN) 2+ can
be measured spectrophotometrically by measuring the intensity of the solution colour and applying Beer’s
Law,
A = k[Fe(SCN)2+]eq
(3)
Solutions are prepared with known initial concentrations, [Fe 3+]0 and [SCN−]0. From the stoichiometry of
reaction (1), the equilibrium concentrations of Fe3+ and SCN- can be obtained according to
[Fe3+]eq = [Fe3+]0 − [Fe(SCN)2+]eq
(4)
[SCN−]eq = [SCN−]0 − [Fe(SCN)2+]eq
(5)
Substitution of the concentrations in equation (2) provides the formation constant for Fe(SCN)2+.
The equilibrium is carried out in the presence of 0.50 M HNO 3 in order to maintain a constant ionic
strength and to minimize hydrolysis by Fe3+. Both of these factors affect the value of the formation
constant of Fe(SCN)2+.
25
Procedure
1.
The Calibration Curve
Take eight clean, dry medium sized test tubes and number them from 1 to 8.
Take these test tubes (on a test tube rack) to the side bench marked A. Record the concentrations
of Fe(NO3)3, KSCN and HNO3 provided in the burettes at bench A.
In each of the eight labeled test tubes, mix the volumes of the Fe(NO 3)3, KSCN and HNO3 specified
in Table 1 below. Shake them well to mix or stir with a glass rod. It is normal that the liquid level
between two test tubes is different since the diameters of test tubes are often quite variable.
Table 1:
Test Tube No.
1
2
3
4
5
6
7
8
Solution Volumes for the Calibration Curve
Volume of Fe(NO3)3, mL
3.00
3.00
3.00
3.00
3.00
3.00
3.00
3.00
Volume of KSCN, mL
1.00
2.00
3.00
4.00
5.00
6.00
7.00
0.00
Volume of HNO3, mL
6.00
5.00
4.00
3.00
2.00
1.00
0.00
7.00
Test tube 8 is the reference solution
Calculate the concentrations of SCN− in these eight solutions and record their values. This
is important as the experiment cannot be conducted without these values (see the section
on Calculations). (note: [FeSCN]2+eq ≈ [SCN–]o after reaction in this part of the experiment)
Record the temperature of each of the eight solutions. The variation in temperature should not
exceed 2C. Make sure the thermometer is rinsed and dry before inserting it in the test tube.
Two cuvettes have been provided. These cuvettes should be handled only by the top edges of the
ribbed sides. Fill one of them with the reference solution (8). Remember, do not fill the cuvettes
more than three-quarters full. Cap the cuvette.
Rinse the other cuvette with the solution in test tube 1, discard the rinsate, fill the rinsed cuvette
with the solution and cap it.
Log onto the network.
Make sure the Vernier’s Colorimeter sensor is connected to the CH1 input of the active LabQuest
mini interface (green LED on). Double click on the Logger Pro 3.8.3 icon on the computer desktop
Open the Chemistry of Solutions folder and then double click on the file 06_equilibrium
constant. This experiment file is configured for prompted data acquisition, with the prompt being
the concentration values calculated earlier.
The colorimeter must now be calibrated. Wipe the cuvette containing the reference solution (from
test tube 8) with tissue and place it in the cell well of the colorimeter.
Remember, always cap the cuvette before inserting it into the cell well of the colorimeter.
26
Make sure to position the cuvette so that one of its clear faces is directed toward the white
reference mark on the cell well. Close the cell well. On the calorimeter, use the two arrows to
select the light wavelength to 470 nm.
Press the CAL button on the calorimeter and a red light will blink 10 times. The colorimeter has
been calibrated.
Remove the reference solution from the cell well and replace it with the cuvette containing solution
1 after wiping it with tissue.
In the Logger Pro toolbar, click on Collect. Monitor the absorbance on the status bar. When it has
stabilized, click on Keep. Enter the appropriate concentration of SCN − in the entry box that
appears. The button STOP should never be pressed throughout the experiment.
Remove the cuvette from the cell well and discard its contents. Rinse the cell twice with solution 2
and measure its transmittance. Repeat with the remaining solutions.
After recording the absorbance of the last solution, click on Stop.
Click in the Graph Window and then double click on the title “Absorbance vs. molar concentration”.
Select the Graph options tab and provide the graph with a suitable name in the dialog box that
appears then click OK.
Click on the regression line button to obtain the best straight line that fits the data. Note that the
regression line does not pass through the origin.
Print the graph and save the file to a Flash drive.
2.
Determination of the Equilibrium Concentration of Fe(SCN)2+
Take six clean, dry medium sized test tubes and number them from 1 to 6. Take them in a test
tube rack to the side bench with the solutions marked B.
Record the concentrations of the solutions in the burettes at bench B.
Prepare solutions by mixing the volumes of Fe(NO3)3, KSCN and HNO3 specified in Table 2 below.
Table 2:
Solution Volumes for the Determination of the Formation Constant of FeSCN 2+
Test Tube No.
1
2
3
4
5
6
Volume of Fe(NO3)3, mL
5.00
5.00
5.00
5.00
5.00
5.00
Volume of KSCN, mL
1.00
2.00
3.00
4.00
5.00
0.00
Volume of HNO3, mL
4.00
3.00
2.00
1.00
0.00
5.00
Test tube 6 is the reference solution
Record the temperature of each of the six solutions. The variation in temperature should not
exceed 2C of those of Part 1.
The button “Collect” SHOULD NEVER BE PRESSED at this part of the experiment. If so, the
calibration curve of part 1 will be lost.
27
Recalibrate the colorimeter as in Part 1 using the reference solution of part 2: Test tube 6.
Absorbance reading is made again at the 470 nm position.
Read the absorbances from the status bar of each of the remaining solutions and record the values
in the laboratory data sheet. Proceed from the lowest concentration to the highest (from test tube 1
to 5).
The calibration curve obtained in Part 1 will now be used to obtain the equilibrium concentrations of
Fe(SCN)2+ in each of the solutions.
Click on the Analyze menu item and choose Interpolate. Move the mouse on the interpolate line
until it reaches the absorbance values of each of the solutions recorded in your lab data sheet.
Record the concentrations corresponding to these values (from the x-value in the dialog box).
Remove the cuvette from the cell well and rinse both cuvettes thoroughly with distilled water. Place
them in the beaker beside the spectrophotometer. Do not put them in your drawer.
Log off the network.
Calculations
The Calibration Curve
1. Noting that the total volume for each solution is always 10.00 mL, calculate the [SCN -]0 for each of
the test tubes. These concentrations are equal to [Fe(SCN)2+]eq for each of the solutions.
Determination of the Equilibrium Concentration of Fe(SCN)2+
2. Calculate [Fe3+]o and [SCN−]o for each of the solutions.
3. Determine [Fe3+]eq and [SCN−]eq according to equations (4) and (5).
4. Calculate the value for K for each solution and the average.
28
Name:
Section:
Partner’s name:
Date
CHEMICAL EQUILIBRIUM I: THE EQUILIBRIUM CONSTANT DATA SHEET
Fe3+ (aq) + SCN− (aq)
Part 1 (SOLUTIONS A) :
Fe(SCN)2+ (aq).
Determination of the Calibration Curve
Reagent Concentrations
[Fe(SCN)2+]eq, M
Solution
VKSCN, mL
0.8
1
1.00
[KSCN], M
2.0E-4
2
2.00
2.0E-5
4.0E-5
[HNO3], M
0.5
3
3.00
6.0E-5
4
4.00
8.0E-5
5
5.00
1.0E-4
6
6.00
1.2E-4
7
7.00
1.4E-4
8
0.00
0.0
[Fe(NO3)3], M
Temperature, C
24.0
Determination of the Formation Constant for Fe(SCN)2+
Part 2 (SOLUTIONS B):
Reagent Concentrations
[Fe(NO3)3], M
2.0E-3
[KSCN], M
2.0E-3
[HNO3], M
0.5
Solution Absorbance
Temperature, C
[Fe(SCN)2+]eq, M [Fe3+]0, M
[SCN−]0, M
24.0
[Fe3+]eq, M
[SCN−]eq, M
K
1
0.070
1.36E-5
1.0E-3
2.0E-4
9.9E-4
1.9E-4
7.2E1
2
0.156
3.63E-5
1.0E-3
4.0E-4
9.7E-4
3.6E-4
1.0E2
3
0.256
6.12E-5
1.0E-3
6.0E-4
9.4E-4
5.4E-4
1.2E2
4
0.325
7.86E-5
1.0E-3
9.2E-4
7.2E-4
1.2E2
5
0.393
9.70E-5
8.0E-4
1.0E-3
1.0E-3
9.0E-4
9.0E-4
1.2E2
Average K : ______________________
1.0E2
29
30
Titration Curves
Objectives
A strong base, NaOH is used to obtain the titration curves of a strong acid and of a weak acid. These
curves will be used to obtain the initial concentrations of both acids as well as the pK a of the weak acid.
Introduction
A titration involves the addition of a measured volume of one solution, the titrant, from a burette to
another solution, the analyte. The reaction of the titrant with the analyte is usually known, therefore the
technique of titration is commonly used to determine solution concentrations.
A plot of pH vs. the volume of added titrant is known as a titration curve. In this experiment, the titration
curves of two acids, HCl, and a weak acid, denoted by HA, are obtained. The titrant in both cases is the
solution of a strong base, NaOH, of known concentration. The schematic titration curve for a weak acid
by a strong base is presented in Fig. 1 below. The strong acid-strong base titration curve is qualitatively
similar to that depicted in Fig. 1, although the jump in pH near the equivalence point is markedly sharper
for strong acid-strong base titration curves.
Figure 1:
Titration Curve for a Weak Acid by Strong Base
pH
equivalence
point
pH
7
pKa
0.5Veq
Veq
Volume of Strong Base Added (mL)
The titration of a weak acid, generically denoted by HA, by a strong base proceeds according to the
reaction,
HA (aq) + OH– (aq) → A– (aq) + H2O (l),
(1)
31
The addition of NaOH (having the form OH–), drives the reaction in equation (1) until all of the HA is
converted to its conjugate base, A–. This is known as the equivalence point. At the equivalence point the
hydrolysis of A– (the reverse of equation (1)) becomes the principal chemical reaction. At the equivalence
point,
nHA consumed = nNaOH added
(2)
Here, nNaOH is the number of moles NaOH delivered at the equivalence point, and is calculated by,
nNaOH = [NaOH]Veq,
(3)
and nHA is the number of moles of weak acid, HA, initially present. V eq is the volume of base required to
attain the equivalence point (see Fig. 1), and [NaOH]is the initial concentration of the strong base.
The conjugate base of a strong acid is incapable of hydrolysis, thus at the equivalence point the pH = 7.
In the case of a weak acid, the conjugate base, A–, will engage in hydrolysis, resulting in a pH > 7 at the
equivalence point.
At the half-equivalence point of a weak acid-strong-base titration curve,
V = ½ Veq,
(4)
the number of moles of NaOH added is exactly ½nHA. Consequently, one half of the original
concentration of HA has been converted to A –, at least initially. Thus at the half-equivalence point, the
equilibrium acid/conjugate base concentrations are approximately equal.
[HA] ≅ [A− ]
(5)
pH = pKa.
(6)
and therefore
Procedure
PART 1:
TITRATION CURVE FOR A STRONG ACID BY A STRONG BASE
Record the concentration of the standard NaOH solution.
Rinse and fill a burette to the 0.00 mL mark with this standard NaOH solution and clamp the filled burette
to a ring stand.
Rinse a 10 mL pipette with small quantities of the HCl solution, discarding the rinsates.
Pipette 10.00 mL of the HCl solution into a 250 mL beaker and add approximately 150 mL distilled water.
Make sure the pH electrode is attached to the LabQuest Mini interface with the green led on.
Log onto the computer network.
Double click on the Logger Pro icon. Open the Chemistry of Solutions folder and then the file
07_titration. If the window “sensor confirmation” opens, make sure that the selected sensor is pH and
click connect.
Rinse the pH sensor with distilled water. Carefully insert the pH sensor into the pH = 4 buffer solution,
allowing about 2 cm clearance above the bottom of the beaker.
Click on Experiment, choose Calibrate… then click on the LabQuest Mini: 1 CH1: pH tab.
32
Click on Calibrate Now. The pH value in the box is not active on calibration. Wait for 10 seconds then
set Value 1 to 4.0. Click on Keep.
Rinse the pH sensor with distilled water. Carefully insert the pH sensor into the pH = 7 buffer solution.
When the second reading stabilizes, set Reading 2 to 7.0. Click on Keep then click on DONE.
Carefully insert the pH sensor into the acid solution prepared previously.
Click on Collect. Once the pH reading on the status bar has stabilized, click on Keep and enter the
volume of NaOH added in the entry box. You should always write the value you read on the burette. It
should start from 0.00 and increase to 50.00.
Add enough of the NaOH titrant until the pH of the solution has been raised by about 0.15 units. When
the pH reading has stabilized (10 to 15 seconds), click on Keep, and enter the number read on the
burette (i.e. total volume of NaOH added).
Continue adding NaOH solution in increments that raise the pH about 0.15 units and enter the burette
reading after each addition.
When a pH value changes rapidly, add the NaOH dropwise. Enter a new burette reading after each
addition.
After a pH value of approximately 10 units has been reached, beginning adding larger increments of
NaOH to raise the pH by about 0.15 units. Enter a new burette reading after each addition.
Continue this procedure until a pH value of 12 has been attained. Click on Stop. Rinse off the pH
electrode with distilled water and immerse its tip in the water solution.
One way to obtain the coordinates of the equivalence point is to measure the slopes of the titration curves
as a function of volume of titrant added. The point at which the slope has its maximum value defines the
equivalence point.
Click on the file menu item Data and choose New Calculated Column. In
the dialogue box that appears, type the name: “Slope”. In the equation
section select Function> , Calculus and Derivative. In the equation box you
will select derivative(). On the section Variables (columns)> select pH and
press Done. Finally, on the graph, click on pH and choose slope. Your graph
should reflect the plot of Slope vs. Volume. Autoscale the graph by clicking
on the autoscale button. Using the Examine feature, (i.e. click on the
appropriate Toolbar button) to identify the coordinates of the equivalence
point.
Autoscale button
Examine button
Print the Graph window and then save the file to a Flash drive. Close the file.
PART 2:
TITRATION CURVE FOR A WEAK ACID BY A STRONG BASE
Re-open the 7_titration file and repeat the procedure for the titration of the weak acid. Do not forget to
record the unknown number of the weak acid. Once the equivalence point has been identified, use the
Examine feature to determine the pH of the half-equivalence point.
Print and save the graph to a Flash drive.
Log off the network.
33
Calculations
1. Identify the equivalence points for each acid.
2. Calculate the initial acid concentrations.
3. Identify the half equivalence point for the weak acid and report its pKa.
34
Name:
Section:
Partner’s Name:
Date:
TITRATION CURVES
DATA SHEET
Part 1: Strong Acid-Strong Base Titration
[NaOH], M
VHCl , mL
pH at equivalence point
VNaOH at equivalence point, mL
[HCl]0, M
Part 2: Weak Acid-Strong Base Titration
Unknown No:
[NaOH], M
VHA, mL
Initial pH
pH at equivalence point
VNaOH at equivalence point, mL
[HA]o , M
pKa
35
36
Chemical Equilibrium II: The Solubility Product
Objective
The solubility product of potassium hydrogen tartrate, KHC4H4O6, is measured and the effect of a
common ion on its solubility is examined.
Introduction
Potassium hydrogen tartrate is sparingly soluble in water and dissolves to produce potassium and
hydrogen tartrate ions according to
H2O
K+ (aq) +
KHC4H4O6 (s)
HC 4H4O6− (aq)
(1)
For this process, the equilibrium constant, known as the solubility product, is denoted by Ksp in the
expression
−
Ksp = [K+][ HC 4H4O6 ]
(2)
Here, the ion concentrations are those values for the saturated solution.
The hydrogen tartrate ion,
HC 4H4O6− , has an ionizable hydrogen ion. Thus its concentration can be
determined by titration with a strong base,
HC 4H4O6− (aq) + OH– → C4H4O62 − (aq) + H2O (l).
From the concentration (solubility) of
(3)
HC4H4O6− thus determined and the stoichiometry of the dissociation
equilibrium (1), the solubility product may be calculated.
In this manner, the solubility of KHC 4H4O6 in pure water and in four KNO3 solutions ranging in
concentration from 0.010 M to 0.040 M will be determined in order to verify that K sp is effectively
independent of the presence of other species and to explore the effect of a common ion, K +, on the
solubility of KHC4H4O6.
Procedure
Wash a burette then rinse and fill it with the standard NaOH solution. Record the concentration of the
NaOH solution.
Measure approximately 100 mL distilled water in a clean beaker and add to it about 1 g KHC 4H4O6 Stir the
solution vigorously for 5 minutes.
Record the temperature of the solution.
37
Filter this solution. Since the solution is saturated, it is essential that the filter paper, funnel and collection
beaker be dry prior to filtration.
Decant (pour) some of the filtrate (filtered solution) into another dry beaker and rinse a 10 mL pipette with
the filtrate. Discard the rinsate and the contents of the secondary (second beaker) beaker.
Transfer more filtrate into the secondary beaker and pipette 10.00 mL of the filtrate into a clean 250 mL
Erlenmeyer flask.
Add about 50 mL distilled water, 5-8 drops of phenolphthalein and titrate with the standard NaOH solution
to the faint pink end-point.
Record the volume of NaOH used in the titration.
Record the concentrations of the four KNO3 solutions.
Obtain 100 mL of each of these four solutions, and repeat the foregoing procedure with the KNO 3
solutions as the solvent.
Calculations
−
1. Calculate the [ HC 4H4O6 ] for each of the original, saturated solutions from the known [NaOH] and
the volume of NaOH used in each titration.
2. Calculate the total [K+] for each solution, noting that generally there are two sources of K+ so that
[K + ]tot = [K + ]KNO3 + [K + ]KHC4H4O6
In this calculation, assume that KNO3 acts as if it is 86% ionized.
3. Calculate Ksp values for each solution and their average.
38
Name:
Section
Partner’s Name
Date
Chemical Equilibrium II: The Solubility Product
Data Sheet
[NaOH], M
1.
Temperature, C
Solvent: Distilled H2O
Trial 1
Trial 2
Final burette reading, mL
Initial burette reading, mL
Volume of NaOH used, mL
Moles of NaOH used, mol
Moles of
HC 4H4O6− , mol
Volume of KHC4H4O6, L
Solubility of KHC4H4O6, M
Ksp of KHC4H4O6
39
2.
Solvent: KNO3 (aq)
[KNO3] solvents
0.0100 M
0.0200 M
0.0300 M
0.0400 M
Final burette reading, mL
Initial burette reading, mL
Volume of NaOH used, mL
Moles of NaOH used, mol
Moles of
HC 4H4O6− , mol
Volume of KHC4H4O6, mL
−
[ HC 4H4O6 ], M
[K+]Tot, M
Ksp of KHC4H4O6
Average Ksp for all titrations (including those with distilled water as the solvent) ______________
40
The Rate Law in Chemical Kinetics
Objective
The reaction,
2−
2−
2 I− (aq) + S2O8 (aq) → I2 (aq) + 2 SO4 (aq)
is studied in order to obtain its rates law and constant.
Introduction
The rate law for a generic chemical reaction,
aA + bB → products
(1)
R = k[A][B].
(2)
can be written as
Here,
R
k


 + 
=
=
=
=
=
the reaction rate
the rate constant
the order of the reaction with respect to A
the order of the reaction with respect to B
the overall order of the reaction.
The specific rate law governing a reaction cannot be predicted; it must be determined experimentally.
One technique for doing so is the method of initial rates.
If a series of experiments are performed whereby the initial concentration of reactant B, denoted by [B] 0,

is held constant, then the term k[B] 0 will be approximately constant. Equation (2), written in logarithmic
form, then becomes
LnR0 = Ln[A]0 + constant.
(3)
Thus, by varying the initial concentrations of a, [A] 0, and measuring the corresponding initial rates, R 0, the
order of the reaction with respect to A, , may be determined from the slope of the graph of LnR0 vs.
Ln[A]0.
In a similar fashion, by holding [A]0 constant and varying [B]0 in a series of experiments,  can be
determined.
This technique will be used to study the oxidation of iodide ion by peroxydisulphate ion according to
2−
2−
2 I− (aq) + S2O8 (aq) → I2 (aq) + 2 SO4 (aq).
(4)
41
The presence of I2 produced in reaction (4) is detected by adding starch to the reaction since even in
trace quantities, I2 forms a deep blue coloured complex with starch. Thus if the two reagents in reaction
(4) are mixed in the presence of starch, a deep blue colour would appear almost immediately.
This process occurs too quickly to be of much use in rate studies, therefore another very fast reaction is
allowed to occur in the reaction flask, namely
2−
2−
2 S2O3 (aq) + I2 (aq) → 2 I− (aq) + S4O6 (aq).
(5)
2−
The thiosulphate ion, S2O3 , quickly uses up I2 produced in the main reaction, preventing any
accumulation of I2. Thus free iodine, continually being produced in reaction (4), is used up in reaction (5)
2−
until the thiosulphate ion, S2O3 , is depleted, whereupon iodine produced in reaction (4) combines with
starch to create the characteristic blue coloured solution
The appearance of the blue coloured solution signals the first appearance of unreacted iodine. This
2−
permits an evaluation of the initial reaction rate R 0. The blue colour can only occur when S2O3 is
depleted, thus
𝑅𝑂
−
1 ∆[𝐼 − ]
2 ∆𝑡
= −
O
∆[𝑆2 𝑂82− ]
∆𝑡
=−
O
1 ∆[𝑆2 𝑂32− ]
2
∆𝑡
(6)
Procedure
Record the concentrations of all reagents.
Mix in Flask 1 the volumes of KI, KCl and Na2S2O3 specified in Run 1 of Table 1 below.
Mix in Flask 2 the volumes of (NH4)2S2O8 and (NH4)2SO4 specified in Run 1 of Table 1 below. Add 2-3
drops of starch.
Measure the temperature of each solution. If they differ by more than 1C, warm or cool the solutions as
appropriate. Record the temperature.
Pour the contents of Flask 2 into Flask 1 quickly and start the stopclock simultaneously.
Swirl the solutions gently until the solution turns to a dark colour. Record the time.
Repeat the procedure for Run 1. The time obtained should be within 2 seconds of that of the previous
trial.
Perform the reaction for each of the subsequent runs using the volumes specified for each flask in
Table 1. It is necessary to do these only once.
42
Flask 1
Flask 2
Run
KI, mL
KCl, mL
Na2S2O3, mL
(NH4)2S2O8, mL
(NH4)2SO4, mL
1
30.0
0.0
15.0
30.0
0.0
2
25.0
5.0
15.0
30.0
0.0
3
20.0
10.0
15.0
30.0
0.0
4
15.0
15.0
15.0
30.0
0.0
5
10.0
20.0
15.0
30.0
0.0
6
30.0
0.0
15.0
25.0
5.0
7
30.0
0.0
15.0
20.0
10.0
8
30.0
0.0
15.0
15.0
15.0
9
30.0
0.0
15.0
10.0
20.0
Note: 2-3 drop of starch are added to flask 2
Calculations
2−
1. Calculate the number of moles of S2O3 initially present. This equals twice the number of moles of
2−
S2O2−
8 reacted. Calculate the [ S2O8 ] reacted. Note that the total volume of the solution is always
0.0750 L.
2. Determine the initial reaction rate from equation (6), where t is the time it takes for the reaction
mixture to turn dark.
2−
3. Calculate [I−]0 and [ S2O8 ]0 values for each run.
2−
4. Using the data from Runs 1 − 5, plot LnR0 vs. Ln[I-]0. In these runs Ln[ S2O8 ]0 is constant. From the
slope of the graph, obtain the order of the reaction with respect to I–.
2−
5. Using the data from Run 1 and Runs 6 − 9, plot LnR0 vs. Ln[ S2O8 ]0, Ln[I–]0 being held constant.
2−
The slope of the graph yields the order of the reaction with respect to S2O8 .
Tabulate the data.
6. Calculate the value of the rate constant, k, for each run and its average value. Include the units of k.
43
44
Name:
Section
Partner’s Name
Date
The Rate Law in Chemical Kinetics
Data Sheet
Reagent Concentrations
[KI], M
[Na2S2O3], M
[(NH4)2S2O8], M
T (K)
Run
Time, s
[I– ]0, M
Ln[I– ]0
2−
[ S2O8 ]0, M
2−
Ln[ S2O8 ]0
R0, Ms-1
LnR0
k
1
2
3
4
5
6
7
8
9
Average k:
(including units)
45
46
Activation Energy in Chemical Kinetics
Objective
The rate at which chemical reactions proceed are highly temperature dependent. In this experiment, the
rate of the reaction,
2−
2 I– (aq) + S2 O2−
8 (aq) → I2 (aq) + 2 SO4 (aq)
is measured at various temperatures in order to determine its activation energy. This reaction is also
carried out a catalyst in order to determine its effect on the reaction rate.
Introduction
The temperature dependence of chemical reactions can generally be approximated by the Arrhenius
expression for the rate constant, k,
Lnk = −
Ea 1
+ LnA.
R T
(1)
Here R and T are the universal gas constant and the absolute temperature respectively. The quantity A
is related to the frequency of collisions between the reactant molecules, while E a is the activation energy.
In order for a reaction to occur, colliding reactant molecules must pass through a state whereby their
mutual interactions are strong enough to break bonds and to rearrange themselves before breaking apart
into product molecules. The point at which rearrangement occurs is known as the activated complex.
The activation energy, Ea, is the potential energy of the system at the activated complex relative to the
potential energy of the pre-collisional reactant molecules. Thus Ea is a minimal or threshold energy which
must be supplied to the reactant molecules in order for the reaction to proceed.
The reaction,
2−
2 I− (aq) + S2 O2−
8 (aq) → I2 (aq) + 2 SO4 (aq)
(2)
obeys a second-order bimolecular rate law,
R = k[I − ][S2 O2−
8 ].
(3)
In this experiment, the temperature dependence of the rate constant will be investigated. If the
concentrations of I− and S2 O2−
8 are held constant through a series of rate measurements over various
temperatures, the temperature dependence of k can be determined. Accordingly, as equation (1)
demonstrates, a plot of Lnk vs. T− yields a straight line; from its slope, the activation energy can be
extracted.
47
Procedure
Record the concentrations of KI, Na2S2O3, (NH4)2S2O8 and (NH4)2SO4.
The procedure is the same as for Run 9 of Experiment: The Rate Law in Chemical Kinetics.
Flask 1:
30.0 mL KI
15.0 mL Na2S2O3
Flask 2:
10.0 mL(NH4)2S2O8
20.0 mL (NH4)2SO4
2–3 drops starch
Immerse both flasks in the temperature baths until both are at the bath temperature.
temperature of the solutions.
Record the
Mix the solutions with a glass rod (NEVER use a thermometer to stir a solution) and start the stopclock
simultaneously. Keep the flask in the bath.
Record the time it takes the solution to turn a dark colour.
Repeat this procedure for all the other temperature baths, including a run at room temperature.
Repeat the procedure for a run at room temperature except that 1–2 drops of a catalyst must be added to
Flask 1 before mixing the contents of the two flasks (starch is also added in flask 2). Record the time and
temperature.
Calculations
1.
Calculate the initial reaction rates for each temperature study as in equation (6) of the preceding
experiment, The Rate Law in Chemical Kinetics.
2.
Calculate the rate constant, k, at each temperature.
3.
Plot Lnk vs. T−1. From the slope of the line, determine Ea.
4.
Calculate the ratio of the reaction rates for the catalyzed to the uncatalyzed reactions at room
temperature.
48
Name:
Section
Partner’s Name
Date
Activation Energy in Chemical Kinetics
Data Sheet
Reagent Concentrations
[KI], M
Na2S2O3, M
(NH4)2S2O8, M
Flask 1
Flask 2
Run
VKI, mL
VNa 2 S 2 O 3 , mL
V(NH4 )2 S2O8 , mL
V(NH4 )2 SO4 , mL
Starch
All
30.0
15.0
10.0
20.0
3-5 drops
Condition
Run
Temperature, C
Time, s
1
2
Uncatalyzed
3
4
5
Catalyzed
6
49
Calculation
[I−]0, M
[S2 O2−
8 ]0, M
Initial moles S2 O2−
3 , mol
[S2 O2−
8 ], M
Condition
Run
1
2
Uncatalyzed
3
4
5
Catalyzed
Ea, kJmol–1
R cat
Runcat
50
6
T, K
T−1, K-1
Rate, Ms-1
k, Lmol-1 s-1
Lnk
Indirect Determination of a Thermodynamic Quantity
Objective
The solubility of potassium nitrate in water at various temperatures is used to determine its standard
enthalpy of solution
Introduction
Changes in thermodynamic quantities cannot be measured directly for many physicochemical processes.
If, however, a process does not go to completion, but instead attains a state of equilibrium, measurement
of the equilibrium constant, K, may permit the evaluation of several thermodynamic functions.
At equilibrium
G = −RT LnK
(1)
or
LnK = −
∆G°
RT
(2)
.
Using the definition of the free energy,
G = H − TS.
(3)
Substitution of equation (3) into equation (2) yields the Van’t Hoff equation,
LnK =
−
∆𝐻°
𝑅
∙
1
𝑇
+
∆𝑆°
(4)
𝑅
If it is assumed that H and S are not very strongly dependent on temperature, the terms H/R and
S/R will be approximately constant over moderate temperature ranges. With these assumptions, the
Van’t Hoff equation permits the evaluation of H provided that the equilibrium constants can be
measured at various temperatures.
In this experiment, the solubility of potassium nitrate, KNO 3, in water will be used to determine H so ln from
the Van’t Hoff equation. Potassium nitrate dissociates according to the equation,
H2O
KNO3 (s)
K+ (aq) + NO−
3 (aq)
(5)
This dissociation equilibrium is examined at several temperatures. The equilibrium constant for equation
(5) is known as the solubility product, Ksp,
K sp = [K + ][NO−
3]
Since there are no sources of K+ or
(6)
NO3− apart from KNO3 in this experiment,
[K + ] = [NO−
3] ≡ S
(7)
51
from the stoichiometry of reaction (5). Furthermore, inasmuch as equilibrium has been established, the
solution is saturated so that S is the solubility of the salt. With the result of equation (7), the solubility
product may be rewritten
K sp = S 2
(8)
Ln K = 2 LnS
(9)
or
With this consideration, van’t Hoff’s equation (4) takes the form
LnS = −
∆H 1
∙ + constant
2R T
(10)
Here, H is the enthalpy of solution of KNO3 associated with the chemical reaction (5).
LnS vs. T − yields a straight line with a slope of −H / 2R, by virtue of equation (10).
A plot of
Procedure
Put a 400 mL beaker filled with about 300 mL of tap water to heat on a hot plate.
Weigh about 8 g KNO3 to three decimal places. Transfer the KNO3 to a clean, dry boiling tube.
Wash and fill a burette with distilled water.
Add 5.00 mL of the distilled water to the KNO3 in the boiling tube and warm the mixture in the hot bath
and stir gently until all the KNO3 has dissolved.
Calculate the concentration of this solution in units of g KNO 3/100 g H2O. This is essential as the
experiment cannot be conducted without this value (see the Calculations section).
Make sure the temperature sensor is connected to the CH1 input of the LabQuest mini interface and that
the LabQuest interface is on (green LED on).
Log onto the computer network.
Double click on the Logger Pro icon to start the application. Open the Chemistry of Solutions folder and
double click on the file 09_Indirect thermo. . If the sensor confirmation box opens, select the DirectConnect Temp sensor and press connect. The file is set to collect data with entry. The entries are the
calculated KNO3 solution concentrations. Click on the Collect button.
Allow the solution to cool while stirring with the temperature sensor. When the first crystals (needles of
KNO3) begin to form, click on Keep. Enter the concentration into the entry box that appears.
To the boiling tube, successively add 1.00 mL, 2.00 mL, 2.00 mL, 5.00 mL and 5.00 mL of distilled water
from the burette. Remember to calculate the concentration of the KNO3 upon each successive addition of
water. Following each addition, warm the KNO3 to dissolve it and on cooling, record the concentration
and temperature at which the first crystals appear.
After the last addition, click on Stop.
Wash the temperature sensor to remove any KNO3 crystals attached to it.
52
Set up two calculated columns, one to calculate LnS and the other to calculate 1/T. Be sure to use the
absolute temperature in the latter calculation.
Procedure to create the two columns
To create the column: Ln S
Menu: Data → New calculated column
Name: Ln S
Equation: Function> → ln()
Variables (column)> → solubility
Done
To create the column: 1/T
Menu: Data → New calculated column
Name: 1/T
Equation: 1/(273.15+Variables (column)> → temperature)
Done
Use Logger Pro to graph LnS vs. 1/T and to obtain the regression line for this data. Provide a suitable title
for this graph.
Print the graph window and save the file to a Flash drive.
Log off the network
Calculations
1. Assuming that the density of water is 1.0 gmL−, calculate the solubility of KNO3 according to
S=
mass of KNO3
mass of H2 O
∙ 100
2. Plot LnS vs. 1/T. This is known as the van’t Hoff plot.
3. Obtain H from the slope of the van’t Hoff plot.
4. Report the molar enthalpy of solution for KNO3.
53
54
Name:
Section:
Partner’s Name:
Date:
INDIRECT DETERMINATION OF A THERMODYNAMIC QUANTITY
DATA SHEET
Mass of KNO3, g
Total volume of water added, mL
Run 1
Run 2
Run 3
Run 4
Run 5
Run 6
5.00
6.00
8.00
10.00
15.00
20.00
Temperature of saturated solution, K
1 / T, K−
Solubility of KNO3, g KNO3/100 g H2O
LnS
Slope of van’t Hoff Plot
Molar enthalpy of solution, kJmol−
55
56
COMMON LABORATORY EQUIPMENT
Beaker
Test Tube Holder
Test Tube
Erlenmeyer Flask
Graduated Cylinder
10 mL Pipette
Medicine Dropper
Glass Stirring Rod
Evaporating Dish
Crucible & Crucible dish
Glass Funnel
Watch Glass
Plastic Weighing Dish
Scoopula
Wash Bottle
Thermometer
57
58
IA
IIA
IIIB
IVB
VB
VIB
VIIB
VIII
IB
IIB
IIIA
IVA
VA
VIA
VIIA
0
1
2
H
He
1.01
4.00
3
4
5
6
7
8
9
10
Li
Be
B
C
N
O
F
Ne
6.94
9.01
10.81
12.01
14.01
16.00
19.00
20.18
Transition Metals
11
12
13
14
15
16
17
18
Na
Mg
Al
Si
P
S
Cl
Ar
22.99
24.31
26.98
28.09
30.97
32.07
35.45
39.95
19
20
21
22
23
24
25
26
27
28
29
30
31
32
33
34
35
36
K
Ca
Sc
Ti
V
Cr
Mn
Fe
Co
Ni
Cu
Zn
Ga
Ge
As
Se
Br
Kr
39.10
40.08
44.96
47.87
50.94
52.00
54.94
55.85
58.93
58.69
63.55
65.39
69.72
72.61
74.92
78.96
79.90
83.80
37
38
39
40
41
42
43
44
45
46
47
48
49
50
51
52
53
54
Rb
Sr
Y
Zr
Nb
Mo
Tc
Ru
Rh
Pd
Ag
Cd
In
Sn
Sb
Te
I
Xe
85.47
87.62
88.91
91.22
92.91
95.94
98
101.07
102.91
106.42
107.87
112.41
114.82
118.71
121.76
127.60
126.90
131.29
55
56
57
72
73
74
75
76
77
78
79
80
81
82
83
84
85
86
Cs
Ba
La*
Hf
Ta
W
Re
Os
Ir
Pt
Au
Hg
Tl
Pb
Bi
Po
At
Rn
132.91
137.33
138.91
178.49
180.95
183.84
186.21
190.21
192.22
195.08
196.97
200.59
204.38
207.2
208.98
209
210
222
87
88
89
104
105
106
107
108
109
110
111
112
Fr
Ra
Ac**
Rf
Unp
Unh
Uns
Uno
Une
Uun
Uuu
Uub
223
226.03
227.03
261
262
266
264
269
268
269
272
277
*Lanthanides
**Actinides
58
59
60
61
62
63
64
65
66
67
68
69
70
71
Ce
Pr
Nd
Pm
Sm
Eu
Gd
Tb
Dy
Ho
Er
Tm
Yb
Lu
140.12
140.91
144.24
145
150.36
151.96
157.25
158.93
162.50
164.93
167.26
168.93
173.04
174.97
90
91
92
93
94
95
96
97
98
99
100
101
102
103
Th
Pa
U
Np
Pu
Am
Cm
Bk
Cf
Es
Fm
Md
No
Lr
232.038
231.036
238.029
237.048
244
243
247
247
251
252
257
258
259
262