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Inorganic Chemistry II
Laboratory
William Carey University
Summer, 2012
2
Table of Contents
Introduction to the Chemical Laboratory: Safety……………………. 3
Calorimetry……………………………………………………………... 12
Determine the Gas Constant, R ……………………………………….. 24
Physical Properties- Melting Point ………………………….……….. ..29
Aspirin Synthesis & Analysis…………………………………….…….. 37
Chemical Kinetics ……………………………………………………… 45
Equilibrium ……………………………………………………………... 59
pH Measurements ……………………………………………………… 70
Solubility Product Constant and Common Ion Effect (Ksp)…………. 79
Acid/Base Titration……………………………………………………… 89
The Determination of a Molar Mass…………………………………… 93
Visible Spectrophotometry…………………………………………… 102
3
An Introduction to the General Chemistry Laboratory:
Safety
Objectives
--Learn basic safety rules and procedures
--Check into the laboratory
Laboratory Safety
Experience has demonstrated the necessity of maintaining a constant awareness of the
hazards of experimental work performed in the study of chemistry. The best way for a
student to protect him/herself and co workers is to incorporate safety as an integral part of
each task.
You are expected to exercise good judgment and common sense in preventing any
hazardous situations. This safety handbook has been compiled to help you plan your
work in a safe manner. Remember that these safe practices are not hard and fast and that
all circumstances are not covered. If under exceptional circumstances good judgment
indicates certain of these practices should not be followed, then make certain you are not
creating a hazard by selecting a different procedure.
4
I. Laboratory Safety Rules
A. Eye Protection
All students must wear safely goggles or safety glasses while doing work in the
laboratory. When in the lab, your first act should be to put your safety glasses on and
your last act prior to leaving the lab should be to take them off. If you need a break from
wearing them, then go into the hall and take them off.
B. Horseplay
Horseplay and practical joking of any kind are strictly forbidden. People who engage in it
in the lab will be ejected and will receive a grade of F (zero).
C. Working Alone
No one is allowed to perform experimental work in a lab unless a second person is
present or nearby. The instructor must be present in general chemistry.
D. Work Authorization
Unauthorized experiments are forbidden. Extra work or original work is encouraged, but
before any experiment is performed in the lab, approval must be given by the Instructor
in charge.
E. Safety Precautions
Before performing any experiment, become familiar with any safety hazards which may
be present such as flammability or toxicity, and take necessary precautions.
F. Reporting Accidents and Fires
All accidents resulting in Injury, property damage, or fire must be reported immediately
to an Instructor.
G. Eating
Preparation, storage, or consumption of food or drink in work areas is strictly forbidden
because of the danger of contamination with toxic/poisonous substances; Hands should
be thoroughly washed before leaving the lab and before handling food.
H. Smoking
Smoking is not permitted in the labs, stockrooms, or classrooms. Flammable liquids,
vapors, and gases create a definite fire hazard.
I. Appearance
No sandals or open-toed shoes are allowed in lab because of frequent breakages and
spills. Long hair must be pulled back. Students must check out with the lab Instructor to
ensure that the lab bench is in the appropriate order.
5
II Emergency Procedures
A. In case of accident or Illness
-Render prompt first aid,
-Call the instructor immediately for assistance
-If injury or illness appears to be serious, have someone call an ambulance
-Report all accidents that cause injury, no matter how minor, immediately, to your
Instructor.
B. In case of fire
-Immediately notify the instructor
-If the fire is small and easily extinguished, use the lab extinguisher at once.
-If the fire is large and difficult to extinguish, pull the alarm switch and call the fire
department. Immediately evacuate the area and building.
-If your clothing should catch fire, STAY CALM, and get under the nearest shower
while calling for help.
-If the fire alarm sounds, immediately shut down your experiment and evacuate the
building.
-Tag any used fire extinguisher as empty and take it to your instructor for replacement.
C. First Aid
If any chemical gets into your eyes or mouth or on your skin, go quickly to the nearest
eye wash sink or shower and wash with as much water as possible. If the eye is involved,
hold the eyelids open with your fingers and allow water to run freely over the eyeball for
at least 10 minutes. This is essential to prevent permanent eye damage. Call your
Instructor immediately. It is best to assume that all substances except pure water are
harmful when in contact with any part of the body.
If you are cut, apply direct pressure on the cut to stop the bleeding. All cuts, bums, and
other injuries should be reported to your instructor at once.
6
III. GENERAL LABORATORY PRACTICES
A. Safety Check
- Locate all exits from the lab and building.
- Locate the fire extinguishers, safety showers, and eye wash fountains.
- Locate the nearest telephone for use in art emergency.
- Locate the power line cut-offs and the nearest fire alarm switch.
- For laboratory safety it is essential that the student include in the experimental design a
provision for power or water failure which could cause an accident situation or an unsafe
condition to develop. For example, consider the possibility of the loss of cooling water to
a condenser of the loss of power to an expensive instrument. Plan to meet such an
eventuality.
- Ongoing experiments/operations are not to be left unattended. Unsafe conditions or
practices cause virtually all accidents. Anyone observing unsafe practices or situations
should notify the instructor.
B. Housekeeping (Removal of the Hazard)
The continuous practice of good housekeeping is essential to the prevention of accidents,
fires, and injuries. Students are expected to keep their benches neat and orderly; a
cluttered laboratory is a dangerous place in which to work.
- Keep benches, balances, tables, hoods, and floors clear of all materials not in use.
- Keep adequate passageways to exits.
- Carefully clean up spills and broken glass. Broken glass should go in the garbage cans.
Use proper waste disposal techniques for spilled materials.
- Keep all chemical containers clean, properly labeled, and capped.
- DO NOT use eye wash fountains for drinking or for disposal of chemicals.
- DO NOT borrow equipment.
C Handling Chemicals
Chemicals can be dangerous unless properly handled. Before working with any chemical,
know its properties. Hazardous chemicals include those which are flammable, toxic,
corrosive, and/or reactive, Use of a fume hood is mandatory when undesirable vapors are
produced. The Merck Index is a good source of information on many common chemicals
and their properties.
A very important safety practice is to keep all material properly labeled. The label should
show the following information: Chemical name and structure; Brief notation of any
hazard; Date of purchase, preparation, or transfer to present container, and for solutions,
the concentration and name of preparer.
7
Safety Precautions
- Confine hair, No sandals (open toed shoes) permitted.
- Use great care l transporting chemicals and keep reagent containers clean
- Avoid contact of chemicals with skin; use plastic/rubber gloves as necessary.
- Never taste a chemical or pipette by mouth, and avoid breathing vapor
- Avoid contamination by never returning excess portions of reagent back into the stock
bottles.
- Always add concentrated chemicals to water, never the reverse.
- Keep flammable solvents away from heat and flame.
- Use caution when working with mercury. The equilibrium concentration of the vapor
over liquid mercury at room temperature is 20 times the threshold toxic limit. (Hatters
went mad from mercury chloride poisoning.)
- Clean up spills immediately. Use sodium hydrogen carbonate (bicarbonate) to
neutralize acids, and dilute acetic acid to neutralize bases.
- Never look directly into a test tube or beaker, or point them toward anyone. The
formation of a gas could cause the contents to come flying out, especially when heating.
- When using a reagent from a bottle always read the label twice to be certain you
have the correct reagent. When pouring solutions from bottles, hold the label towards
your hand so any excess solution does not drip on the label, if drips occur, wipe the
outside of the bottle.
- Always place the cap back on reagent bottles. Leave the reagent bottles on the reagent
table where you found them.
D. Chemical Waste Disposal
Most water soluble wastes may be poured down the sink with cold tap water. Specific
disposal instructions will be given for liquid wastes not miscible with water. Special
containers will be provided for water insoluble solids.
E. Handling Compressed Gas.
Know the cylinder contents and its properties. Always transport cylinders carefully, using
a wheeled cart. Keep the cylinder strapped down in the cart or at the bench. To remove
gas through a regulator, make sure all valves are closed, and then slowly open all valves,
starting with the cylinder master valve.
**General chemistry students will not handle compressed gas without the instructor’s
supervision.
8
F Handling Laboratory Glassware
- Always carry glass tubing in a vertical position. Protect your hands with a towel when
cutting tubing
- Use glycerin to lubricate the surface of glass tubing and thermometers before inserting
them into rubber stoppers, and protect your hands with a towel. Never push on glass! If a
stopper is difficult to remove, cut it off the glass rather than risk breaking the glass
or ramming the glass through your hand - it can happen!
- Broken or cracked glassware is very dangerous and should be discarded.
- Never heat graduated cylinders, volumetric flasks, bottles, or funnels over a flame. A
watch glass made of Pyrex or kimax brand glass may be heated with a cool blue flame.
G. Fire Prevention
Three components are necessary for a fire to start: fuel, oxidizing agent, and an ignition
source (usually heat). Most fires can be prevented by keeping fuel and oxidant away from
hot ignition sources.
Fire can usually be extinguished using the same principles that tend to avoid it.
- Reduce the air supply by smothering - cover the vessel or apply C02.
- Shut off or reduce the fuel supply.
- Cool the fuel below its ignition temperature.
- Lower the concentration of the fuel by diluting it with an inert material.
The Four Classes of Fires and How to Extinguish Them
Class A: burning wood, paper, cloth, etc.; use water, foam, soda-acid, or C02.
Class B: burning oil, greases, paints, etc.: use foam, CO2 or dry chemical.
Class C: live electrical equipment: extinguish with CO2 or dry chemical.
Class D: active metals, such as sodium, potassium, magnesium, etc.; extinguish by
smothering with dry soda ash or dry chemical extinguisher.
9
H. The Laboratory Instructor’s Role in Matters of Safety
- To teach and enforce safety as an integral part of the course.
- To see that all safety rules are obeyed, and to set a proper example.
- To remain in the laboratory at all limes when students are present.
Laboratory Safety: I have carefully read and studied the Safety Handbook. I will follow
the rules for laboratory safety. I am prepared to take a quiz on the rules; and I understand
that I must have a passing score in order to continue in the lab. The safety test will be
given at the beginning of the next laboratory.
Signature __________________ Date _________________ Sec. # ______
10
Chemistry 111-112
Desk Drawer Equipment List
Desk No.: _________ Section No.: _________ Name:___________________
Quantity Description Check-in Check-out
3 Beakers, l00mL _______ _______
2 Beakers, 250mL _______ _______
1 Beaker, 400mL _______ _______
1 Crucible & Cover _______ _______
1 Evaporating Dish, small _______ ______
3 Erlenmeyer Flasks, 250mL _______ _______
1 Florence Flask, 500mL _______ _______
1 Forceps _______ _______
1 Funnel _______ _______
1 Graduated Cylinder, 50mL _______ _______
1 Graduated Cylinder, l0mL _______ _______
5 Medicine Droppers _______ _______
2 Stirring Rods _______ _______
10 Test Tubes, 13 x 100mm _______ _______
3 Test Tubes, 18 x 150mm _______ _______
1 Test Tube Brush _______ _______
1 Test Tube Holder _______ _______
1 Test Tube Block _______ _______
1 Triangle ________ _______
1 Watch Glass ________ _______
Unit Drawer
1 Bunsen Burner ________ _______
1 Pair Crucible Tongs ________ _______
1 Pneumatic Trough ________ _______
1 Pinch Clamp ________ _______
1 Ring with Ring Stand ________ _______
1 Set Test Tube Clamps ________ ________
11
1 Test Tube Rack _________ _________
1 Wire Gauze _________ _________
12
13
Calorimetry
Objectives
--To determine a calorimetric constant
--To use a simple solution calorimeter to obtain heats of reaction
--To use Hess’s Law of Heat Summation to determine an unknown ΔH value
Introduction
All chemical and physical changes in matter are accompanied by a transfer of heat
(energy): heat may be evolved (exothermic) or absorbed (endothermic). A device called a
calorimeter will be used to measure the quantity and direction of heat flow occurring
when matter undergoes a change. The heat flow for chemical reactions is quantitatively
expressed as the heat (or enthalpy) of reaction, ΔH, at constant temperature and pressure.
A negative ΔH value would indicate an exothermic reaction and a positive ΔH value
indicates an endothermic reaction. The heat of reaction can be expressed per gram or per
mole of substance; it is then called the specific heat of reaction or the molar heat of
reaction, respectively.
Procedure
Part A - Preparation and Calibration of the Calorimeter
Prepare a calorimeter by nesting one Styrofoam cup inside another. Use a cardboard
square as a lid. Insert the thermometer through the hole in the lid; the solution should be
stirred gently with a thermometer. The thermometer should not be removed from the
solution to take a reading.
Calorimeter Calibration: The purpose of this exercise is to determine the amount of
energy absorbed by your calorimeter. (See Chart in Special Tables Section)
1. Allow the hot water at your bench to run for several minutes. The temperature will
reach approximately 40 oC. Collect 50m1 of the hot tap water in a graduated cylinder.
2. Measure the mass of a clean, dry calorimeter to 0.01g.
3. Add 50ml of room temperature water to the calorimeter and determine the mass of the
calorimeter and water.
4. Allow the heated water to cool to between 30 and 35 degrees.
5. Measure and record (on the data sheet) the temperature of each of the two vessels of
water.
6. Immediately pour the hot water into the calorimeter and begin stirring carefully.
7. Record the temperature reading (page 10 special tables) every 30 seconds until the
temperature remains constant--while stirring. (About 5 minutes)
8. Weigh the calorimeter and the contents to determine the amount of hot water added.
14
9. Steps for Calculating Calorimeter constant:
Step 1 - Calculate the energy lost by hot water (q).
q = (mass hot water) (specific heat of water) (ΔT)
Specific Heat (of Solution) = 4.18 J/g °C
ΔT = Tf -Ti
Step 2 - Calculate the energy gained by cool water (q).
q = (mass cool water) (specific heat) (Δ T)
Step 3 - Calculate energy absorbed by the calorimeter (q).
q = heat lost by warm water - heat gained by cool water
q = (absolute value Step 1) - (absolute value Step 2)
Step 4 - Calculate the Calorimeter Constant (Q).
Q = energy absorbed by calorimeter temperature change of the cool water
Important Note
1. The calorimeter constant is specific for your calorimeter. You must use your Q value
in future calculations.
2. If you get a negative number for your Q value, use zero as your calorimeter
constant.
15
Part B - Change in Heat Determination
Caution: HCI, NH3 and NaOH can cause chemical burns. If a spill occurs, wash the
contaminated area and report the incident to your laboratory instructor.
1. Obtain exactly 50ml of the 2.0 M solution of HCl in a clean, dry graduated cylinder.
2. Obtain exactly 50m1 of the 2.0 M solution of NaOH.
3. Measure the temperature of each of these solutions. Rinse and dry the thermometer
before making each determination.
4. If the temperatures of the two solutions are different immerse the warmer solution into
cool tap water until the solutions are within 2 degrees of each other.
5. Record the average temperature of the solutions before mixing.
6. Add all of the acid to the calorimeter and then add the entire base to the calorimeter.
(Make sure you keep the graduated cylinders with acid and base separate and don’t cross
contaminate.)
7. Immediately place the top on the calorimeter and stir gently with the thermometer
taking care not to scrape the sides and bottom of the calorimeter.
8. Record the temperature to the nearest 0.1 °C after 30 seconds and every 30 seconds
thereafter for 4 minutes.
9. Plot the temperature vs. time data to determine the final temperature. This extrapolated
high temperature is used to calculate z (See Graph)
Note:
(a) Plot the Average Initial temperature at time = 0.
(b) Draw a straight line of best fit through the highest temperature readings until you
reach time = 0.
(c) Extrapolate back to the y-axis to determine T final
16
10. Calculate heat of solution (q)
(Thermal energy change of Calorimeter) + (Thermal energy change of solution)
(Your Calorimeter Constant) (ΔT) + (Total Mass Solution) (Sp. Heat) (ΔT)
Note: (a) Specific Heat of Water = 4.180 J/g °C
Density of water is 1.0 g/ml (Calculate total mass from total volume)
(c) Temperature Change (ΔT = Tf – Ti)
Tf = from graph Ti = average before mix
Useful Information:
1. If Tf > Ti, the energy change is exothermic.
2. If Tf < Ti, the energy change is endothermic.
3. If your results indicate an exothermic process, you must record q as a negative number.
If endothermic, q is recorded as a positive number.
11. Calculate the Molar Enthalpy of Reaction, (ΔH)
AH = q/moles of base used (moles = molarity x volume)
(Include the appropriate sign)
12. Repeat steps 1 - 11 for NH3 and HC1 combination.
17
Part C - Use Hess’s Law to predict ΔH value for
NaOH + NH4Cl—--> NaC1 + NH3 + H2O
You may have noticed that there is a relationship among the equations that we are dealing
with in this experiment.
NaOH + HC1 ---> NaCl + H2O ΔH (known from Part B)
NH3 + HC1 ---> NH4Cl ΔH (known from Part B)
Target Equation: NaOH + NH4Cl—--> NaC1 + NH3 + H2O ΔH unknown
If you add the first two equations in a certain manner (forward or reverse reaction) the
target equation is obtained.
Hess’s Law of Heat Summation - If a reaction is carried out in a series of steps, the ΔH
for the reaction will be equal to the sum of the enthalpy changes for the individual steps.
Important Points to Remember:
(a) If a reaction is reversed, the sign of the ΔH must be changed (negative to positive or
positive to negative).
(b) If a reaction is multiplied or divided by a number, the ΔH is multiplied or divided by
the same number.
(c) If a substance occurs on both sides of the yield mark, it can be cancelled.
18
Complete the following:
1. Use Hess’s law and your experimentally determined ΔH values for NaOH-HC1
and NH3 - HCl reactions to calculate the ΔH value for this target reaction.
(Show your work.)
NaOH + NH4Cl—--> NaC1 + NH3 + H2O ?ΔH
2. Calculate the % deviation:
Using the calculated ΔH value, for your target equation above, as your experimental
value (EV) and -3.9 kJ/mole for the standard value (SV), calculate the % deviation.
% deviation = SV-EV/SV
19
Name: _______________________________ Lab Section: _________________
Lab Partner:___________________________ Lab Desk: ___________________
W. Data
A. Calorimeter Calibration
1. Mass of dry calorimeter ___________________ g
2. Mass of Calorimeter and Water ___________________ g
3. Mass of Water ( #2 - #1) ___________________g
4. Temperature of Water before Mixing ___________________
5. Temperature of Hot Water before Mixing ___________________
6. Final Equilibrium Temp. After Mixing ___________________
7. Mass of Calorimeter and Contents ___________________g
8. Mass of Hot Water (#7- #2) ___________________g
9. Calorimeter Constant ___________________J / °C
Show all work for #9 below.
Remember: If your calorimeter constant is negative, record zero.
Step 1: ________________________ Step 3: _______________________
Step 2: ________________________ Step 4: _____________________
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B. Enthalpy Change (NaOH – HCl)
1. Concentration of HCI _____________________
2. Concentration of NaOH _____________________
3. Concentration of NH3 _______________________________
NaOH - HCI NH3 - HCl
4. Volume Acid added : __________ _________
5. Volume Base Added __________ _________
6. Average Initial Temperature (Before Mixing) __________ _________
7. Final Extrapolated Temp. (From Graph) __________ ________
8. Moles of Base Used __________ ________
9. ΔH (kj/mole) __________ ________
Show your work for 7 - 9 below.
21
C. Hess’s Law and ΔH Determination:
Reaction 1 ____________________________ ΔH =
Reaction 2 ____________________________ ΔH =
Total ____________________________ ΔH =
Net Equation ____________________________ ΔH =
22
23
PRE-LABORATORY ASSIGNMENT:
Name___________________________ 1. A metal sample weighing 45.2 g and at a temperature of 100.0°C was placed in 38.6 g of
water in a calorimeter at 25.2°C. At equilibrium the temperature of the water and metal
was 33.0°C.
a. What was ∆t for the water? (∆t = tfinal – tinitial)
__________°C
b. What was ∆t for the metal?
__________°C
c. How much heat flowed into the water?
__________ joules
d. Taking the specific heat of water to be 4.18 J/g°C, calculate the specific heat of
the metal. (q = –mmcmΔTm = mwcwΔTw)
__________ joules/g°C
e. What is the approximate molar mass of the metal? ( )
__________ g
2. When 2.0 g of NaOH were dissolved in 49.0 g water in a calorimeter at 24.0°C, the
temperature of the solution went up to 34.5°C.
a. Is this solution reaction exothermic? _________ Why?
b. Calculate qwater. (q = mcΔT)
_________ joules
c. Find ∆H for the reaction as it occurred in the calorimeter (∆H = qrxn = -qw)
∆H = _________ joules
d. Find ∆H for the solution of 1.00 g NaOH in water.
∆H = _________ joules/g
e. Find ∆H for the solution of 1 mole NaOH in water.
∆H = _________ joules/mol
24
f. Given that NaOH exists as a Na+ and OH- ions in solution, write the equation for
the reaction that occurs when NaOH is dissolved in water.
g. Given the following heats of formation, ∆Hf, in kJ per mole, as obtained from a
table of ∆Hf data, calculate ∆H for the reaction in Part f. Compare your answer
with the result you obtained in Part e.
NaOH(s), -425.6; Na+, -240.1; OH-, -230.0
∆H = _________kJ
25
Determine the Gas Constant, R
Objective
--To calculate R by collecting a gas at known volume, pressure, and temperature.
--To further a general understanding of the relationship of pressure volume and
temperature.
Introduction
In this experiment, magnesium metal will be reacted with HCI producing hydrogen gas.
From the balanced equation, the amount of hydrogen produced cart is calculated (n). The
pressure inside the reaction will be maintained at atmospheric pressure (P). The volume is
determined by allowing the gas formed in the reaction to displace the equivalent amount
of water which can be measured with some degree of accuracy (V). Finally, the
temperature is that of the surroundings which can be recorded as room temperature (T).
With this date one can determine the constant(R) in the equation 1W = nRT.
Procedures
1. Assemble the apparatus shown below.
2. Place .5 to .85g of Magnesium measured to three decimal places in a small vial.
3. The middle l000mL Erlenmeyer flask and tubing should be filled with tap water and
the tube clamped so that the water inside is trapped in the tube leaving no air spaces.
Then place the end of the filled tube into the clean, dry l000mL flask.
4. Use a ring stand and clamp to raise the top of the 25OmL Erlenmeyer flask to the level
to the l000mL flask. Use forceps to place the vial of Mg into the 25OmL flask without
spilling the Mg. (As shown in the diagram)
5. Carefully, use a long stem funnel to slowly add 4OmL 3.OM HC1 to the bottom of the
25OmL Erlenmeyer flask. Do not allow the Mg to contact the acid.
CAUTION: HCL is a strong acid. It will cause severe burns and is a danger to the
eyes. Wear eye protection.
26
6. Carefully, replace the stoppers and make sure they are secure. As the reaction proceeds
you and your partner must hold the stoppers in place to assure that the pressure does not
blow the top off. If gas leaks out, you must start over.
7. Remove the Clamp on rubber tubing Tip the vial over and spill the Mg into the HC1
will securing the stoppers.
Caution: Extremely exothermic reaction - touch only the neck of the 25OmL
Erlenmeyer flask.
8. In the reaction will proceed for a few minutes. When it has stopped allow the system to
equilibrate for approximately half an hour.
9. To equalize pressure, make H2O levels in the 2 large Erlenmeyer flask equal - raise the
flask with less H2O volume up. Hold in place approximately 10 seconds, clamp off
tubing.
10. Determine as accurately as possible using a 200mL graduated cylinder the volume of
water in the collecting vessel.
27
Name: _______________________________ Lab Section: _________________
Lab Partner:___________________________ Lab Desk: ___________________
Data and Calculations
Data
a. mass of Mg g
b. volume water displaced mL
c. Room temperature (System Temp.) °C.
d. Vapor Pressure of system in mmHg mmHg
(from chart)
e. Barometric Pressure ____________mmHg
Show the proper calculation for each step with respective units.
a. Volume of water displaced. (L)
b. Volume of gas produced. (Vol. gas = vol. water displaced)
c. Pressure of hydrogen gas = (Barometric pressure - water vapor pressure)
d. Pressure of hydrogen gas in bar.
1 atm = 1.01325 bar
28
e. Absolute temperature of system (K).
f. Balanced reaction equation of Mg and HC1.
g. Moles of Mg used.
h. Moles of hydrogen gas produced. (from g above)
I. Calculate R using the ideal gas equation and your recorded data to 4 significant digits
for the variables P, V, n, and T.
PV = nRT (units = L bar/mole K)
J. Calculate the percent deviation from the standard value of R.
R = 0.083 1442 L bar/mol K
% = [std. value - experimental value / std. value] X 100
29
Pre-laboratory Assignment
Use the gas constant, 0.0831442 L bar / mol K, and the following conversion factors:
76 cm Hg/atm, 1.0 1325 bar / atm, 1000 mL/L, 1 atm/14.7 psi, 1 Torr/ 1 mm Hg, 1
cm3/mL; to calculate answers to the following problems.
1. What is the volume per mole of gas at 25.00°C and 1.00 atm?
2. How many moles are in 64.21 cm³ 02 at 22.3°C and 0.953 1 bar?
3. What pressure will 192 mol of C02 exert at 25°C when it occupies 2.50L?
4. What is the temperature of 0.01937 mol H2 at 121.7 Ton when it occupies 4.137L?
5. What is the volume of 1.000 mol Ar at -19 1°C and 23.94 cm Hg?
30
Physical Properties- Melting Point
Objectives
--Experimentally determine the melting point of pure substances
--Experimentally determine the effect of purity on melting point
Introduction:
How Are Melting Points Used?
The melting point (mp) of a pure compound is defined as that temperature at which the
solid and liquid phase of the compound are in equilibrium at a certain pressure, usually 1
atmosphere. Most crystalline organic compounds have characteristic melting points that
are sufficiently low (30-300° C) to be conveniently determined with simple equipment.
Organic chemists routinely use melting points (1) to get an indication of the purity of
crystalline compounds and (2) to help in identification of such compounds.
Pure crystalline compounds usually have a sharp melting point. That is, the melting-point
range (the difference between the temperature at which the sample begins to melt and the
temperature at which the sample is completely melted) is small (narrow), usually within
1°. Impurities, even when present in small amounts, usually depress the melting point
and broaden the melting-point range. A wide melting point range (more than 5° C)
usually indicates that the substance is impure; a narrow melting point range (0.5-2° C)
usually indicates that the substance is fairly pure. However, there are some exceptions to
both of these generalizations. If two samples have identical structures, their mixture
melting point is not depressed and the melting-point range is not broadened. Small
differences in melting point (on the order of 2-3° C) may also result from variations in
technique, in thermometer accuracy, and in the amount of experience possessed by the
person doing the melting-point determination.
General Technique for Melting-Point Determination:
To determine the melting point of a crystalline substance, a small amount of the finely
powdered material is placed into a thin-walled capillary tube that is sealed at one end.
The capillary tube is inserted into a melting point apparatus in which the temperature can
be measured when heated. The first trial with the sample is to determine an approximate
melting point of the sample. This will be done with heating at a rate of 10° per minute.
To obtain an accurate melting point (A new sample will be set up just like the first), it is
necessary to heat SLOWLY, proceeding through the melting point at no faster than 1° per
minute. Two temperatures are recorded: the temperature at which the substance begins to
liquefy and the temperature at which it becomes completely liquefied. The observed
melting-point range is the interval between these two temperatures.
The observed melting-point range can be influenced not only by the purity of the material
but also by the size of the crystals, the amount of material, the density of its packing in
the tube, and the rate of heating. A finite time is required to transfer heat from a hot liquid
through the walls of the capillary tube and throughout the mass of the sample. When the
31
heating fluid is heated too quickly, its temperature rises several degrees during the time
required for melting to occur. This can result in an observed range that is higher than the
true one.
When the temperature of the heating fluid approaches the melting point of the sample, it
is essential for good results to raise the temperature slowly and at a uniform rate, about 1-
2° C/min. The sample should be small, finely powdered, and packed tightly in a thin-
walled capillary tube of small diameter. The column of solid in the capillary tube should
be just high enough to be seen clearly during melting (about 1-2 mm).
I. Melting Point
A melting point can be used to identify a substance and to get an indication of its purity.
The melting point of a solid is the temperature at which the solid exists in equilibrium
with its liquid state. Experimentally, it is extremely difficult to establish the exact
temperature at which this equilibrium is established; therefore, the temperature range
over which liquid and solid are found to coexist is called the melting point or melting
point range. For example, a solid may be reported to have a `melting point' of 100-
101oC; this means that, on heating slowly, the first droplet of liquid was observed at
100oC and the last crystal of solid disappeared at 101
oC.
A pure crystalline compound usually has a sharp melting point and it melts completely
over a narrow temperature range of not more than 1-2oC. The presence of impurities
produces a lowering of the melting temperature and produces an increase in the width of
the melting point range.
To determine the melting point of a crystalline substance, a small amount of the finely
powdered crystals is introduced into a thin-walled capillary tube; the latter is placed on an
electrically heated "hot-stage" and heated. Two temperatures are recorded, the
temperature at which the substance begins to liquefy and that at which it becomes
completely liquefied. The melting point range is the interval between these two
temperatures. The average of the temperature at which the solid begins to turn to liquid
and the temperature when the entire solid has turned to liquid is the melting point.
To obtain good results it is essential to heat the melting point hot-stage slowly at a
uniform rate, about 2 degrees per minute. To minimize lag in the melting process and
heat transfer, a small amount of material, finely powdered and densely packed in a thin-
walled capillary of small diameter, should be used. The height of solid in the capillary
tube should be just enough to permit good observation of the behavior on melting (about
3-4 mm).
The data obtained from melting point determination is often useful because one can make
observations about the purity of a sample based upon this information. A small melting
point range of less than two degrees Celsius usually indicates a pure substance. Melting
point ranges greater than two degrees Celsius indicate some sort of impurity. Adding an
impurity to a compound lowers the melting point. This allows for comparison of samples
to determine if they are the same compound or different compounds. If two compounds
are mixed and the mixture has the same melting point as one of the components of
32
the mixture, then the two compounds are the same. If the melting point of the
mixture is lower, the two components are different.
To determine the melting point of sample, the sample is finely ground using a mortar and
pestle, a small amount is placed in a capillary tube and packed.
A Mel-Temp™ (shown below) is commonly used to measure the melting points of
compounds. A sample is loaded into a capillary tube and placed into the Mel-Temp™.
While the tube is being viewed through an eyepiece, the Mel-Temp™ gradually heats the
sample. By carefully observing the temperature range at which the sample turns from a
solid into a liquid, the melting point is determined.
Operation
Turn the Mel-Temp power on and adjust the temp control to the desired rate of
heating.
• In order to obtain an accurate mp, it is necessary to heat
S L O W L Y, at a rate of 2-3oC/min.
Thermometer
Sample slot
Eye piece
33
• Heating too fast may lead to inaccurate results because of insufficient time for
heat transfer.
Procedure
Caution: Wear eye Protection!
Part A: Pure Substances
1. Set up the basic melting point apparatus as shown above.
2. In a clean mortar and pestle, grind a small amount of Vanillin.
3. Place a small (pea sized) mound of Vanillin onto a glass plate and using a scoopula
gently force the Vanillin into a capillary tube. This can be accomplished by laying the
tube onto the plate and sliding it into the pile of Vanillin, while keeping the pile from
moving with the scoopula.
As you fill the tube gently tap the capillary tube to cause the particles to pack in the
bottom of the tube. Then, place the tube open-side-up at the top of an approximately one
meter glass tube with its end flat on the bench top (if you are tall) or floor (if you are
short) and drop the capillary tube down the glass tube allowing It to bounce and pace the
particles tightly. Repeat this procedure until the final packed capillary tube has about .3 to
.5 cm of sample. Save the ground sample for later use.
34
4. Insert the capillary tube in the Mel-Temp.
6. Monitor the temperature and sample carefully. Just before the sample begins to melt,
it will constrict. After this happens pay close attention and record the temperature
when the first sign of liquid is evident. Continue to watch the sample until all solid has
disappeared and record this as the final temperature. The average of these numbers is
the melting point.
7. Repeat steps 1 - 6 using Acetamide as the sample of which you will determine the
melting point. If the range is >2oC for either pure sample, repeat the procedure.
8. What was the melting point range in Vanillin and Acetamide?
What does this tell us about the melting point range of pure substances?
Were the melting points of the two compounds close?
Part B: Mixtures
1. Prepare three mixtures of the Vanillin arid Acetamide,
1 part Vanillin: 9 parts Acetamide
5 parts Vanillin: 5 parts Acetamide
9 parts Vanillin: 1 Part Acetamide
To prepare these use the tip of your scoopula to combine 1 part of Vanillin with 9 similar
parts Acetamide and mix thoroughly. This is approximately 10% Vanillin and 90%
Acetarnide. Make the other mixtures in the same fashion.
2. Determine the melting points of each mixture as you did in Part A. 4
3. Fill in the chart on the data sheet provided.
4. What is the melting point range for the three mixtures? How are they related?
5. Graph Melting Point vs. Percent Composition using the graphing program introduced
by your instructor.
You should have five points (Pure Vanillin, 10% Vanillin -90% Acetamide, 50%
Vanillin-50% Acetamide, 90% Vanillin-1O% Acetamide, and Pure Acetamide).
6. What observations can be made about the relationship of melting point to percent
purity?
7. How could one use this technique to determine if two unknown compounds are the
same substance?
35
Melting Point
Melting Point Indicates Purity in Two Ways
1. The Purer the Compound, the Higher the Melting Point
2. The Purer the Compound, the Narrower the Melting Point Range
3. Melting point of A decreases as impurity B is added
4. Eutectic Point is the Solubility Limit of B in A; Thus, it is the
Lowest Melting Point of an A/B mixture
(Note: Sharp melting point – no range at eutectic point)
Solid + Liquid
Solid +
Liq
uidLiquid A + B
Solid A + B
Range
Complete
Melting
First Drop
of Liquid{
mpA
mpB
0% B 0% A
Tem
pera
ture
Eutectic
mpB > mpA
A eutectic point is the manifestation of a eutectic reaction which is an isothermal,
reversible reaction between two (or more) solid phases during the heating of a system, as
a result of which a single liquid phase is produced.
36
Name: ________________________________ Lab section: ____________________
Lab Partner: ___________________________ Lab Desk: _______________________
Data
RANGE
Vanillin ____ to _____ melting point = ______________
Acetamide ____ to _____ melting point = _____________
8. ______________________________________________________________________
Vanillin Acetainide Range melting point
10% to 90% _____ to _____ ________
50% to 50% _____ to _____ ________
90% to 10% _____ to _____ ________
6. __________________________________________________________________
7. __________________________________________________________________
5. Graph on separate graphing paper.
37
Pre Lab Questions:
1. Why are melting points of substances determined?
a._____________________________________________________________________.
b._____________________________________________________________________.
2. A pure substance has a _________________________________________ melting
point.
3. An impure substance added to a pure substance does two things …
a. _____________________________________________________________________.
b._____________________________________________________________________.
4. To determine an accurate melting point on a sample the heating should be done at a
rate of __________________________________________.
38
Aspirin Synthesis & Analysis
Objectives
--To prepare aspirin
--To qualitatively determine the purity of the prepared aspirin
--To calculate the percent yield of the prepared aspirin sample
Introduction
Pure aspirin, acetylsalicylic acid, molecular mass 180.2 u, is an organic ester which is an
important derivative of an acid. The equation for the reaction of salicylic acid and acetic
anhydride is shown below.
Aspirin plays an important role in medicine in today’s society as a pain-killer (analgesic)
and as a fever-reducing agent (anti-pyritic). The purity of an aspirin sample can be
determined qualitatively from its melting point. Pure organic compounds tend to melt
over a very small temperature range (<20 ) and higher temperatures where as compounds
containing impurities have a tendency to melt over a wider range of temperature and the
melting point is lowered. The degree to which the melting point is lowered depends on
the nature and concentration of the impurities.
An important part of this lab is the analysis for the percent yield of the acetylsalicylic
acid that you prepared. To carry out this analysis you must complete the following steps.
Limiting Reagent
The limiting reagent is the reactant not in excess that will limit the final amount of
product you obtain. You should not have more product than you have grams of limiting
reagent.
39
Calculation
Given 4.32 g of acetic anhydride, solve for the number of grams of salicylic acid required
to completely react with the given grams of acetic anhydride. Convert grams of one
reactant to needed grams of the other reactant.
If you do not have 5.84 g of salicylic acid, which is needed to completely react with 4.32
g of acetic anhydride. Then salicylic acid is the reactant that limits the amount of product.
Theoretical Yield
The theoretical yield is the maximum amount of product that may be obtained based on
your limiting reagent and the balanced chemical equation, assuming no loss of product.
Calculation
Actual Yield
This is the actual amount of product collected and weighed alter the reaction has been
conducted. Do not be shocked if the actual amount of aspirin is less than that predicted by
the theoretical yield.
Possible reasons include:
1. Reaction doesn’t go to completion due to side reactions or the establishment of
equilibrium
2. All of the product may not have been recovered. (i.e. product sucks to sides of
flasks).
3. Student error - poor lab technique
40
What could have occurred if you have more product mass than you expected (i.e. actual
yield is greater than theoretical yield)
1. The reaction goes virtually 100% and the product has trapped Impurities from
the by-products in the reaction solution.
2. Product not completely dry
3. Starting reactant mass error
It is extremely important to get the specified amount of reactant. Be careful to be
accurate in measuring out sample.
Percent Yield
Percent yield compares the actual yield to the theoretical yield. This value tells you how
well you performed your experiment.
Procedure
A. Preparation of Aspirin
1. Mix the Starting Materials and Heat
Use the tare function on the balance and weigh (+/- 0.01g) 2.0 grams of salicylic acid
in a dry 125 mL Erlenmeyer flask. Carefully add 4.0 mL of acetic anhydride (located
in the hood) and gently swirl contents of the flask. Add 5 drops of concentrated
sulfuric acid to the mixture, swirl, and heat the flask in a boiling water bath for 10
minutes.
41
2. Cool to Crystallize the Aspirin
Remove the flask from the water bath (be careful, it is hot) and add 10 mL of ice water
to the flask to decompose any excess acetic anhydride. Chill the solution in an ice bath
by putting a 1 inch layer of ice in the bottom of a 400 or 600 mL beaker, set the flask
on the ice and packs more ice around the flask. Swirl the flask to help speed
crystallization. Continue to chill the solution until crystallization is complete,
occasionally stirring with a clean glass stirring rod to decompose residual acetic
anhydride. Upon completion of crystallization, add 30 mL of water and swirl for 60
seconds.
3. Separate the Solid Aspirin from the Solution
Set up a vacuum infiltration apparatus (see figure). Make sure you clamp the filter flask
to a ring stand to prevent the flask from turning over. Use the heavy wall tubing to
attach the filtration flask to the vacuum outlet. Place a filter paper circle in the top of
the Buchner funnel, turn the vacuum on and seal the paper with a little water from your
wash bottle.
42
Swirl the contents of the flask containing your prepared aspirin and carefully pour the
mixture into the funnel at a moderate rate. Take care to evenly distribute the contents
over the filter paper, avoiding the edges of the filter paper as much as possible.
4. Washing the Aspirin
Pour 15 mL of cold water over the crystals on the filter paper to wash. Allow the
vacuum to pull air through the crystals for at least five minutes to dry the crystals as
best possible.
5. How Much Did You Prepare
Determine the mass of a watch glass. Carefully remove the filter paper with product
from the funnel, transfer the product to the watch glass, and determine the mass of the
watch glass and product.
B. Qualitative Analysis
Take three small 10cm test tubes. To each test-tube add lml of ethanol 1 drop of 1% iron
chloride solution. To one test-tube add a few crystals of salicylic acid. Iron (III) Chloride
Is used as a test for phenols. Salicylic add has both a phenol group as well as a carboxylic
acid group. To the second test-tube add a few crystals of your product from part A. To
the third tube add a few crystals of pure acetylsalicylic acid. Record your observations on
your data sheet. In your conclusion explain what these observations indicate.
43
Name: _______________________________ Lab Section: _________________
Lab Partner:___________________________ Lab Desk: ___________________
Calculation Data Sheet
Be sure to show all work on calculations and record final answers on data sheet.
1. Limiting Reagent
a. Mass of SA needed to react completely with (your mass here) of AA.
b. Mass of AA needed to react completely with (your mass here) of SA.
c. Based on calculations in above, what is the limiting reagent in this reaction?
. .
2. Determine Theoretical Yield
3. Calculate Percent Yield
44
Name: _______________________________ Lab Section: _________________
Lab Partner:___________________________ Lab Desk: ___________________
Data
A. Preparation of Aspirin
(1) Mass of Salicylic Acid g
(2) Mass of Acetic Anhydride (4.0 ml X 1.08g/ml) g
(3) Mass of Watch Glass g
(4) Mass of Aspirin + Watch Glass g
(5) Mass of Aspirin g
B. Analysis
Record the color changes observed.
. . .
SA Product ASA
C. Calculations
(1) Limiting Reagent Substance
and grams used . g
(2) Theoretical Yield of Aspirin g
(3) Actual Yield of Aspirin g
(4) Percent Yield of Aspirin %
45
Prelaboratory Assignment
Name____________________ Section_______ Date___________
1. Zinc oxide, ZnO, which is used as an antiseptic and as a white pigment in paint can be
prepared by burning zinc. A sample of zinc metal of mass 125g is burned in excess
oxygen. The amount of ZnO collected was 131g. Calculate the percent yield of ZnO.
2. Carbon disulfide, CS2 is a very flammable substance that reacts with 02 to form CO2
and SO2 When 9.34g of CS2 was burned in excess oxygen, 5.OOg of CO2 was collected.
Calculate the percent yield of CO2.
3. What is the percent yield of Al2 (SO4)3 if 27.Og of A1 (OH)3 is reacted with excess
H2SO4 and 55.0g of Al2 (SO4)3 is collected?
4. In the preparation of aspirin 7.50g of salicylic acid was reacted with excess acetic
anhydride. The student collected 7.50g of aspirin. Calculate the percent yield of aspirin.
46
Chemical Kinetics
Objectives
--To determine the rate law for a chemical reaction.
--To determine the effect of temperature change on rate for a chemical reaction.
--To use graphical and algebraic analysis of the experimental data.
Introduction
Chemical kinetics is the study of -chemical reaction rates, how reaction rates are
controlled, and the pathway by which a reaction proceeds from -its reactants to its
products.
There are several factors which Influence reaction rates:
--the nature of the reactants
--the surface area of the reactants
--the temperature of the chemical system
--the presence of a catalyst
--the concentration of the reactants
In this exercise we will determine the rates for several chemical reactions and observe the
effects temperature and concentration changes have on these reaction rates. We will
summarize these concentration change effects in an experimentally derived rate law.
The rate of a chemical reaction is expressed as a change in the [ ] of a reactant or product
as a function of time. Consider the following equation.
aA + bB cC
The rate of this reaction is related, by some exponential power, to the concentration of
each reactant.
Rate = k[A]m
[B]n
This expression is called the rate for the reaction.
--K is called the specific rate constant which varies with temperature, but is not affected
by the concentrations of the reactants.
47
--[A] [B] represent the molar concentrations of species A and B.
--The superscripts m and n are called the orders of the reaction with respect to each
reactant. The most common reaction orders include 0, 1, and 2; m and n are always
determined experimentally. If the reaction order is zero with respect to any reactant,
this is an indication that the change in [ ] of this particular reactant has no effect on the
reaction rate and is therefore omitted from the rate law expression.
Example: Rate = k[A]2 [B]
0
Rate = k[A]2
You will determine the values of m and n using two different methods.
1. Graphical analysis
2. Algebraic determination
Experimental Procedure
Procedure A
1. Solution Preparation
One partner should prepare Solution A while the other partner prepares Solution B. (very
accurate throughout this lab.)
Solution A
a. Clean an l00mL volumetric flask; add 10mL of distilled 1120, swirl, and drain;
repeat 3 times. The flask does not have to be dry.
b. Measure out 2.00g of sodium hydrogen su1f NaHSO
c. Measure out 0.300g of sodium sulfite, Na
d. Place both solids into the clean volumetric flask.
e. Fill the flask half-full of distilled H20 and swirl the contents until the solids
have dissolved.
£ Add more distilled H until the bottom of the meniscus is on the l00mL mark.
Use a dropper for the last few drops.
g. Parafilm the flask and invert 4 or 5 times with careful shaking to thoroughly
mix the solution, (Critically Important!)
h. Transfer the solution into a clean, thy 25OmL beaker and label the beaker
standard solution “A”.
48
Solution B
a. Same as (a) above.
b. Use a clean l0mL graduated cylinder and measure 9.0mL of 37% formaldehyde
(formalin) solution.
c. Put the 9.0mL of formalin into your clean l0mLvolumetric flask.
d. Fill the volumetric flask half-full using distilled H20, and swirl to mix the
contents.
e. Add more distilled H20 until the bottom of the meniscus is on the l00mL mark.
Use a dropper for the last few drops.
f. Parafilm the flask, Invert several times (4 or 5) with careful shaking to
thoroughly mix the solution. (Critically Important!)
g. Transfer the solution into a clean, dry 25OmL beaker and label the beaker
standard solution “B”.
Caution. Formaldehyde is toxic. If you get any on your hands, wash them
thoroughly. As a precaution wash your hands after solution preparation and
before you leave the lab.
2. Preparation of Test Solutions
The reaction will be performed using seven combinations of standard solution A and
standard solution B. A constant total volume of l00mL will be used in each case. The
table (next page) summarizes the preparation of the 7 test solutions.
--Run one test reaction at a time.
--Measure the volumes of the standard solutions A and B with a clean pipet.
--Transfer standard solution A to a clean, dry 250mL beaker.
--Transfer standard solution B to a clean, dry l00mL beaker.
--You may use a clean 50mL graduated cylinder to transfer the appropriate distilled H
volumes to test beakers A and B.
--You will find the phenolphthalein in dropper bottles.
49
Composition of Test Solutions
Make sure the separate solutions are mixed thoroughly using a clean, dry stirring rod.
3. Preparation for the Reaction
--The reaction begins when the formalin (Soln. B) is added to Solution A.
--Be prepared to time the reactions in seconds (You will need a watch with a second
hand.)
--Place dilute Solution A (250mL beaker) on a white sheet of paper so the color change
(faint pink) is more easily detected.
--As one student mixes the solutions, the other notes the time.
4. Timing the Reaction
--Rapidly add Solution B to Solution A. Don’t “slosh”. START TIME and swirl the
contents of the mixture.
--Be ready to STOP TIME at the appearance of color throughout the entire reaction
mixture
--Record the time lapse (start to finish) to the nearest second on the report sheet.
5. Run all test reactions using the above procedures.
--Use appropriate volumes (as listed in table).
--Be sure to use clean glassware (rinse all glassware twice with a small portion of
distilled H
--Repeat all test reactions for a second trial.
50
Procedure B
The Effect of Temperature on Reaction Rate
1. Increase Temperature: Test Solution #8
--Using the volumes listed for Test Reaction #1 in the Table on the previous page repeat
procedure A parts #2 and #3.
--Before the solutions are mixed, place both beakers in a pan of hot tap H20. Hold
beakers to prevent them from turning over.
--Place a thermometer in beaker “A” solution. When the temperature is at least 10°C
above room temperature, mix the solutions with timing as• before.
--Record the temperature at the moment of mixing.
_____°C _____K
-- Record time lapse.
2. Decrease Temperature: Test Solution #9
--Repeat the above procedure but pack Ice around the beakers to cool the solutions to
somewhere in the 5°C range before mixing.
--Record temperature at the moment of mixing.
_____°C _____K
--Record time lapse.
Calculations for Determining the Rate Law
Perform the calculations, one step at a time. As you read through this section, complete
the appropriate calculation and record your answer for each test solution on the Report
Sheet.
1. Calculate Rate: The rate (in sec-1
) is calculated as the reciprocal of the time.
Calculate the Rate for 9 Test Reactions. Show a calculation set-up for the rate of Test
Solution #1 only, on the calculation data sheet.
51
2. Calculate Molar Concentrations: Solution A, HSO3-
Solution B, H2CO
A. Original Standard Concentrations
No calculation necessary here.
B. Dilute Test Solutions: Solution [A] and Solution [B]
Use the Formula M1V1 = M2V2
Show a sample calculation on the calculation data sheet for Reaction 2 Solution [B].
Note Original Standard Concentration = 1.20M H2CO
52
3. Determination of the Rate Order, m and n
A. Exponent m: Graphical Analysis
--If you plot time vs. [A] and discover that the plotted data appears to be a straight
line through the origin, the value of the exponent m=0, zero order in relationship
to [A], has no influence on rate and is therefore omitted from the rate law
expression.
--If you plot Rate vs. [A] and discover that the plotted data appears to be a straight
line through the origin, the value of the exponent m=1, the reaction is 1st order in
relationship to [A] if the reaction is 1st order in a particular reactant, the
concentration change and rate change will be proportional.
--If you plot Rate vs. [A]2 and discover that the plotted data appears to be a straight
line through the origin, the value of the exponent m=2, the reaction is 2nd order in
relationship to [A]. When the rate law is second order in a particular reactant,
doubling the concentration increases the rate by a factor of 22 = 4; tripling the [ ]
causes the rate to increase by a factor of 32= 9.
Determine the value of the Exponent m
1. Plot Time vs. [A] for Test Solutions 1, 5, 6, and 7 only. Is this a zero order case?
If the answer is yes go on to the determination of the Exponent n. If your answer
to the question is NO, try Step 2 below.
2. Plot Rate vs. [A] for Test Solutions 1, 5, 6, and 7. Is this a first order case? If not
go to Step 3 below.
3. Plot Rate vs. [A]2 for Test Solutions 1, 5, 6, and 7. Is this a second order case?
Include all graphs in your final write-up. Record the determined value of m on the
Report Sheet.
B. Exponent n: Algebraic Determination
1. Take the ratio of the rates from two test reactions in which the [B] is changing
and [A] is held constant.
53
2. Using the rate law, rate = k[A]m
[B]n plug in the appropriate concentrations from
the same test reactions (in the same order) used in Step #1.
Show a work set-up using your values for rates and [ ]’s on the calculation data sheet.
Use Test Reactions #1 and #2 as shown above. Record the calculated value of n on the
Report Sheet.
4. Determination of k, the Specific Rate Constant for the Reaction:
--Using your rate law expression, substituting in (and using) the values you have found
for the exponents, rearrange the equation to obtain k, the rate constant.
--After you rearrange the equation, k will be on the left of the = mark and all other
factors will be on the right.
--Remember that [X]0 = 1, therefore any zero order reactant is not included in the
expression.
Calculate k for all test reactions that are not zero order. Show a calculation set-up for
one of the k values you determined on the calculation data sheet. Record all results on
the Report Sheet.
5. Calculate the average value of k and standard deviation. Show your work on the
calculation data sheet. Record answers on the Report Sheet. (You will need the standard
deviation calculation found In the Appendix.)
54
Name: _______________________________ Lab Section: _________________
Lab Partner:___________________________ Lab Desk: ___________________
DATA
55
Post Laboratory Questions:
1. What happens to the reaction rate as the temperature of the test solution is increased?
2. What happens to the reaction rate as the temperature of the test solution is decreased?
56
Kinetics Calculation Data Sheet
Name: .
1) Rate Calculation:
Test Solution #1
2) Dilute Test Solution:
Test Reaction #2 Solution [B]
3) Exponent n: (Algebraic Determination)
Use data from Test Reactions #1 & #2
4) K Calculation:
Show one calculation for any test reaction that is not zero order.
57
5) Average Value of K Calculation: (No Zero Order Reactions Included)
6) Standard Deviation Calculation:
Step 1 =
Step 2 =
Step 3 =
Step 4 =
Step 5 =
58
59
60
Equilibrium
I. Objectives
--To observe the effects concentration and temperature have on a system at
equilibrium.
--To observe the effect that a strong acid or strong base addition has on a buffered
and a non-buffered system. (The Common Ion Effect)
I. Introduction
Many chemical reactions do not go to completion, (do not produce 100 percent
yield product). These reactions come to a point of chemical equilibrium. At the
point of equilibrium there are two reactions taking place at the same time in
opposite directions. The two reactions are indicated by two arrows in the equation
<---->. When a system has reached equilibrium, the rate of the forward reaction is
equal to the rate of the reverse reaction, and the concentrations of all species
remain constant.
Equilibrium is a dynamic (continuous) state that will continue unless some stress
is applied to the system. This stress may be in the form of concentration changes,
temperature changes, or pressure changes when one of the species in the system is
a gas. The French chemist Henri Louis Le Chatelier (1850-1 936) explained what
happens to the equilibrium when one of these stresses is applied. Le Chatelier’s
principle states: A system at equilibrium responds to an imposed stress in such a
way as to relieve the stress.
Example System at Equilibrium
Applied Stress:
1. An increase in nitrogen gas.
2. A decrease in ammonia gas concentration.
3. An increase in temperature.
4. A decrease in pressure.
Results
1. Adding N2(g) causes the reaction to shift to reduce the amount of N2 added. As a result
the reaction shifts to the right producing more ammonia (and consuming some hydrogen)
until a new equilibrium is established.
61
2. Decreasing NH3(g) causes a stress on the product side of the system. The N2 arid H2
react to replace the NH3 that has been removed; this relieves the stress and the
equilibrium is shifted to the right
3. The example system in the forward direction is exothermic--heat is given off and
shown as a product. If the temperature is increased the reaction will shift in the direction
that will use up some energy and lower the temperature (i.e. the reactant side). The shift
is to the left.
4. A decrease in pressure would cause a shift to the side showing the greatest number of
moles of gas molecules--to the left (reactants).
Summary of the Effects of Changes on System Equilibrium
Change
Concentration of one of
the reactants is increased
Concentration of one of
the reactants is
decreased
Concentration of one of
the products is increased
Concentration of one of
the products is
decreased
Pressure is increased (by
decrease in volume)
Equilibrium
Shifts to the right
Shifts to the left
Shifts to the right
Shifts to the right
Shifts in the direction
that produces the
smaller # of gas
molecules. If the # of
gas molecules is the
same on both sides of
the equation, the
equilibrium will not be
affected.
K
No effect
No effect
No effect
No effect
No effect
62
Change
Pressure is decreased
(by an increase in
volume)
Temperature is
increased
Temperature is
decreased
Equilibrium
Shifts in the direction
that produces the larger
# of gas molecules. If
the # of gas molecules is
the same on both sides
of the equation, the
equilibrium will not be
affected
Exothermic reactions
shift to the left;
endothermic reactions
shift to the right
Exothermic reactions
shift to the right;
endothermic reactions
shift to the left
K
No effect
Increases for
endothermic reactions;
decreases for exothermic
reactions
Decreases for
endothermic reactions;
increases for exothermic
reactions
III. Experimental Procedure
Perform this experiment with a partner. At each numbered superscript in the
procedure STOP arid record your observations on the Report Sheet Discuss your
observations with your lab partner. The chemical equation will be a great help to you
in this experiment.
A. Chromate-Dichromate Equilibrium
The position of some equilibria are easily shifted by adjusting the acidity (or
basicity) of the solution.
1. Pour 3mL of 0. 1M potassium .chromate into a small clean test tube.
2. Acidify the system by adding 1M sulfuric acid drop-wise until a color change is
observed. Account for your observations.
3. Make the solution basic by adding 1M sodium hydroxide drop-wise until a
color change is observed.
63
Disposal Information: Dispose of the waste chromate-dichromate solution in the “waste
chromium salts” container located in the hood.
B. Acetic Acid-Acetate Ion Equilibrium (A Buffer System
The effect of adding an ion or ions common to those already present in a system at a
state of dynamic equilibrium is called the common ion effect the effect is observed
below.
1. In a small clean test tube, add a few drops of universal indicator to 5mL of 0.1M
HC2H3O2: note the color and pH. Prepare the buffer solution by adding 0.1g of
NaC2H3O2 to the HC2H3O2 test tube. Agitate to dissolve. Compare the color of the
solution with the pH color chart for the universal indicator. Divide the solution into two
equal volumes in separate small test tubes and label #1 and #2.
2. Prepare two water systems by placing 2.5mL of distilled H2O in two small test tubes
and label #3 and #4. Add several drops of universal Indicator to each. What is the H2O
pH based on the pH color chart?
3. Test the buffer and water systems with NaOH by adding lmL (20 drops) of 0.1M
NaOH to test tubes #1 and #3. Compare the colors (and pH) of the two solutions.
4. Test the buffer and systems with HCl by adding lmL (20 drops) of 0.1M HCl to test
tubes #2 and #4. Compare the colors (and pH) of the two solutions.
5. Explain the effect a buffered system (as compared to a non-buffered system) has on pH
change when a strong acid or strong base is added to it.
C. Ammonia Equilibrium
If the [ ] of one of the species in an equilibrium system changes, the equilibrium position
tends to shift in a way that compensates for the changes.
1. Pour 3mL of distilled H2O into a small clean test tube and add one drop of the
indicator phenolphthalein. Observe the color. (Note: Clear is not a color.) Add 4 drops
of IM HCl, mix thoroughly, and observe the color. Add 9 drops of IM NaOH, mix
thoroughly, and observe the color.
2. Pour 3mL of 0. 1M ammonia solution into a small clean test tube; add 1 drop of
phenolphthalein and observe the color.
3. Add lM ammonium chloride drop-wise with mixing until a color change is observed
64
D. saturated Ammonium Chloride Equilibrium
The effect a temperature change has on a system at equilibrium depends on what type of
reaction you are investigating (i.e. exothermic or endothermic). In an exothermic reaction
heat is considered to be a product. In an endothermic reaction heat is thought of as a
reactant.
1. Four 3rnL of saturated ammonium chloride solution into a small clean test tube. Add
l2M HCl (caution: avoid skin contact) drop-wise with mixing until a change occurs. What
change do you observe?
2. Place the test tube from part 1 in a H2O bath and heat until a change occurs. Record
your observations.
3. To a clean, dry test tube add enough solid ammonium chloride to cover the bottom of
the test tube to a depth of approximately 5cm. Hold the bottom of the test tube in the
palm of your hand. Add 3mL of distilled H2O to the test tube. Record what you feel and
see as the ammonium chloride dissolves.
65
Report Sheet--Le Chatelier’s Principle
Name . Lab Sec. . Date .
A. Chromate-Dlchrolflate Equilibrium
1. The color of the CrO42-
ion. .
2. Color change with H2SO4 addition. .
3. The addition of the H2SO4 caused a shift toward reactants or products?
.
4. Effect of NaOH. Explain. .
.
B. Acetic Acid-Acetate Ion Equilibrium (A Buffer System)
5. Color of the universal indicator in HC2H3O2 .
6. Color of the universal Indicator after addition of Na C2H3O2 .
pH .
Effect of NaC2H3O2 on the equilibrium .
.
66
9. When a strong acid or a strong base is added to a buffered solution, acetic acid-sodium
acetate, how does the magnitude of the pH change compare to the pH change that
occurs when the strong acid or strong base is added to pure H20?
C. Ammonia Equilibrium
10. Color of phenolphthalein in H2O .
11. Color of phenolphthalein in acidic solution .
12. Color of phenolphthalein in basic solution .
13. Color of phenolphthalein in aq NH3 soln. .
67
14. Color of phenolphthalein upon addition of NH4Cl .
The addition of NH4Cl caused the (NH4+) to .
This ↑ or ↓ in the [NH4+] caused a shift in the equilibrium toward the .
Noting the color of the NH3/MH4Cl solution (#14), would you say the shift caused the
solution to be less basic or more basic? .
D. Saturated Ammonium Chloride Equilibrium
15. What change was observed? .
This indicates a shift toward the .
The addition of what ion caused the shift?____ ________ ____.
16. Observed change .
A shift occurs toward the .
What caused this shift? .
17. The approximate temperature of the system is (hot--cold?) .
Is the dissolution process of ammonium chloride endothermic or exothermic?
.
Now that you know what type of energy change takes place when NH4Cl Is dissolved,
write a chemical equation using heat as a reactant or product to explain the shift in #16 in
terms of concentration.
68
Pre-laboratory Questions
1. Define chemical equilibrium.
2. What do the arrows mean in the equation for the production of ammonia by the Haber
process?
3. State Le Chatelier’s principle.
4. Predict the direction of the shift in equilibrium:
A. CO32-
(aq) + H2O (l) <----> HCO32-
(aq) + OH- (aq)
Solid NaHCO3 is dissolved in the solution. ____
Hydrochloric acid is added to the solution._ ___
B. C(s) + H2O(g) + energy <----> CO(g) + H2(g)
Carbon monoxide is removed as it forms .
The system is cooled .
69
C. NaNO3(s) <----> Na+(aq) + NO3
2-(aq)
Nitric acid solution is added to the system .
Solid sodium chloride is dissolved in the solution .
5. Give formulas for the following compounds.
a) sodium hydroxide
b) acetic acid
c) hydrochloric acid
d) sodium acetate
e) potassium dichromate
f) nitric acid
g) potassium chromate
h) ammonium chloride
70
Chemicals used:
0. 1M potassium .chromate
1M sulfuric acid
1M sodium hydroxide
Universal indicator
0.1M HC2H3O2
NaC2H3O2
0.1M NaOH
0.1M HCl
Phenolphthalein
IM HCl
0. 1M ammonia
1M ammonium chloride
Saturated ammonium chloride
12M HCl
Solid ammonium chloride
71
72
73
74
75
76
77
78
79
80
Solubility Product Constant and Common Ion Effect
Objectives
--To determine the solubility product constant, Ksp, of Ca(OH)2
--To determine the solubility product constants Ksp, of Ca(OH)2 in the presence of added
Ca2+
Introduction
There are many common salts which• dissolve in water to a very limited degree. We
classic these substances as slightly soluble or insoluble salts. No salt is actually 100%
insoluble in water. A saturated solution of a slightly soluble salt produces a dynamic
equilibrium between the salt and its dissociated ions. It is important to remember that
because these salts are ally slightly soluble the ion concentration will be low in the
equilibrium.
Example In a saturated CaCO3 solution an equilibrium between solid CaCO3 and Ca2+
and CO2-
ions in solution is reached. Due to the low solubility of solid CaCO3 (few Ca2+
and CO2-
ions) we say that the equilibrium lies far to the left.
CaCO3(s) <----> Ca2+
(aq) + C032-
(aq)
We write the equilibrium constant expression for the above reaction in the following
form:
Ksp = [Ca2+
] [CO32-
] = 4.5 x 10-9
--In this expression the Ksp Is called the Solubility Product Constant for the slightly
soluble salt, CaCO3
--The brackets around the ions refer to the molar concentration of each Ion.
--The 4.5 x i0 is the actual Ksp value. The Ksp value is small which means that the molar
concentrations of Ca and C0 are low.
Note: If you know the Ksp value and the Ions dissociate in a 1:1 ratio:
[Ca2+
] = [Co32-
] = ×√4.5 × 10-9
(Ksp value)
= 6.7x10
81
You may have noticed that we omitted the solid CaCO3 in our equilibrium constant
expression. This is a common practice because we consider the concentrations of pure
solids to be constant and therefore do not show them in equilibrium expressions.
Experimental Procedure
A. Ksp of Ca(OH)2
1. Preparation of Solution:
--Dissolve 0.4g of Ca(OH) in l50mL of distilled H20.
--Stir for 5 seconds out of every 30 seconds for 5 minutes to ensure an
equilibrium has been reached.
2. Filter the Solution:
--Place a filter paper circle in the top of a Buchner funnel.
--Attach a length of heavy wall rubber tubing from a filter flask to the vacuum
outlet.
--Clamp the filter flask to a ring stand (bench bar) so the flask does not turn over
during filtration.
--Attach the funnel to the flask arid turn on the vacuum control valve.
--If your filter flask is dry, then proceed to filter all of your Ca(OH)2 solution.
--If the filter flask is wet inside, filter approximately 20mL of the Ca(OH)2
solution, turnoff the vacuum valve, disconnect the flask and swirl the filter
solution around the entire interior wall to pick up any moisture present. Discard
this liquid in the sink. Reconnect the filtering system and filter the rest of the
solution.
3. Volumetric Transfer:
--In two clean Erlenmeyer flasks place 20mL (use a 20mL volumetric pipet) of
the filtered Ca(OH)2 solution.
--Add four drops of phenolphthalein to each flask.
4. Titrate the Solutions:
--Prepare your buret as indicated in the appendix.
82
--Place —30 mL HC1 of about 0.1 M concentration in the buret. See the Reagent
bottle for the precise concentration.
--Titrate each of the samples recording all buret readings. Initial volumes and
final volumes, to the second decimal place.
--It is wise to keep the titrated solution from the previous run to compare the
colors at the endpoint.
B. Ksp and the Common Ion Effect
1. Go to the reagent table and take —60 mL of the Ca(OH)2 solution that is
already 0.0200 M in Ca2+
2. Repeat procedures 2 - 4 from Part A using this new solution.
83
Calculation Report Sheet
Name . Section # .
Procedure A Calculations:
Record all final answers on your data sheet. Show a sample calculation below for the first
run only
1. Using the Concentration (M) of the HC1 used in your titration X the mL of acid used,
calculate the moles of acid used.
2. Use the moles of acid from #1 above and a mole/mole factor to calculate the moles of
OH- ion present.
H+(aq) + OH
-(aq) <----> H2O (l)
3. Use the moles of OH- ion from #2 and the volume of Ca(OH)2 solution pipetted out to
calculate the concentration (M) of the OH- ion.
84
4. Calculate the concentration of the Ca2+
Ion by using the answer in #3 X a mole/mole
factor. Be sure to notice the mole ratio of Ca2+
ions to OH- Ions in Ca(OH)2
5. Write the Ksp expression for your reaction. Calculate the Ksp for each sample
6. Calculate Average Ksp.
85
Procedure B Calculations
Record all final answers on the data sheet. Show a sample calculation below for the first
run only
1 - 3 are done exactly as Procedure A.
4. Use the answer to#3 X a mole/mole factor (remember the ratio of Ca2+
to OH- in
Ca(OH)2) to calculate the Ca2+
ion concentration provided by the dissolved Ca(OH)2
solution.
5. Calculate the total concentration of Ca2+
ion by adding the answer in #4 to the initial
Ca2+
ion concentration provided by the 0.0200 M Ca2+
already present.
6. Write the Ksp expression for your reaction. Calculate Ksp for each sample Use total
[Ca2+
]
7. Calculate Average Ksp.
86
Data Sheet
Name: .
Part A
87
Post Laboratory Questions:
1. In Procedure B the concentration of the Ca2+
ion was Increased compared to the Ca2+
ion in Procedure A. According to Le Chatelier’s principle what should happen to the
OH- Ion concentration? Did this happen?
2. Based on LeChatelier’s Principle, the addition of a common ion would have what effect
on solubility of a slightly soluble substance?
88
Prelaboratory Assignment
1. A Write the equilibrium equation for the dissolving of lead bromide. Indicate
aqueous Ions with (aq) as a subscript and solids with (s) as a subscript.
B Write the Ksp expression for lead bromide.
C. If the concentration of lead Ion is found to be 0.0116 M, what is the
concentration of bromide ion? Show a work set-up.
D. Calculate the Ksp. Show a work set-up.
89
2. Write the equilibrium constant expression for the reaction:
Fe3+
(aq) + HSCN(aq) FeSCN2+
(aq) + H+(aq)
3. Nitrogen reacts with hydrogen to produce ammonia according to the equation
3 H2(g) + N2(g) 2 NH3(g)
Calculate the equilibrium constant if an equilibrium mixture contains 0.45 mol
nitrogen, 0.60 mol hydrogen, and 0.25 mol ammonia per liter at 673 K.
90
Acid/Base Titration
Standardization of HCl
Objectives
--Learn basic titration techniques
--Practice solution preparation
--Practice stoichiometric calculations with titration data
Introduction
Titration is the process of measuring the exact volume of a solution of known
concentration that is required to react completely with a measured volume of a solution of
unknown concentration or a known mass of unknown solid. A solution of accurately
known concentration is called a standard solution. To be considered standard, the
concentration of the solute in the solution must be known to four significant figures.
When an acid or base solution is prepared from stock acid or base, the concentration is
approximately known. However, due to relative purity and limitations on the accuracy
and precision of measuring quantities for solution preparation, the solution’s exact
concentration is not known. For this reason, the process of standardization is used.
Standardization is a laboratory process in which the exact concentration of a solution is
obtained by comparing the concentration of the solution to a primary standard, a dry
substance of known purity.
To standardize a solution, the solution is used to titrate a known mass of primary
standard. The titration continues until the standard is completely reacted with the solute
in the solution. To determine when the equivalence point or the point at which the entire
standard has been reacted, is reached an indicator is used. The indicator is a pH sensitive
dye that changes color as the pH changes. In this experiment an HC1 solution will be
standardized against primary standard sodium carbonate using methyl purple as the
indicator. Methyl purple will change to a gray color as you approach the equivalence
point and will be violet in color when the equivalence point is reached.
Procedures
Part A - Preparation of 0.1 M HC1 Solution
1. Prepare 250m1 0.1 M HCl solution from the 6 M stock solution provided.
Calculation:
(Conc. Stock Soln)(Volume Stock Soln.) = ( Conc. Soln. ) (Volume Soln.)
(6 M) (?ml) = (0.1 M) (250 ml)
91
2. Prepare the solution by adding approximately l00ml distilled water to the 250 ml
volumetric flask, pipetting the calculated volume of stock 6 M HC1 solution into the
flask, and diluting to the mark with distilled water. Apply parafilm to the flask top and
invert and shake to assure mixing.
Part B - Standardization of HC1 Solution
1. Clean and Dry three 125ml Erlenmeyer flasks.
2. Mass three portions of Na2CO3 between .2 and .25 g and record the masses to three
decimal places and identl1 each flask ( ie.; .241 g).
3. In each flask, dissolve the dry solid with approximately 25 ml of distilled water. Add
three drops of methyl purple indicator to each flask.
Note: remember to clean buret as discussed and use 2-3 ml titrant as final rinse.
4. Titrate each sample as discussed in the introduction to the laboratory. Be careful to add
titrant slowly and mix well as you titrate. Do not over-shoot the endpoint. Note that there
will be a change to gray color before the final endpoint of violet.
5. Record the initial and final buret readings on the data sheet.
6. Repeat this process for all three samples.
7. Calculate the concentration of the standardized acid solution from the reaction as
shown below.
8. Determine the average concentration of the HC1 solution.
92
9. Determine the relative precision of your results.
Standard Deviation
Step 1 = Compute the average of the three values.
Step 2 = Subtract each individual value from the average.
Step 3 = Sum the square of the deviations in step two.
Step 4 = Divide step three by n (the number of observations).
Step 5 = The square root of this value is standard deviation.
93
Name: _______________________________ Lab Section: _________________
Lab Partner:___________________________ Lab Desk: ___________________
Data
Standard Deviation (Show your Calculations)
Step 1 =
Step 2 =
Step 3 =
Step 4 =
Step 5 =
94
The Determination of a Molar Mass
Objective
--To determine the molar mass of an unknown solute by observing the difference
between the freezing points of a pure solvent and a solution.
Introduction
A solution consists of a solvent (commonly considered to be the substance present in the
greatest amount) and one or more solutes (the substance(s) dissolved in the solvent). The
addition of a solute to a solvent produces characteristic changes in several physical
properties of the solvent. It lowers the vapor pressure of the solvent, it increases the
boiling point of the solvent, and it decreases the freezing point of the solvent. These
physical properties are called colligative properties because the changes that occur
depend only on the number of solute particles dissolved in the solvent, not the kind of
solute particles
Examples of colligative properties are easily found:
--Salt is used to freeze Ice cream because a salt-ice-water mixture is at a lower
temperature than an Ice-water mixture alone.
--Antifreeze (ethylene glycol) added to the cooling system of your car prevents freeze-up
in the winter and boil-over in the summer. This prevention is possible because the
antifreeze-water solution has a lower freezing point and a higher boiling point than pure
water.
The magnitude of the change in a freezing point, ΔTf, the colligative property that will be
used in this experiment to determine molar mass, is directly proportional to the molality
(m) of the solute in solution. This proportionality is a constant called the freezing point
depression constant Kf, and is specific for a particular solvent. For example, one mole of
solute particles dissolved in 1 kilogram of water lowers the freezing point by 1.86oC, if
the solvent is 1 kilogram of benzene the freezing point depression is 5.12°C. See Table I
for a list of some common solvents and their Kf values.
95
1
Solid Solute Molar Mass Determination
Sample Calculation:
If 6.133g of an unknown solute, dissolved in 55.9g of benzene, results in a solution that
freezes at 2.32°C, determine the molar mass of the solute.
Freezing point of Benzene = 5.5°C.
96
The freezing points of the pure p-dichlorobenzene and the p dichlorobenzene
solution are obtained from a cooling curve--a plot of temperature vs. time (Figure
Page 2). The cooling curve for the pure solvent reaches a plateau at the freezing
point; extrapolation of the plateau to the temperature axis determines the freezing
point. The cooling curve for the solution does reach a plateau; but continues to
decrease slowly as the p-dichlorobenzene freezes. This freezing point is
determined at the intersection of two straight lines drawn through the data points
above and below the freezing point.
Experimental Procedure
A Freezing Point of P-Dichlorobenzene (Solvent
1. Prepare a water bath. Fill a 400mL beaker almost full of tap H Support the beaker
using an Iron support ring plus wire gauze on a ring stand. Gently heat the H20 to near
80°C; it does not need to boil.
2. Determine the mass of PDB (solvent). Get a large test tube from the reagent table. Clean
and dry the test tube. Determine the mass of your large, clean, empty, dry test tube to the
nearest 0.ooig and record the value on your report sheet. Add approximately 15g (1/3 of
tt) of PDB to the test tube and remeasure the mass of the test tube and contents to the
nearest 0.001g. Record the value on your report sheet.
3. Melt the solvent. Clamp the test tube to the ring stand and immerse it in the H bath so
that all the PDB is lower than the H surface. Heat the H so all the PDB completely
melts. Carefully Insert a thermometer into the test tube; stir gently to ensure all the
PDB has melted.
4. STOP heating the solvent. When the temperature of the PDB reaches approximately
65°C stop heating. Remove the test tube from the water (IT IS HOT! Use a piece of
pleated paper towel to secure the top of the test tube). Dry the wet test tube with a paper
towel, being very careful not to spill the HOT PDB on you or your partner.
5. Data for the freezing point of PDB. As the PDB cools in the air record the temperature.
Gently and continuously stir (very important) with the thermometer. When the
97
temperature has dropped to 60°C begin recording the temperature every 30 seconds
until the PDB has solidified to a point where you can no longer stir it. Near the freezing
point crystals will begin to appear and they will increase in amount as cooling proceeds.
Make a note of when crystals of the solid appear in the liquid. Be sure to estimate all
temperatures to O.1°C.
6. Plot the Data On linear graph paper, plot the temperature (°C vertical axis) vs. time
(sec. horizontal axis) to obtain the cooling curve for pure p-dichlorobenzene.
Freezing Point of P-Dichlorobenzene and Unknown Solute
Obtain an unknown from the reagent table and record the number on the report sheet.
1. Mass of Solid Solute. Go to the balance room! Determine and record the mass
of solute (about 2 grams) to the nearest 0.001g). Add the unknown solute to
the test tube of PDB.
2. Data for the Freezing Point of Solution. Determine the freezing point of this
solution in the same way as that of the pure solvent. When the solution nears the
freezing point of the pure PDB, record the temperature to the nearest 0.1°C at
more frequent time intervals (approximately 15 seconds). A “break” in the
curve occurs as the freezing begins, although it may not be as sharp as that for
the pure PDB.
3. Plot the D on the Same Graph. Plot the temperature vs. time data on the same
graph as that for the pure PDB. (Use a different colored writing instrument.)
Draw straight lines through the data points above and below the freezing point;
the intersection of the two straight lines is the freezing point of the solution.
4. Repeat with Additional Solute. Add (another 2 grams) to the large test tube of P
and Sample #1. Determine the total mass of unknown that has been added to the
large test tube. Repeat the freezing point determination procedure as above. Plot
the temperature vs. lime data on the same graph.
Disposal Information. Dispose of the waste PDB solution in the “waste (PDB)” container
found in the hood.
Note: If you remelt the solid, PDB disposal is easier.
98
Name: _______________________________ Lab Section: _________________
Lab Partner:___________________________ Lab Desk: ___________________
A. Freezing Point of P-Dichlorobenzene (Solvent)
1. Mass of test tube + PDB (g) .
2. Mass of test tube (g) .
3. Mass of PDB (g) .
4. Freezing point, from cooling curve (°C) .
B. Freezing Point of P-Dichlorobenzene + Unknown Solute
Unknown Solute Number .
Solution I
1. Mass of vial+ Unknown (g) .
2. Mass of vial after removal of solute (g) .
3. Total mass of solute (g) (Bi - B2)
.
4. Freezing point, cooling curve (°C) .
Solution II
5. Mass of empty vial after removal of solute (g)
.
6. Mass of solute added (g) (B2 - B5)
.
7. Total Mass of solute sample (g) (B3 + B6)
.
8. Freezing point, cooling curve (°C) .
99
100
Cooling Curve Data Sheet
Name .
*Remember to make the necessary time-temp recording change for the solutions.
(Procedure B-2).
101
102
Post-laboratory Assignment
Name . Section # .
Show full calculational set-ups.
1. A 15.68g sample of an unknown solute was dissolved in 100.4g of naphthalene. The
solution froze at 72.8°C. What is the molar mass of the unknown solute?
2. If 3.011g of p-dichlorobenzene is dissolved in 5O.7g of camphor, what will be the
melting point of the solution? (Note: melting point and freezing point are the same.)
3. Why is it not necessary to start the temperature axis at 0°C for the plots in this
experiment?
103
Visible Spectrophotometry
Determination of Iron in Water
Objectives
--To develop solution preparation skills.
--To learn to operate the Spectronic 20 visible spectrometer.
--To use a standard Beer’s law curve to determine the concentration of an unknown.
Introduction
Spectroscopic methods of analysis are based upon the measurement of the
electromagnetic radiation (light) absorbed by a sample at a given wavelength. Absorption
spectroscopy is based upon the decrease in power of a beam of electromagnetic radiation
as it passes through a sample. This principle is useful because the amount of light
absorbed by a sample is directly related to the concentration of the absorbing species and
to the path length (how far the light must travel through the sample). Two basic terms are
used in visible spectroscopy to describe the change in power of the list beam: (1.)
Absorbance the amount of light absorbed as the beam passes through the sample, and
(2) Transmittance the amount of light that is allowed to path through the solution. Both of
these are shown on the spectrophotometer readout.
Because the relationship of absorbance and concentration is linear, a standard curve can
be created; from this curve an unknown concentration can be determined. To create a
standard curve one must prepare at least three different standard solutions within the
range of interest. Then, the absorbance or transmittance of each solution is determined
using the Spectronic 20. This data allows a standard curve to be created by plotting
Absorbance versus Concentration. The slope-intercept formula (y = mx + b) for the
resulting line is determined and after determination of the Absorbance (y} of the
unknown, the Concentration (x) of the solution can easily be calculated using this
formula.
In the experiment that follows you will develop a red-orange complex between Fe (III)
and orthophenanthroline. The intensity of the color will be directly related to the
concentration of iron in the solution. Absorbance measurements will be taken at a
wavelength setting of 510 nm on the spectrophotometer.
104
Experimental Procedure
A. Preparation of Standards
From the 50 ppm stock iron solutions prepare 0.25 ppm. 1.0 ppm, and 2 ppm iron
solutions,
1. To evaluate the volume of stock iron solution to pipet into the l00mL volumetric
flask. Use the dilution formula below.
(Conc. stock soln.) (? mL stock soln.) = (desired conc.) (l00mL)
2. To prepare each standard solution, transfer the calculated volume of stock iron
solution into the l00mL volumetric flask. Add 2.00 mL of sodium acetate, 2.00 mL
of hydroxylamine hydrochloride and 5.00 mL of orthophenanthroline to each
standard. Using distilled water, dilute the resulting solution to 100 mL, cover with
para-film, and Invert to mix.
3. Obtain the 5mL of the unknown. Pipette the volume specified on the test tube label
and add the remaining reagents as above.
Note: The unknown is used In place of the stock Iron solution above.
4. Prepare the spectrophotometer as shown in the instructions below
B. Instrument Preparation
(See Diagram that follows for Hints on Use)
a. Turn on the instrument using the zero control on the front of the meter.
Allow the instrument to warm up for at least one-half hour.
b. Use the wavelength control on top of the meter to set the instrument to the
desired wavelength
Calibrate the Instrument
a. With nothing in the closed sample compartment, use the zero control to
obtain a 0 %T reading. Position your eye so that the meter needle is
directly above its reflection in the mirror scale.
105
b. Rinse a small test tube (cuvette) with a blank solution or deionized water.
Fill the cuvette about one-half full with the blank solution, eliminate any air
bubbles, and gently dry the outside with a tissue.
c. Place the cuvette into the sample compartment, rotating the cuvette so that
the alignment mark on the sample compartment lines up with the line on the
cuvette. Verify that the cuvette is inserted to its maximum depth.
d. Close the sample compartment, and use the full-scale control on the front of
the meter to obtain a 100 %T reading.
Use the Instrument
a. Rinse and half fill the cuvette with the solution to be measured.
b. Insert the cuvette as before, close the sample compartment and quickly
obtain the %T reading. Remove the cuvette.
c. Repeat the steps as before for your remaining solutions to be measured.
Remove the cuvette when finished.
d. If the instrument will not be used within one to two hours, turn it off.
Otherwise turn the zero set control fully counterclockwise without turning
off the instrument.
e. Wash and rinse all the cuvettes. Dry them in art inverted position.
5. Determine the absorbance of each of the solutions.
6. Plot a graph of Absorbance vs. Concentration. Determine the concentration of the
unknown solution as described in the Prelab session. Determine the slope of the line
of best fit.
Using the Spectronic 20
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107
Name: _______________________________ Lab Section: _________________
Lab Partner:___________________________ Lab Desk: ___________________
Data
Concentration Absorbance
.25 ppm .
1.0 ppm .
2.0 ppm .
Unknown .
Unknown Number = .
Volume of unknown pipetted .
Unknown Concentration .
Slope .
