laboratory manual for general chemistry...

CHE 1402

Last update: June 2011 1

School of Science & Engineering

LABORATORY MANUAL FOR

GENERAL CHEMISTRY II

Last Update: December 2015

CHE 1402

Last update: June 2011 1

Name: ________________________

Section: ________________________

LABORATORY MANUAL FOR

GENERAL CHEMISTRY II

Last Update: December 2015

CHE 1402

Lab Manual i

TABLE OF CONTENTS

Preface ii

Laboratory Safety iii

Experiment 1:

Boyle’s law and molar mass of a vapor 1

Experiment 2:

Determination of R: the gas law constant 6

Experiment 3:

Colligative properties: freezing point depression and molar mass 10

Experiment 4:

Rates of chemical reactions, a clock reaction 16

Experiment 5:

An equilibrium constant 25

Experiment 6:

Titration curves of polyprotic acids 34

Experiment 7:

Determination of the solubility-product constant for a slightly soluble salt 38

Experiment 8:

Molar solubility, common-ion effect 43

Experiment 9:

Determination of orthophosphate in water 48

Experiment 10:

Galvanic cells 52

Appendix 56

CHE 1402

Lab Manual ii

PREFACE

Chemistry is an experimental science that relies upon accurate measurements and observations

from scientists. Thus, it is important that students of chemistry do experiments in the laboratory

to more fully understand that the theories they study in lecture and in their textbook are

developed from the critical evaluation of experimental data. The laboratory can also aid the

student in the study of Science by clearly illustrating the principles and concepts involved.

Finally, laboratory experimentation allows students the opportunity to develop techniques and

other manipulative skills.

The laboratory is designed to support and illustrate chemical concepts studied in the lecture

portion of the course, as well as to introduce important laboratory techniques and encourage

analytical thinking. The sequence of experiments in this laboratory manual is designed to follow

the lecture notes. However, we cannot guarantee the synchronization of lectures and laboratory

experiments. For this, certain experiments may come before or after the material that has been

covered in the lecture.

The lab manual contains background information and procedures for the experiments you will

perform as part of your General Chemistry II course – CHE1402.

Along with concepts and chemistry covered in the lecture, the laboratory portion of the course

will present some additional chemistry, both theoretical and practical (e.g. water analysis).

Questions are presented throughout each experiment. It is important that you try to answer each

question as it appears in the manual, as it will help you understand the experiment as you do it.

There is a lot of interesting chemistry to explore in the CHE 1402 Laboratory Manual. It is our

hope that you enjoy learning it and that it will enhance your understanding of the material

presented in lecture.

Remember that what you get out of your laboratory experience will be directly

proportional to what you put into it.

CHE 1402

Lab Manual iii

SAFETY IN THE LABORATORY Safety in the laboratory must be emphasized. The compounds you will work with do have some hazards

associated with them. Therefore, it is important to follow the safety rules outlined in this lab manual. You

should assume that all compounds encountered in the laboratory are toxic and handle them accordingly.

Safety goggles for eye protection are recommended and lab coats are to be worn by all students at all

times when entering the laboratory. Many chemicals, common in chemical laboratories, will make holes

in clothing. Always wash your hands thoroughly when leaving the laboratory. The location and use of the

safety equipment in laboratory were already discussed in CHE1401 and will be reminded by your

instructor the first day the laboratory class meets. You should become familiar with the proper use of the

safety shower, eye-wash fountain, fire blanket and fire extinguisher.

Report any accidents which occur immediately to the laboratory supervisor (Dr. S. El Hajjaji).

Safety rules to be strictly followed by all students.

1. Wear goggles when required.

2. Do not touch chemicals with your hands. Spatulas will be provided for handling solid materials.

3. Do not eat or drink in the laboratory.

4. Do not taste any chemical.

5. Do not smell any chemicals directly. Use your fingers to waft the odor to your nose.

6. Do not pipet solutions by mouth. Rubber pipet bulbs are provided at each lab station.

7. Do not put flammable liquids near an open flame.

8. When heating a test tube, make certain that the open end of the tube is directed away from the

students.

9. When finished with your Bunsen Burner for a given portion of an experiment, turn it off.

10. Do not sit on the lab benches.

11. Do not engage in games in the laboratory. Failure to follow this rule will result in immediate

dismissal from the lab and subsequent conduct action.

12. Do not pour any chemicals into a sink without authorization from the instructor.

13. Notify your instructor if a mercury spill should occur.

14. All broken glassware should be cleaned up immediately. The instructor should be notified of all

breakage, especially if a thermometer is involved.

15. Do all reactions involving malodorous, noxious or dangerous chemicals in a fume hood.

16. If a chemical gets on your skin, immediately wash the affected area with large quantities of water.

The instructor should be notified; no matter how insignificant the incident might seem.

17. When pouring one liquid into another, do so slowly and cautiously. To dilute an acid, pour the acid

into the water; never pour water into an acid.

18. No student shall be permitted to work alone in the lab, you should be supervised by a laboratory

instructor (or the lab technician during make up sessions).

19. Exercise good housekeeping practices in the laboratory. Be sure that the lab benches remain free of

disorder during the experiment. In the event of a spill, clean the area immediately and be sure to use a

wet sponge to wipe off the work station at the end of the lab session.

20. Know what you have to do before entering the lab. Read the experiment carefully before coming to

the laboratory.

For more information, a booklet titled “Student’s Chemistry Laboratory Safety Manual” will be

provided to you in your first lab session. Please get acquainted with it.

Be cautious and think about what you are doing !

CHE 1402

Lab Manual 1

Figure 1.1

EXPERIMENT 1

Boyle’s Law &

Molar Mass of a Vapor

OBJECTIVES

To observe how changes in pressure, for a fixed amount of a trapped gas at constant temperature, can affect

the volume of the gas.

To determine the molar mass of a gas based on a knowledge of its mass, temperature, pressure, and volume.

Relates to chapter 10 of “Chemistry the Central Science, 12th Ed.”.

APPARATUS AND CHEMICALS

gas-law demonstration apparatus

balance

125-mL Erlenmeyer flask

600-mL beaker

boiling chips

250-mL graduated cylinder

Bunsen burner

ring stand and iron ring

buret clamp

wire gauze

3-mL sample of a volatile

unknown liquid

DISCUSSION

The Effect of Pressure on the Volume of a Gas The effect of the pressure on the volume of a gas can be determined by using a gas

buret connected to a U shaped manometer as shown in Figure1.1. When the stopcock

is opened, air can enter in the buret, and the level of mercury will be equal in both

arms of U shaped manometer. If the stopcock is then closed, a fixed volume of air is

trapped in the buret at atmospheric pressure. Raising the leveling of mercury increases

the pressure on the gas (enclosed air); the new pressure on the gas corresponds to the

atmospheric pressure plus the difference in height of the mercury in both arms.

At constant temperature, the volume of a given amount of gas is inversely proportional

to the pressure. This is Boyle's law that may be expressed mathematically as

2211 VPVP [1]

Therefore, one can calculate the volume of a gas at any pressure, provided that

the initial volume of the gas at a given pressure is known.

Behavior of Gases: Molar Mass of a Vapor

The ideal gas law can be expressed as follows:

nRTPV [2]

R= 62400 mL-mmHg/mol-K or, in other units, R = 0.0821 L-atm/mol-K. The number of moles of a

substance (n) equals its mass in grams, m, divided by the number of grams per mole (that is, its molar mass,

M):

n = m/M. Making this substitution into Equation [2] gives

RTM

mPV

[3]

CHE 1402

Lab Manual 2

The units of P, V, T, and m must be expressed in units consistent with the value of R.

The first part of this experiment will be a demonstration of Boyle's law. You will vary the pressure of a fixed

amount of trapped air at constant temperature, using a special gas buret.

The second portion of this experiment concerns the determination of the molar mass of a volatile liquid using

equation [3]. A small quantity of liquid sample is placed in a pre-weighed flask and vaporized so as to expel

the air from the flask, leaving it filled with the vapor at a known temperature (temperature of water). The flask

plus vapor is then cooled so that the vapor condenses. The flask plus condensed vapor plus air is then

weighed. The mass of air, being nearly identical before and after cancels out and allows one to determine the

mass of the vapor. The above data, in conjunction with the volume of the flask, permit the calculation of the

molar mass.

PROCEDURE :

A. Verification of Boyle’s law

Step 1:

As shown in Figure 1.3a, with each leg of a U-tube manometer exposed to the atmosphere, the height of

liquid in the columns is equal. At this moment, close the stopcock and record the height h1 that exists

between the stopcock and the level of mercury. h1 corresponds to the height occupied by the gas (air) that is

trapped in the left leg of the manometer, at pressure P1 (measure P1 by means of a barometer).

Step 2:

Next, shift the position of the right leg, either upwards or downwards; this will create a difference of

pressure between the trapped gas in the left leg and the atmospheric pressure in the right leg. The new

height and pressure of the trapped gas are named h2 and P2. Two situations can then occur (Figure 1.3b):

- h2 > h1 which means that P2 < P1. In that case, the difference of heights between the right leg and the left

leg, GP, called the gauge pressure, is negative.

- h2 < h1 which means that P2 > P1. In that case, the difference of heights between the right leg and the left

leg, GP, is positive.

Figure 1.3a Figure 1.3b

Gas is expanded Gas is compressed

CHE 1402

Lab Manual 3

Figure 1.2

Once P1 and GP recorded, you can calculate P2: P2 = P1+ GP.

Finally, verify Boyle’s law:

2211 VPVP

2211 hSPhSP

2211 hPhP

You will consider that the law is verified if the percent error is lower than 5%.

Percent Error:

100%2211

2211

hPhP

hPhPerror

B. Molar Mass of a Vapor

Overview: A sample of a volatile liquid is added to a pre-

weighed flask. The flask is submerged in a boiling water

bath to vaporize the liquid. Because an excess amount of

liquid is used, the volume of vapor produced is greater

than the volume of the flask.

Upon heating, the vapor that is created initially pushes the

air out of the flask and then the vapor begins exiting the

flask until the pressure inside the flask is equal to the

atmospheric pressure. The mass of the vapor remaining in

the flask is obtained by reweighing the flask. Additional

measurements are made to determine the pressure,

temperature and volume of the sample.

Procedure: Place a rubber stopper (inside which a glass

tubing is inserted) in the mouth of a clean, dry 125-mL

Erlenmeyer flask; (see Figure 1.2). Remove the rubber

stopper and add approximately 2 mL of unknown to the

flask. Clamp the flask at the top of the neck and immerse

it as deeply as possible in a 600-mL beaker nearly full of

water. Place some boiling chips in the water and heat to

boiling. Record the temperature of the boiling water (3)

and the barometric pressure (4). As the water boils,

watch the liquid in the flask. As soon as all of the liquid

(including any that has condensed in the neck) has vaporized (about 3 min.), remove the flask by means of the

clamp and set it aside to cool. After the flask has cooled to room temperature, wipe it dry and remove any

water that may adhere to the rubber stopper. Weigh the flask, cap, rubber band, and condensed unknown

liquid (5). Calculate the weight of the condensed liquid (6). Remove the cap and fill the flask completely with

water. Measure the volume by pouring the water into a large graduated cylinder (7). Calculate the molar mass

of the unknown using Equation [3] (8).

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Lab Manual 4

REVIEW QUESTIONS

Before beginning this experiment in the laboratory, you should be able to answer the following questions: 1 How does the pressure of an ideal gas at constant volume change as the temperature increases?

2 How does the volume of an ideal gas at constant temperature change as the pressure increases?

3 How does the volume of an ideal gas at constant temperature and pressure change as the number of molecules changes?

4. Write the ideal-gas equation and give the units for each term when R = 0.0821 L-atm /mol-K.

5. Show by mathematical equations how one can determine the molar mass of a volatile liquid by measurement of the pressure, volume, temperature, and weight of the liquid.

6. If 0.75 g of a gas occupies 300 mL at 27°C and 700 mm Hg of pressure, what is the molar mass of the gas?

7. A sample of nitrogen occupies a volume of 300 mL at 30°C and 700 mm Hg of pressure. What will be its volume at STP?

8. Consider Figure 12.1. If the height of the mercury column in the leveling bulb is 30 mm greater than that in the gas buret and atmospheric pressure is 670 mm, what is the pressure on the gas trapped in the buret?

9. Consider Figure 1.3a. If the level of the mercury in the leveling bulb is lowered, what happens to the

volume of the gas in the gas buret?

10. Show that Boyle's law, Charles's law, and Avogadro's law can be derived from the ideal-gas law.

11 Methane bums in oxygen to produce CO2 and H2O.

CH4 (g) + 2 O2 (g) 2 H2O (l) + CO2 (g)

If 3.7 L of gaseous CH4 is burned at STP, what volume of O2 is required for complete combustion? What volume of CO2 is produced?

12 Calculate the density of O2 at STP, (a) using the ideal-gas law and (b) using the molar volume and molar mass of O2. How do the densities compare?

CHE 1402

Lab Manual 5

Experiment 1

Boyle’s Law & Molar Mass of a Vapor

Name(s)

Date Laboratory Instructor

REPORT SHEET

A. Boyle's Law : Effect of Pressure at Constant Temperature

Trial 1 Trial 2

1. First pressure (atmospheric), barometer reading, mm Hg (P1) _______ _______

2. First height, in mm (h1) _______ _______

3. Second height in mm (h2) _______ _______

4. Difference in mercury levels (+ or -), mm Hg (GP) _______ _______

5. Second pressure, mm Hg (P2) _______ _______

6. Percent error _______ _______

(Show calculations)

B. Molar Mass of a Vapor

1. Unknown liquid number ____________

2. Wt. of flask + rubber stopper + cap ____________ g

3. Temperature of water when liquid boiled ____________ °C

4. Pressure of the flask = PATM = ___________ mbar = ___________ mm Hg

5. Wt. of flask + rubber stopper + cap + condensed vapor ____________ g

6. Wt. of condensed vapor ____________ g

7. Volume of flask ____________ mL

8. Molar mass of vapor ____________ g/mol

(Show calculations)

CHE 1402

Lab Manual 6

EXPERIMENT 2

Determination of R:

The Gas-Law Constant

OBJECTIVE Determination of the ideal-gas-law constant R.


APPARATUS & CHEMICALS

0.30 g CaCO3 250 mL beaker

5 mL of 4 M HCl rubber stoppers

test tube rubber tubing

ring stand balance

clamp 125 mL Erlenmeyer flask

pinch clamp thermometer

250-mL Erlenmeyer flask barometer

rubber bulb 100 mL graduate cylinder

DISCUSSION

The ideal-gas equation, nRTPV is very useful in describing the behavior of ideal gases under normal

condition (room temperature and atmospheric pressure). However, the behavior of a real gas can be described

by the van der Waals equation:

nRTnbVV

naP

2

2

where a and b are constants characteristic of a given gas. The term nb is a correction for the finite

volume of the molecules. The term 2

2

V

na is a correction to the pressure which takes into account the

intermolecular attractions.

In this experiment you will determine the numerical value of the gas-law constant R, in its common

units of L-atm/mol-K. This will be done using both the ideal-gas law and the van der Waals equation

together with measured values of pressure, P, temperature, T, volume, V, and number of moles, n, of

an enclosed sample of CO2. An error analysis will then be performed on the experimentally

determined constant. The CO2 will be prepared by the decomposition of CaCO3 in presence of a concentrated solution of HCl:

CaCO3 (s) + 2 HCl (aq) CO2 (g) + H2O (l) + CaCl2 (aq)

The CO2 can be collected by displacing water from a bottle and the volume of gas can be determined from the

volume of water displaced. Through use of Dalton's law of partial pressures, the vapor pressure of water, and

atmospheric pressure, the pressure of the gas may be obtained. Dalton's law states that the pressure of a

mixture of gases in a container is equal to the sum of the pressures that each gas would exert if it were

present alone: i itotal PP

CHE 1402

Lab Manual 7

Since this experiment is conducted at atmospheric pressure, vaporOHCOatmospheretotal PPPP22

.

PROCEDURE

Place 0.30 g of CaCO3 in a test tube and carefully insert another smaller test tube in it containing 5mL of 4 M

HCl (be sure NOT to mix CaCO3 and HCl before the experiment). Assemble the apparatus illustrated in

Figure 2.1 but do not attach the test tube. Be sure that tube B does not extend below the water level in the

bottle. Fill glass tube A and the rubber tubing with water by loosening the pinch clamp and attaching a rubber

bulb to and applying pressure through tube B. Close the clamp when the tube is filled. Attach tube B as shown

in Figure 2.1. When you attach the test tube, half-fill the beaker with water, insert glass tube A in it, open the

pinch clamp, and lift the beaker until the levels of water in the bottle and the beaker are identical; then close

the clamp, discard the water in the beaker, and dry the beaker. The purpose of equalizing the levels is to

produce atmospheric pressure inside the bottle and test tube.

Set the beaker with tube A in it on the desk and open the pinch clamp. A little water will flow into the beaker,

but if the system is air tight and has no leaks, the flow will soon stop, and tube A will remain filled with

water. If this is not the case, check the apparatus for leaks and start over again. Leave the water that has

flowed into the beaker in the beaker; at the end of the experiment, the water levels will be adjusted, and this

water will flow back into the bottle.

By opening the pinch clamp, mix the solids (CaCO3) in the test tube with the liquid HCl, be sure that none of

the mixture is lost from the tube. You will notice a slow but steady stream of gas produced, as evidenced by

the flow of water into the beaker. When the rate of gas evolution slows considerably and no more CO2 is

evolved stop the experiment and close the pinch clamp.

Empty the water from the beaker into a 100 mL graduated cylinder and measure its volume which is equal to

the volume of CO2 produced (assuming that the density of water is 1 g/mL).

Record the barometric pressure. The vapor pressure of water at various temperatures is given in Table 2.1.

Calculate the gas-law constant, R, from your data, using the ideal-gas equation nRTPV as well as the

van der Waals equation nRTnbVV

naP

2

2

. 22 /59.3

2molatmLa CO

molLb CO /0427.02

OHatmCO PPP22

3

3

32

CaCO

CaCO

CaCOCOM

mnn

2COV Vwater displaced

Using the ideal gas equation:

Tn

VPRnRTPV

CO

COCO

2

22

Using the van der Waals equation:

Tn

bnV

V

naPRnRTnbV

V

naP

CO

COCO

CO

CO

CO

2

22

2

2

2 2

2

2

2

CHE 1402

Lab Manual 8

TABLE 2.1 Vapor Pressure of Water at

Various Temperatures

Temperature

(°C)

H2O vapor pressure

(mm Hg)

15 12.8

16 13.6

17 14.5

18 15.5

19 16.5

20 17.5

21 18.6

22 19.8

23 21.1

24 22.4

25 23.8

REVIEW QUESTIONS

Before beginning this experiment in the laboratory, you should be able to answer the following questions: 1. Under what conditions of temperature and pressure would you expect gases to obey the ideal-gas equation?

2. Calculate the value of R in L-atm/mol-K by assuming that an ideal gas occupies 22.4 L/mol at STP.

3. Why do you equalize the water levels in the bottle and the beaker?

4. Why does the vapor pressure of water contribute to the total pressure in the bottle?

5. What is the value of an error analysis?

6. Suggest reasons, on the molecular level, why real gases might deviate from the ideal gas law.

7. Newly devised automobile batteries are sealed. When lead storage batteries discharge, they produce hydrogen. Suppose the void volume in the battery is 100 mL at 1 atm of pressure and 25°C. What would be the pressure increase if 0.05 g H2 were produced by the discharge of the battery? Does this present a problem? Do you know why sealed lead storage batteries have not been used in the past?

8. Why is the corrective term to the volume subtracted and not added to the volume in the van der Waals equation?

9. A sample of pure gas at 20°C and 670 mm Hg occupied a volume of 562 cm3. How many moles of gas

does this represent? (HINT: Use the value of R that you found in question 2)

10. A certain compound containing only carbon and hydrogen was found to have a vapor density of 2.550 g/L at 100°C and 760 mm Hg. If the empirical formula of this compound is CH, what is the molecular formula of this compound?

11. Which gas would you expect to behave more like an ideal gas Ne or HBr? Why?

Figure 2.1

CHE 1402

Lab Manual 9

Experiment 2

Determination of R:

The Gas-Law Constant

Name(s)


REPORT SHEET

1. Mass of CaCO3 = ____________ g

2. Moles of CaCO3 = ____________ = Moles of CO2

3. Volume of 4M HCl = ____________ mL

4. Volume of water collected = ____________ mL = Volume of CO2(g) released

5. Barometric pressure = ____________ mbar = ____________ mm Hg

Show calculation overleaf

6. Temperature of water = ____________ oC = ____________ K

7. Vapor pressure of water = ____________ mmHg

8. Pressure of the gas (show calculations)

9. Gas-law constant, R, from ideal-gas law (show calculations)

10. R from the van der Waals equation (show calculations)

Accepted value of R _______ Justify your choice overleaf

CHE 1402

Lab Manual 10

EXPERIMENT 3

Colligative Properties:

Freezing-Point Depression & Molar Mass

OBJECTIVE Determination of the molar mass of a substance by using the colligative properties of a solution.



ring and ring stand unknown solid (2 g)

clamp naphthalene (50 g)

wire gauze 2-hole rubber stopper with slit

thermometer towel

large test tube wide-mouth glass bottle

wire stirrer weighing paper

Bunsen burner and hose

600-mL beaker

DISCUSSION

A solution is a homogeneous mixture of a solute in a given solvent. Both solute and solvent are components

of the solution. The solute is the smallest amount present in the solution while the solvent is the largest

amount present in the solution. Since the solution is primarily composed of solvent, physical properties of a

solution resemble those of the solvent. Some of these physical properties called colligative properties, are

independent of the nature of the solute and depend only upon the solute concentration. The colligative

properties include vapor-pressure lowering, boiling point elevation, freezing point lowering, and osmotic

pressure. The vapor pressure is just the escaping tendency of the solvent molecules. When the vapor pressure

of a solvent is equal to atmospheric pressure, the solvent boils. At this temperature the gaseous and liquid

states of the solvent are in dynamic equilibrium, and the rate of molecules going from the liquid to the

gaseous state is equal to the rate of molecules going from the gaseous state to the liquid state. It has been

found experimentally that the dissolution of a non volatile solute (one with very low vapor pressure) in a

solvent lowers the vapor pressure of the solvent, which in turn raises the boiling point and lowers the freezing

point. This is shown graphically by the phase diagram given in Figure 3.1.

Figure 3.1

CHE 1402

Lab Manual 11

You are probably familiar with some common uses of these effects: Antifreeze is used to lower the freezing

point and raise the boiling point of coolant (water) in an automobile radiator; and salt is used to melt ice.

These effects are expressed quantitatively by the colligative-property law, which states that the freezing point

and boiling point of a solution differ from those of the pure solvent by amounts that are directly proportional

to the molal concentration of the solute. This relationship is expressed by Equation [1] for the freezing-point

lowering and boiling-point elevation:

KmT [1]

where T is the freezing-point lowering or boiling-point elevation, K is a constant that is specific for each

solvent, and m is the molality of the solution (number of moles solute per/ 1000 g solvent). Some

representative constants, boiling points, and freezing points are given in Table 3.1. For naphthalene, the

solvent used in this experiment, the molal freezing-point depression constant (Kfp) has a value of

6.9218 °C/ m.

Table 3.1: Molal freezing point and boiling point constants

Solvent Freezing

point

(°C)

Kfp

(°C/m)

Boiling

point

(°C)

Kbp

(°C/m)

CH3COOH (acetic acid) 16.6 3.90 118.1 2.93

C6H6 (benzene) 5.4 5.12 80.2 2.53

CHCI3 (chloroform) -63.5 4.68 61.3 3.63

C2H5OH (ethyl alcohol) -114.1 --- 78.4 ---

H2O (water) 0.0 1.86 100.0 0.51

C10H8 (naphthalene) 80.6 6.9218 --- ---

C6H12 (cyclohexane) 6.6 20.4 80.7 2.79

Since the molal freezing-point-depression constant is known, it is possible to obtain the molar mass of a

solute by measuring the freezing point of a solution and the weight of both the solute and solvent.

In this experiment you will determine the molar mass of an unknown. You will do this by determining the

freezing-point depression of a naphthalene solution having a known concentration of your unknown. The

freezing temperature is difficult to ascertain by direct visual observation because of a phenomenon called

supercooling and also because solidification of solutions usually occurs over a broad temperature range.

Temperature-time graphs, called cooling curves, reveal freezing temperatures rather clearly. Therefore, you

will study the rate at which liquid naphthalene and its solutions cool and will construct cooling curves, for

both the pure solvent and the solution, similar to the ones shown in Figure 3.2.

CHE 1402

Lab Manual 12

Figure 3.2: Cooling curve.

Figure3.2 shows how the freezing point of a solution must be determined by extrapolation of the cooling

curve. Extrapolation is necessary because as the solution freezes, the solid that is formed is essentially pure

solvent and the remaining solution becomes more and more concentrated. Thus its freezing point lowers

continuously. Clearly, supercooling produces an ambiguity in the freezing point and should be minimized.

Stirring the solution helps to minimize supercooling.

PROCEDURE

A. Cooling Curve for Pure naphthalene Weigh a large test tube to the nearest 0.01 g. Add about 15g of naphthalene and weigh again. The difference

in weight is the weight of naphthalene. Assemble the apparatus as shown in Figure 3.3; be certain to use a

split two-hole rubber stopper. Carefully insert the thermometer into the hole that has been slit. Bend the stirrer

so that the loop encircles the thermometer. Fill your 600-mL beaker nearly full of water and heat it to about

85°C. Clamp the test tube in the water bath as shown in Figure 3. When most of the naphthalene has melted,

insert the stopper containing the thermometer and stirrer into the test tube; make certain that the thermometer

is not resting on the bottom of or touching the sides of the test tube. When all of the naphthalene has melted,

stop heating, remove the beaker of water, and dry the outside of the test tube with a cloth towel. Place the test

tube in a wide-mouth bottle that contains a piece of crumpled paper in the bottom to lessen the chance that

impact of the test tube with the bottle will cause the bottle to break. The purpose of the wide-mouth bottle is

to minimize drafts. Record temperature readings every 30 s while you are stirring. When the freezing point is

reached, crystals will start to form, and the temperature will remain constant. Shortly after this, the

naphthalene will solidify to the point where you can no longer stir it. Your lab instructor will direct you to

perform either procedure B or procedure C.

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Lab Manual 13

Figure 3.3

B. Determination of the Molar Mass of the unknown Using weighing paper, weigh to the nearest 0.01 g about 1.0 g of the unknown.

Replace the test tube in the water bath and heat until all the naphthalene has melted. Gently remove the

stopper, making sure that no naphthalene is lost, and add the unknown to the test tube. Replace the stopper

and stir gently until the entire unknown has dissolved. Remove the water bath, dry the test tube with a towel,

and insert the test tube in a wide-mouth glass bottle containing a crumpled piece of paper. Record the

temperature every 30 s until all the naphthalene has solidified.

C. Determination of the Molar Mass of the unknown Same as Part B, but use about 2.0 g of the unknown.

Cleanup: To clean out the test tube at the end of the experiment, heat the test tube in a water bath until the

naphthalene just melts. Care should be taken not to heat the thermometer beyond its temperature range. Be

careful, because naphthalene is flammable. Remove the stopper and pour the molten naphthalene on a

crumpled wad of paper. When the naphthalene has solidified, throw both the paper and solid naphthalene into

a waste receptacle. DO NOT POUR LIQUID NAPHTHALENE INTO THE SINK!

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Lab Manual 14

REVIEW QUESTIONS

Before beginning this experiment in the laboratory, you should be able to answer the following questions: 1. Distinguish between solute and solvent.

2. List three colligative properties and suggest a rationale for the choice of the word colligative to describe

these properties.

3. Distinguish between volatile and nonvolatile substances.

4. What effect does the presence of a nonvolatile solute have upon (a) the vapor pressure of a solution, (b) the

freezing point, and (c) the boiling point?

5. What is the molality of a solution that contains 1.5 g urea (molar mass = 60 amu) in 200 g of benzene?

6. What is supercooling? How can it be minimized?

7. Calculate the freezing point of a solution containing 6.50 g of benzene in 160 g of chloroform.

8. A solution containing 1.00 g of an unknown substance in 12.5 g of naphthalene was found to freeze at

75.4°C. What is the molar mass of the unknown substance?

9. How many grams of NaNO3 would you add to 250 g of H2O in order to prepare a solution that is 0.200

molal in NaNO3?

10. Define molality and molarity.

CHE 1402

Lab Manual 15

Experiment 3

Colligative Properties:

Freezing-Point Depression

& Molar Mass

Name(s)


REPORT SHEET

1. Weight of naphthalene _____________ g

2. Weight of unknown _____________ g

Cooling curve data:

Pure naphthalene Naphthalene + unknown

Time (s) Temp. (oC) Time (s) Temp. (

oC)

0 0

30 30

60 60

90 90

120 120

150 150

180 180

210 210

240 240

270 270

300 300

330 330

360 360

390 390

420 420

3. Freezing point of pure naphthalene, from cooling curve

4. Freezing point of solution of the unknown in naphthalene T = oC

5. Molality of the unknown (show calculations)

6. Molar mass of unknown (show calculations) _______

CHE 1402

Lab Manual 16

EXPERIMENT 4

Rates of Chemical Reactions

A Clock Reaction

OBJECTIVES To determine the order of a reaction with respect to the reactant concentrations.

To obtain the rate law for a chemical reaction.


APPARATUS & CHEMICALS

burets (2)

1-mL pipets (2)

clock or watch with second hand

125-mL Florence flask

400-mL beaker

test tube

pipet bulb

250-mL Erlenmeyer flasks (4)

buret clamp

ring stand

25-mL pipet

50-mL pipet

thermometer

0.2 M KI (200-mL)

0.4 M Na2S2O3 (100- mL) (freshly prepared)

1% percent starch solution, boiled

0.2 M KNO3 (300-mL)

0.1 M solution of Na2H2EDTA

0.2M (NH4)2S2O8, (200-mL)

(prepared from fresh solid)

DISCUSSION

Kinetics is the study of how fast chemical reactions occur. Among the important factors that affect the rates of

chemical reactions are:

1. Reactant concentration

2. Temperature

3. Catalyst

Before a reaction can occur, the reactants must come into direct contact via collisions of the reacting particles.

The reacting particles (ions or molecules) must have the right orientation and must collide with sufficient

energy to result in a reaction. With these considerations in mind, we can qualitatively explain how the various

factors influence the rates of reactions.

Concentration

Changing the concentration of a solution alters the number of particles per unit volume. The more particles

present in a given volume, the greater the probability of their collision. Therefore, increasing the

concentration of a solution increases the number of collisions per unit time and therefore the rate of reaction.

Temperature

Since temperature is a measure of the average kinetic energy, an increase in temperature increases the kinetic

energy of the particles. An increase in kinetic energy increases the velocity of the particles and therefore the

number of collisions between them in a given period of time. Thus, the rate of reaction increases. Also, an

increase in kinetic energy results in a greater proportion of the collisions having the required energy for

reaction.

Catalyst

Catalysts, in some cases, are believed to increase reaction rates by bringing particles into close juxtaposition

in the correct geometrical arrangement for reaction to occur. In other instances, catalysts offer an alternative

CHE 1402

Lab Manual 17

route to the reaction, one that requires less energetic collisions between reactant particles. If less energy is

required for a successful collision, a larger percentage of the collisions will have the requisite energy, and the

reaction will occur faster. Actually, the catalyst may take an active part in the reaction, but at the end of the

reaction, the catalyst can be recovered chemically unchanged.

Let's examine now precisely what is meant by the expression rate of reaction. Consider the hypothetical

reaction

A + B C + D [1]

Speed of a reaction is measured by the change in concentration of reactants or products with time.

Suppose A reacts with B to form C and D.

For the reaction A and B there are two ways of measuring rate:

1. the speed at which the products appear (i.e. change in concentration of C or D per unit time), or

2. the speed at which the reactants disappear (i.e. the change in concentration of A or B per unit time).

t

D

t

C

t

B

t

A

changethisforrequiredtime

ionconcentratinchangerateaverage

The units for average rate are mol/L-s or M/s.

In general, the rate of the reaction will depend upon the concentration of the reactants. Thus, the rate of our

hypothetical reaction may be expressed as

yxBAkrate [2]

where [A] and [B] are the molar concentrations of A and B, x and y are the powers to which the respective

concentrations must be raised to describe the rate, and k is the specific rate constant. One of the objectives of

chemical kinetics is to determine the rate law. Stated slightly differently, one goal of measuring the rate of the

reaction is to determine the numerical values of × and y. Suppose that we found x = 2 and y = 1 for this

reaction. Then

BAkrate2

[3]

would be the rate law. It should be evident from equation [3] that doubling the concentration of B (keeping

[A] the same) would cause the reaction rate to double. On the other hand, doubling the concentration of A

(keeping [B] the same) would cause the rate to increase by a factor of 4, because the rate of the reaction is

proportional to the square of the concentration of A. The powers to which the concentrations in the rate law

are raised are termed the order of the reaction. In this case, the reaction is said to be second order in A and

first order in B. The overall order of the reaction is the sum of the exponents, 2 + 1 = 3, or a third-order

reaction. It is possible to determine the order of the reaction by noting the effects of changing reagent

concentrations on the rate of the reaction.

It should be emphasized that k, the specific rate constant, has a definite value that is independent of the

concentration. k depends only on temperature. Once the rate is known, the value of k can be calculated.

Reaction of Peroxydisulfate Ion with Iodide Ion

In this experiment you will measure the rate of the reaction S2O8

2- + 2 I- I2 + 2 SO42-

[4]

and you will determine the rate law by measuring the amount of peroxydisulfate, S2O82-

, that reacts as a

function of time. The rate law to be determined is of the form:

CHE 1402

Lab Manual 18

YX

IOSkOSofncedisappearaofrate

2

82

2

82 [5]

or

YX

IOSkt

OS

2

82

2

82

Your goal will be to determine the values of x and y as well as the specific rate constant, k. You will add to

the solution a small amount of another reagent (sodium thiosulfate, Na2S2O3 which will cause a change in the

color of the solution. The amount is such that the color change will occur when 2 × 10-4

mol of S2O82-

has

reacted. For reasons to be explained shortly, the solution will turn blue-black when 2 × 10-4

mol of S2O82-

has

reacted. You will quickly add another portion of Na2S2O3 after the appearance of the color, and the blue-black

color will disappear. When the blue-black color reappears the second time, another 2 × 10-4

mol of S2O82-

has

reacted, making a total of 2× (2 × 10-4

) mole of S2O82

that has reacted. You will repeat this procedure several

times, keeping careful note of the time for the appearance of the blue-black colors. By graphing the amount of

S2O82-

consumed versus time, you will be able to determine the rate of the reaction. By changing the initial

concentrations of S2O82-

and I- and observing the effects upon the rate of the reaction, you will determine the

order of the reaction with respect to S2O82-

and I-.

The blue-black color that will appear in the reaction is due to the presence of a starch-iodine complex that is

formed from iodine, I2, and starch in the solution. The color therefore will not appear until a detectable

amount of I2 is formed according to Equation [4]. The thiosulfate ion S2O32-

that is added to the solution reacts

extremely rapidly with the iodine, as follows:

I2 + 2 S2O3

2- 2 I- + S4O62-

[6]

Consequently, until the same amount of S2O32-

that is added is all consumed, there will not be a sufficient

amount of I2 in the solution to yield the blue-black color (Figure 4.1). You will add 4 × 10-4

mol of S2O32-

each time (these equal portions are termed aliquots), and from the stoichiometry of equations [4] and [6] you

can verify that when this quantity of S2O32-

has reacted, 2 × 10-4

mol of S2O82-

has reacted. Note also that

although iodide, I-, is consumed according to equation [4], it is rapidly regenerated according to equation [6]

and therefore its concentration does not change during a given experiment (Figure 4.1).

S2O82- + 2 I- I2

(+ 2 SO42-)

starchBLUE-BLACK

COLORLESS

2 S2 O

3 2-very fast

S4O62- + 2 I-

[4]

[6]

Figure 4.1 Starch is used as a colored indicator allowing us to know when

reaction [6] is complete.

CHE 1402

Lab Manual 19

Graphical Determination of Rate

The more rapidly the 2 × 10-4

mol of S2O82-

is consumed, the faster is the reaction. To determine the rate of

the reaction, a plot of moles of S2O82-

that have reacted versus the time required for the reaction is made. The

best straight line passing through the origin is drawn, and the slope is determined. The slope, t

nOS

282 ,

corresponds to the moles of S2O82-

that have been consumed per second and is proportional to the rate. Since

the rate corresponds to the change in the concentration of S2O82-

per second, dividing the slope by the volume

of the solution yields the rate of disappearance of S2O82-

, that is,

t

OS

2

82 .

Helpful Comments

1. According to the procedure of this experiment, the solution will turn blue-black when exactly 2 ×10-4

mol

of S2O82-

has reacted.

2. The purpose of the KNO3 solution in this reaction is to keep the reaction medium the same in each run in

terms of the concentration of ions; it does not enter into the reaction in any way.

3. The reaction studied in this experiment is catalyzed by metal ions. The purpose of the drop of EDTA

solution is to minimize the effects of trace quantities of metal ion impurities that would cause spurious

effects on the reaction.

4. You will perform a few preliminary experiments to become acquainted with the observations in this

experiment so that you will know what to expect in the reactions.

5. The initial concentrations of the reactants have been provided for you on the report sheet.

PROCEDURE

Prepare four reaction solutions as follows (one at a time):

Solutions

0.2 M KI (mL) 25.0 25.0 50.0 12.5

0.2 M KNO3 (mL) 48.0 23.0 23.0 35.5

0.4 M Na2S2O3 (mL) 1.0 1.0 1.0 1.0

starch (mL) 1.0 1.0 1.0 1.0

EDTA 1 drop 1 drop 1 drop 1 drop

0.2 M (NH4)2S2O8

(mL)

added at time zero

25.0 mL

50.0 mL

25.0 mL

50.0 mL

Total Volume (mL) 100.0 100.0 100.0 100.0

Plastic dropper

buret

buret

buret

5-mL pipet

5-mL pipet

CHE 1402

Lab Manual 20

Each solution must be freshly prepared to begin the rate study - that is, prepare solutions 1, 2, 3, and 4 one

at a time as you make your measurements.

Equipment Setup

Set up three burets held by a clamp on a ring stand. Use these burets to measure accurately the volumes of the

KI, KNO3 and (NH4)2S2O8 solutions. Use two separate 5-mL pipets for measuring the volumes of the Na2S2O3

and starch solutions.

Rate Measurements

Procedure for solution 1:

Prepare solution 1 in a 250-mL Erlenmeyer flask that has been scrupulously cleaned and dried. Pour 25.0 mL

of (NH4)2S2O8 solution into a clean, dry 100-mL beaker. Be ready to begin timing the reaction when the

solutions are mixed (READ AHEAD). The reaction starts the moment the solutions are mixed!

BE PREPARED! ZERO TIME!

Quickly pour the 25.0 mL of (NH4)2S2O8, solution into solution 1 and swirl vigorously; note the time you

begin mixing to the nearest second. At the instant when the blue-black color appears, 2 × 10-4

mol of S2O82-

has reacted. IMMEDIATELY (be prepared!) add a 1.0 mL aliquot of Na2S2O3 solution from the pipet and

swirl the solution; the color will disappear. If the students fill each of seven clean, dry test tubes with 1.0 mL

of Na2S2O3 solution, they then can add these aliquots to their reactions at the appearance of the blue color

without loss of time.

Record the time for the reappearance of the blue-black color. Add another 1.0 mL aliquot of Na2S2O3 solution

and note the time for the reappearance of the color. The time interval being measured is that between the

appearance of the blue-black color. For good results, these aliquots of Na2S2O3 must be measured as quickly,

accurately, and reproducibly as possible. Continue this procedure until you have added seven aliquots to

solution 1.You are finished with solution 1 when you have recorded all your times on the report sheet. The

time intervals are cumulative.

Procedure for solution 3:

Solution 3 should be treated in exactly the same manner as solution 1. CAUTION: Be on guard-solution 3

will react much more rapidly than solution l.

Procedure for solutions 2 and 4:

Solutions 2 and 4 should be treated in exactly the same manner except that 50.0 mL portions of

(NH4)2S2O8 solutions should be added.

In each of these reactions the final total solution volume is 100 mL.

Calculations Tabulate on the data sheet for each aliquot of Na2S2O3 added to each of the four solutions:

1. The time interval from the start of the reaction (addition of S2O82-

) to the appearance of color for the first

aliquot of S2O82-

and the time interval from the preceding color appearance for each succeeding aliquot

(column2)

2. The cumulative time from the start of the reaction to each appearance of color (column 3)

3. For each solution plot on the graph paper provided the moles of S2O82-

consumed (as the ordinate, vertical

axis) versus time in seconds (as the abscissa, horizontal axis), using the data in columns 3 and 4. Draw the

tangent to the curve at time zero. Calculate the slope of each tangent, and from these calculations answer

the questions on your report sheet.

CHE 1402

Lab Manual 21

REVIEW QUESTIONS Before beginning this experiment in the laboratory, you should be able to answer the following questions: 1. What factors influence the rate of a chemical reaction?

2. What is the general form of a rate law?

3. What is the order of reaction with respect to A and B for a reaction that obeys the rate law

32BAkrate ?

4. Write the chemical equations involved in this experiment and show that the rate of disappearance of [S2O8

2-] is proportional to the rate of appearance of the blue-black color of the starch-iodine

complex.

5. It is found for the reaction A + B C that doubling the concentration of either A or B quadruples the rate of the reaction. Write the rate law for this reaction.

6. If 2 × 10

-4 mol of S2O8

2- in 50 mL of solution is consumed in 188 s, what is the rate of

consumption of S2O82-

?

7. Why are chemists concerned with the rates of chemical reactions? What possible practical value does this type of information have?

8. Suppose you were dissolving a metal such as zinc with hydrochloric acid. How would the particle size of the zinc affect the rate of its dissolution?

9. Assuming that a chemical reaction doubles in rate for each 100 temperature increase, by what factor would the rate increase if the temperature were increased by 40°C?

CHE 1402

Lab Manual 22

Experiment 4

Rates of Chemical Reactions

A Clock Reaction

Name(s)


REPORT SHEET

Solution 1. Initial [S2O82-

] = 0.05 M; initial [I-] = 0.05 M.

Time (s) between Total moles of S2O82-

Aliquot no. appearances of color Cumulative time (s) consumed

1 __________________ _____________ 2.0 × 10-4

2 __________________ _____________ 4.0 × 10-4

3 __________________ _____________ 6.0 × 10-4

4 __________________ _____________ 8.0 × 10-4

5 __________________ _____________ 10 × 10-4

6 __________________ _____________ 12 × 10-4

7 __________________ _____________ 14 × 10-4


] = 0.10 M; initial [I-] = 0.05 M.



1 __________________ _____________ 2.0 × 10-4

2 __________________ _____________ 4.0 × 10-4

3 __________________ _____________ 6.0 × 10-4

4 __________________ _____________ 8.0 × 10-4

5 __________________ _____________ 10 × 10-4

6 __________________ _____________ 12 × 10-4

7 __________________ _____________ 14 × 10-4

CHE 1402

Lab Manual 23


] = 0.05 M; initial [I-] = 0.10 M.



1 __________________ _____________ 2.0 × 10-4

2 __________________ _____________ 4.0 × 10-4

3 __________________ _____________ 6.0 × 10-4

4 __________________ _____________ 8.0 × 10-4

5 __________________ _____________ 10 × 10-4

6 __________________ _____________ 12 × 10-4

7 __________________ _____________ 14 × 10-4


] = 0.10 M; initial [I-] = 0.025 M.



1 __________________ _____________ 2.0 × 10-4

2 __________________ _____________ 4.0 × 10-4

3 __________________ _____________ 6.0 × 10-4

4 __________________ _____________ 8.0 × 10-4

5 __________________ _____________ 10 × 10-4

6 __________________ _____________ 12 × 10-4

7 __________________ _____________ 14 × 10-4

CALCULATIONS

1. Determine the slopes of the tangents and the corresponding rates of reactions.

Complete the table below.

Solution [S2O82-

]0 (M) [I-]0 (M) Tangent slope Rate

1 0.05 0.05 slope 1 = _________ mol/s r1 = ____________ M/s

2 0.10 0.05 slope 2 = _________ mol/s r2 = ____________ M/s

3 0.05 0.10 slope 3 = _________ mol/s r3 = ____________ M/s

4 0.10 0.025 slope 4 = _________ mol/s r4 = ____________ M/s

CHE 1402

Lab Manual 24

2. What effect does doubling the concentration of I- have on the rate of this reaction? (see table)

3. What effect does changing the [S2O82-

] have on the reaction? (see table)

4. Determine the partial orders, x and y, and the overall order of the reaction.

Show calculations

5. From your knowledge of the partial orders (as well as the rate in a given experiment), calculate

the specific rate constant, k, from your data.

Express the rate law of the reaction. ( YX

IOSkRate

2

82 ).

Show calculations

CHE 1402

Lab Manual 25

EXPERIMENT 5

An Equilibrium Constant

OBJECTIVES

To determine the equilibrium constant of a chemical reaction by spectrophotometric measurements.

To use graphing techniques and data analysis to evaluate data.


DISCUSSION

A spectrophotometric method of analysis involves the interaction of electromagnetic radiation EM with

matter. The most common regions of the EM spectrum used for analyses are the ultraviolet, visible, and the

infrared regions. We are most familiar with the visible region of the spectrum, having a wavelength range

from 400 to 800 nm. The colors of nature, the bluebonnet flowers, the red rocks, the green grass, and the

changing colors of the leaves in fall are a consequence of visible light interacting with the compounds that are

present in the material.

Every chemical substance possesses its own unique set of electronic, vibrational, and rotational energy states.

When EM radiation falls incident upon an atom or molecule, the radiation absorbed (the absorbed light) is an

energy equal to the difference between two energy states in the atom or molecule, placing the atom or

molecule in an "excited state". The remainder of the EM radiation passes through the sample (the transmitted

light) and an EM radiation detector detects it.

As the energy absorbed (and transmitted) equals the energy difference between the unique sets of energy

states in an atom or molecule, absorption and emission spectrophotometry methods are used to detect its

presence in a sample.

The energy absorbed, E, by an atom or molecule is related to the wavelength, λ, of the EM radiation.

hc

hE

h is Planck's constant ,and is the frequency of light

When a substance absorbs EM radiation from the visible region of the spectrum, it is usually an electron that

is excited from a lower to a higher energy state. When white light (EM radiation containing all wavelengths of

visible light) passes through the sample, our eyes detect the wavelengths of visible light not absorbed, i.e., the

light transmitted. The light that is absorbed excites electrons of the sample. Therefore, the color we see is

complementary to the one absorbed. If, for example, the atom or molecule absorbs energy from the violet

region of the visible spectrum, the transmitted light (and the substance) appears yellow (violet's

complementary color)-the higher the concentration of violet absorbing atoms or molecules, the more intense

is the yellow.

Thus the eye, one kind of EM radiation detector, detects only the transmitted light. Table 5.1 lists the colors

corresponding to wavelength regions of light (and their complements) in the visible region of the EM

spectrum.

CHE 1402

Lab Manual 26

Table 5.1: Color and Wavelengths in the Visible Region of the Electromagnetic Spectrum

Color Absorbed Wavelength (nm) Color Transmitted

red 750-610 green-blue

orange 610-595 blue-green

yellow 595-580 violet

green 580-500 red-violet

blue 500-435 orange-yellow

violet 435-380 yellow

In this experiment, EM radiation is used to determine the concentration of an absorbing substance in an

aqueous solution. The amount of transmitted light is measured using an instrument called a

spectrophotometer, an instrument that measures light intensities with a photosensitive detector at specific

(but variable) wavelengths. The wavelength at which maximum absorption of the EM radiation by the

absorbing substance occurs is determined and set on the spectrophotometer.

The ratio of the intensity of the transmitted light It to that of the incident light IO, is called the transmittance, T

(Figure 5.2).

This ratio, expressed as percent, is

TI

I t %%1000

[2]

Figure 5.2

Incident light, Io, and transmitted light, It, for a sample of concentration c in a cuvet of thickness l.

The spectrophotometer has a %T (percent transmittance of light) scale. Because it is linear, the %T scale is

easy to read and interpolate. Chemists often perform calculations based on the amount of light absorbed by

the sample, rather than the amount of light transmitted, because the extent of absorption is directly

proportional to the concentration of the absorbing substance. The absorbance, A, of the substance is related to

the intensity of the incident and transmitted light and the percent transmittance by the equation:

TTI

IA

t %

100log

1loglog 0 [3]

Several factors control the amount of EM radiation (light energy) that a sample absorbs:

- Concentration of the absorbing substance, c

CHE 1402

Lab Manual 27

- Thickness of the sample containing the absorbing substance, l (determined by the width of the cuvet)

- Probability of light absorption by the absorbing substance, ε (called the molar absorptivity coefficient

or extinction coefficient)

clA

This equation is commonly referred to as Beer's law.

A is the absorbance, ε (the molar absorptivity coefficient) constant at any given wavelength for a thickness l,

and c is the molar concentration of the absorbing substance.

The absorbance value is directly proportional to the molar concentration of the absorbing substance, if the

same (or a matched) cuvet and a set wavelength are used for all measurements. A plot of absorbance versus

concentration data is linear; a calculated slope and absorbance data can be used to determine the molar

concentration of the same absorbing species in an unknown solution.

Measuring an Equilibrium constant

The magnitude of an equilibrium constant, Kc, expresses the equilibrium position for a chemical system. For

the reaction,

aA + bB xX + yY

the mass action expression, ba

yx

BA

YX equals the equilibrium constant, Kc, when a dynamic equilibrium has

been established between reactants and products. The brackets in the mass action expression denote the molar

concentration of the respective substances. The magnitude of the equilibrium constant indicates the principal

species, products or reactants, that exist in the chemical system at equilibrium. For example, a large

equilibrium constant indicates that the equilibrium lies to the right with a high concentration of products and

correspondingly low concentration of reactants. The value of Kc is constant for a chemical system at a given

temperature.

This experiment determines Kc for a chemical system in which all species are soluble. The chemical system

involves the equilibrium between iron(III) ion Fe3+

, thiocyanate ion SCN-, and thiocyanatoiron(III) ion

FeNCS2+

:

Fe(H2O)6

3+ (aq) + SCN- (aq) Fe(H2O)5NCS2+ (aq) + H2O (l) [4]

Because the concentration of water is essentially constant in dilute aqueous solutions, we omit the water of

hydration and simplify the equation to read

Fe3+ (aq) + SCN- (aq) FeNCS2+ (aq) [5]

The mass action expression for the equilibrium system, equal to the equilibrium constant, is

SCNFe

FeNCSKc

2

2

In Part A you will prepare a set of standard solutions of the FeNCS2+

ion. As FeNCS2+

ion is a deep, blood-

red complex with an absorption maximum at 447 nm, its concentration is determined

spectrophotometrically. The absorbance for each solution is plotted versus the molar concentration of

CHE 1402

Lab Manual 28

FeNCS2+

; this establishes a calibration curve from which the concentrations of FeNCS2+

are determined for

the chemical systems in Part B.

In preparing the standard solutions, the Fe3+

concentration far exceeds the SCN-concentration. This huge

excess of Fe3+

pushes the equilibrium (Equation 5) far to the right, nearly consuming all of the SCN- placed in

the system. As a result the FeNCS2+

concentration at equilibrium approximates the original concentration. In

other words, we assume that the position of the equilibrium is driven so far to the right by the excess Fe3+

that

all of the SCN- is complexed, forming FeNCS

2+

In Part B, the concentrations of the Fe3+

and SCN- ions in the various test solutions are nearly the same, thus

creating equilibrium systems in which there is an appreciable amount of each of the species in the

equilibrium.

The chemical system is prepared by mixing known molar concentrations of Fe3+

and SCN-. By knowing the

initial concentrations of Fe3+

and SCN-, and by measuring the equilibrium concentration of FeNCS

2+

spectrophotometrically, the equilibrium concentrations of Fe3+

and SCN- are calculated. Using these

equilibrium concentrations, the Kc for the system is calculated.

At equilibrium, nFeNCS2+ = nFe3+ reacted = nSCN- reacted

and [Fe3+

]equilibrium = [Fe3+

] initial

_ [FeNCS

2+]equilibrium

[SCN-]equilibrium = [SCN

-]

initial

_ [SCN

-]equilibrium

PROCEDURE

One set of solutions having known molar concentrations of FeNCS2+

is prepared for a plot of absorbance

versus concentration. A second set of standard solutions is prepared to determine unknown molar

concentrations of FeNCS2+

. By carefully measuring the initial amounts of reactants placed in the reaction

systems, the mass action expression at equilibrium can be solved; this equals Kc.

A. A Set of Standard Solutions to Establish a Calibration Curve These solutions are used to determine the absorbance of known molar concentrations of FeNCS

2+. A plot of

the data, known as a calibration curve, is used to determine the equilibrium molar concentrations of FeNCS2+

,

in Part B.

Record on the Report Sheet the exact molar concentrations of the Fe(NO3)3 and NaSCN reagent solutions.

1. Prepare a Set of the Standard Solutions.

Pipet 0, 2, 5, 10, and 15 mL of 0.002 M NaSCN into separate, labeled, and clean 100-mL volumetric flasks

(Table 5.2). Pipet 25 mL of 0.2 M Fe(NO3)3 into each flask and dilute to the 100 mL "mark" with 0.25 M

HNO3. Stir each solution thoroughly to ensure that equilibrium is established.

CHE 1402

Lab Manual 29

Table 5.2: Compositions of Standard Solutions for Preparing the Calibration Curve.

Solution 0.002 M NaSCN

(in 0.25 M HNO3)

0.200 M Fe(NO3)3

(in 0.25 M HNO3)

0.25 M HNO3 [Fe3+

]0

(mol/L)

[SCN-]0

(mol/L)

[FeSCN2+

]equ

(mol/L)

1(blank) 0 mL 25 mL dilute to 100 mL 5ₓ10-2

0 0

2 2 mL 25 mL dilute to 100 mL 5ₓ10-2

4ₓ10-5

4ₓ10-5


10ₓ10-5

10ₓ10-5


20ₓ10-5

20ₓ10-5


30ₓ10-5

30ₓ10-5

2. Initial set-up

Turn on the spectrophotometer and let the light source (Tungsten halogen lamp) warm up for 10 minutes.

Then, set the wavelength at 447 nm by using the UP and DOWN arrow keys (Figure 5.3).

2. Prepare the Blank Solution.

Solution 1 is called the blank solution and will be used to calibrate the spectrophotometer. Rinse a cuvet with

several portions of Solution 1. Dry the outside of the cuvet with a clean Kimwipe or Joseph paper, removing

water and fingerprints. Handle the lip of the cuvet thereafter.

3. Calibrate the Spectrophotometer.

Open the sample compartment and insert the cuvet containing your blank solution.

Select the Absorbance mode by moving the cursor to the ABS mode using the LEFT or RIGHT arrow keys

(Figure 5.3). The primary display will show the absorbance, with ABS units. Press the CAL key to initiate

the calibration routine.

First the routine performs a zero% transmission calibration (by automatically activating an internal shutter -

this part of the routine is therefore independent of the solution in the light path) and the instrument will

display overrange absorbance “1. ABS” during this dark calibration. Then the instrument performs the 0.000

Absorbance calibration on your blank solution and will display “0.000 ABS”.

4. Record the Absorbance of the Standard Solutions.

Empty the cuvet and rinse it thoroughly with several portions of Solution 2. Fill it approximately three-

fourths full. Carefully dry the outside of the cuvet with a clean Kim wipe. Remember, handle only the lip of

the cuvet. Place the cuvet into the sample compartment; read the absorbance and record. Repeat with

Solutions 3, 4, and 5.

Disposal: Discard each test solution and each rinse into the appropriate disposal container.

5. Graph the Data. Plot, on linear graph paper, absorbance A (ordinate) versus [FeNCS2+

] (abscissa) for the

five solutions. Draw the best straight line through the five points and the origin. Ask your instructor to

approve your graph.

Record on the Report Sheet the exact molar concentrations of the Fe(NO3)3 and NaSCN reagent solutions.

CHE 1402

Lab Manual 30

Figure 5.3: Jenway 6320D Spectrophotometer employed in this experiment and its controls.

B. Absorbance of the test solutions

1. Prepare the Test Solutions.

In clean 150-mm test tubes (or 10-mL volumetric flasks), prepare the test solutions in Table 5.3. Use pipets

for the volumetric measurements. Be careful not to “mix” pipets to avoid contamination of the reagents prior

to the preparation. Also note that the molar concentration of Fe(NO3)3 for this set of solutions is 0.002 mol/L,

not the 0.2 mol/L solution used in Part A.

1. used to adjust values on the selected display

2. used to move the cursor horizontally between menu options

3. used to select the displayed menu option

4. initiates a calibration routine

5. Print key. Provides a printout of the current reading with an

incremental sample number. When pressed for the first time

after a calibration the print out will give calibration

information. The incremental sample number will be reset

after a calibration.

1. Primary display area - Transmission, Absorbance,

Concentration

2. Primary display adjust annunciator

3. Secondary display area - Wavelength, Factor

4. Primary display units

5. Secondary display adjust annunciator

6. Operation with PC

7. Menu options - %T ABS CONC FACTOR UNITS

8. Menu pointers (for 7)

CHE 1402

Lab Manual 31

Table 5.3 Compositions of Test Solutions for determination of Kc.

Solution 0.002 M NaSCN

(in 0.25 M HNO3)

0.002 M Fe(NO3)3

(in 0.25 M HNO3)

0.25 M HNO3 [Fe3+

]0

(mol/L)

[SCN-]0

(mol/L)

1’ 1 mL 4 mL dilute to 10 mL 8ₓ10-4

2ₓ10-4


4ₓ10-4


6ₓ10-4


8ₓ10-4

5’ 5 mL 5 mL - 10ₓ10-4

10ₓ10-4

2. Recalibrate the Spectrophotometer.

Use the blank solution (Solution 1) from Part A to check the 0.000 Absorbance.

3. Determine the Percent Transmittance of the Test Solutions. Stir each test solution until equilibrium is reached (approximately 1 minute). Rinse the cuvet thoroughly with

several portions of the test solution and fill it three-fourths full. Clean and dry the outside of the cuvet. Be

cautious in handling the cuvets. Record the absorbance of each test solution.

Disposal. Dispose of the waste thiocyanatoiron (III) ion solutions from the 100-mL volumetric flasks and the

cuvets in the appropriate waste container.

CLEANUP: Rinse the volumetric flasks, the pipets, and the cuvets twice with tap water and once with

deionized water. Discard each rinse in the sink.

4. Use Data to Determine Equilibrium Concentrations.

Using the calibration curve prepared in Part A.5 determine the equilibrium molar concentration of FeNCS2+

,

for each test solution.

5. Do the Calculations.

Complete the calculations as outlined on the Report Sheet. Complete an entire Kc calculation for Test

Solution 1’ (Part B) before attempting the calculations for the remaining solutions.

The equilibrium constant varies from solution to solution and from chemist to chemist in this experiment,

depending on chemical technique and the accumulation and interpretation of the data. Consequently, it is

beneficial to work with other colleagues through your own calculations, and then "pool" your final,

experimental K, values to determine an accumulated "probable" value.

CHE 1402

Lab Manual 32

REVIEW QUESTIONS Before beginning this experiment in the laboratory, you should be able to answer the following questions:

1. What effect does a dirty cuvet (caused by fingerprints, water spots, or lint) have on the percent

transmittance reading for a FeNCS2+

solution? Does this error cause the Kc to be reported as being too high

or too low? Explain.

2. In our calculations, the thickness of the solution (the cuvet) and the molar absorptivity of FeNCS2+

are not

considered. Explain.

3. If the percent transmittance reads less than 3.0%T on the spectrophotometer, how can the procedure be

modified to obtain a higher %T reading?

4. Over a period of time the 0 %T (no sample in the sample compartment) and the l00%T(blank solution in

the sample compartment) may drift from the initial calibration of the spectrophotometer. If the 100% T

reading drifts downward (less than 100%T), how does his error affect, in Part B, the :

a. absorbance readings?

b. [FeNCS2+

], equilibrium?

c. [Fe3+

], equilibrium?

d. [SCN], equilibrium?

CHE 1402

Lab Manual 33

Experiment 5

An Equilibrium Constant

Name(s)


REPORT SHEET Vsolution = ________ L

Solutions 1 2 3 4 5 Comment(s)

[Fe3+

]0 Ci ₓ Vi = [Fe3+

]0 ₓVsolution

[SCN-]0 Ci

’ ₓ Vi

‘= [SCN

-]0 ₓ Vsolution

Volume of Fe(NO3)3 (mL) 4.0 5.0 5.0 5.0 5.0 /

Moles of Fe3+

(at t0) = (L) ₓ [Fe3+

]0

Volume of NaSCN (mL) 1.0 2.0 3.0 4.0 5.0 /

Moles of SCN- (at t0) = (L) ₓ [SCN

-]0

Absorbance : spectrophotometer

[Fe(NCS2+

], equilibrium : calibration curve

Moles of Fe3+

, reacted = ₓ Vsolution (L)

Moles of Fe3+

, equilibrium = -

[ Fe3+

], equilibrium = ÷ Vsolution (L)

Moles of SCN- , reacted = = ₓ Vsolution (L)

11 Moles of SCN- , equilibrium 11 = -

12 [SCN-], equilibrium 12 = 11 ÷ Vsolution (L)

13 eqSCNeqFe

eqFeNCSKc

3

2

13 = ÷ ( ₓ 12)

CHE 1402

Lab Manual 34

EXPERIMENT 6

Titration Curves of Polyprotic Acids OBJECTIVE Determination of dissociation constants of a polyprotic acid.



pH meter with electrodes weighing bottle

potassium acid phthalate (KHP) buret clamp, and ring stand

sodium hydroxide solution about 0.1 M 25-mL pipet

100 mL of approx. 0.1M phosphoric acid 150-mL beaker (3)

standard buffer solution 250-mL beaker

phenolphthalein indicator solution

DISCUSSION Consider the triprotic acid H3PO4 .It undergoes the following dissociations in aqueous solution:

H3PO4 (aq) H2PO4- (aq) + H+ (aq)

43

421

POH

HPOHKa

[1]

H2PO4- (aq) HPO4

2- (aq) + H+ (aq)

42

2

42

POH

HHPOKa [2]

HPO42- (aq) PO4

3- (aq) + H+ (aq)

2

4

3

43

HPO

HPOKa [3]

The acid H3PO4 possesses three dissociable protons, and for this reason it is termed a triprotic acid. If you

were to perform a titration of H3PO4 with NaOH, the following reactions would occur:

H3PO4 + NaOH NaH2PO4 + H2O

[4]

NaH2PO4 + NaOH Na2HPO4 + H2O

[5]

Na2HPO4 + NaOH Na3PO4 + H2O

[6]

The resultant titration curve, when plotted as pH versus milliliters of NaOH added, would be similar to that

shown in Figure 6.1. At the point at which one-half of the protons in the first dissociation step of H3PO4 have

been titrated with NaOH, the H3PO4 concentration is equal to the H2PO4- concentration. Substituting [H3PO4]

= [H2PO4-] into Equation [1] yields Ka1 = [H

+], or pH = pKa1 at this point.

Similarly, at one-half the second equivalence point, one-half of the H2PO4- has been neutralized and [H2PO4

-]

= [HPO42-

]. Substituting this into Equation [2] yields Ka2 = [H+] or pKa2 = pH at this point.

In the same manner, at one-half the third equivalence point, [HPO42-

] = [PO43-

]. Substituting this into Equation

[3], we obtain the expression Ka3 = [H+], or pKa3 = pH.

CHE 1402

Lab Manual 35

The same type of result is obtained for any polyprotic acid. If a titration of the acid is performed with a pH

meter, the dissociation constants may be obtained from titration curves as long as the dissociation constants

exceed the ion product of water Kw, which you should recall is 10-14

for the reaction.

2 H2O H3O+ + OH- Kw =10-14

In practice, if the acidity of the acid being studied approaches that of water, as in the case for the third proton

of H3PO4 for which Ka3, is 4.2 × 10-13

it is difficult to determine the dissociation constant in this manner.

Thus for H3PO4, both Ka1 and Ka2 are readily obtained in this way, but Ka3 is not.

In this experiment you will determine the dissociation constants Ka1 and Ka2 of a phosphoric acid.

PROCEDURE

1. Standardize the pH meter.

2. Fill your buret with 0.1 M NaOH solution.

3. Titrate three separate 50-mL aliquots of the phosphoric acid in three separate 250-mL beakers and plot the

titration curves.

4. Determine the volumes necessary to reach the equivalence points.

5. Using the relations given above, and the relevant equations from this experiment, determine the Ka and pKa

and values for phosphoric acid.

HINT: You can save time if you do your first titration rapidly so that you know the approximate volumes of

the equivalence points; then you can do the next two titrations with large-volume increment away from the

equivalence points and small-volume increments near the equivalence points).

Figure 6.1: Titration curve of H3PO4.

CHE 1402

Lab Manual 36

REVIEW QUESTIONS Before beginning this experiment in the laboratory, you should be able to answer the following questions: 1. What is a polyprotic acid?

2. If 20.2 mL of 0.122 M NaOH is required to reach the first equivalence point of a solution of citric acid (H3C6H5O7), how many mL of NaOH are required to completely neutralize this solution?

3. How many mmol of NaOH will react with 50 mL of 6.2 M H2C2O4?

4. How many moles of H3O+ are present in 50 mL of a 0.3 M solution of H2SO4?

5. Why is it necessary to standardize a pH meter?

6. If the pH at one-half the first and second equivalence points of a dibasic acid is 3.52 and 6.31,

respectively, what are the values for pKa1 and pKa2? From pKa1 and pKa2 calculate the Ka1 and Ka2.

7. Derive the relationship between pH and pKa at one-half the equivalence point for the titration of a weak acid with a strong base.

8. Could Kb for a weak base be determined in the same way that Ka for the weak acid determined in this experiment?

9. If the Ka1 of a diprotic acid is 3.20, what is the pH of a 0.10 M solution of this acid?

CHE 1402

Lab Manual 37

Experiment 6

Titration Curves

of Polyprotic Acids

Name(s)


REPORT SHEET

Determination of pKa values of H3PO4

VNaOH

(mL)

pH VNaOH

(mL)

pH VNaOH

(mL)

pH

pKa1 __________ pKa2 __________

Ka1 __________ Ka2 __________

CHE 1402

Lab Manual 38

EXPERIMENT 7

Determination of the Solubility-Product

Constant for a Slightly Soluble Salt

OBJECTIVE

Determination of the value of the solubility-product constant for a slightly soluble salt.


APPARATUS AND CHEMICALS Buret 0.0024 M K2CrO4

Ring stand and buret clamp 0.004 M AgNO3

Centrifuge 0.25 M NaNO3

75-mm test tubes (3) 100-mL volumetric flasks (4)

Spectrophotometer and cuvets no.1 corks (3)

5-mL pipets (2)

DISCUSSION

Acids, bases, and salts are classified as inorganic substances. When an acid reacts with a base in

aqueous solution, the products are a salt and water, as illustrated by the reaction of H2SO4 and

Ba(OH)2:

H2SO4 (aq) + Ba(OH)2 (aq) BaSO4 (s) + 2 H2O (l) [1]

In general most common salts are strong electrolytes. The solubilities of salts span a broad spectrum,

ranging from slightly soluble to very soluble. This experiment is concerned with heterogeneous

equilibria of slightly soluble salts. In order for a true equilibrium to exist between a solid and

solution, the solution must be saturated.

For instance barium sulfate is a slightly soluble salt, and in a saturated solution this equilibrium may

be rewritten as:

BaSO4 (s) Ba2+ (aq) + SO42- (aq)

[2]

The equilibrium constant for Equation [2] is:

Ksp = [Ba2+

] [SO42-

] [3]

The value of Ksp is a constant at constant temperature.

The solubility product for a slightly soluble salt can easily be calculated by determining the

solubility of the substance in water. Suppose, for example, we determined that 2.42 × 10-4

g of

BaSO4 dissolves in 100 mL of water. The molar solubility of this solution (that is, the molarity of the

solution) is 1.04 × 10-5

M.

CHE 1402

Lab Manual 39

ML

molmolgg

mL

gs 5

44

1004.1100.0

/39.2331042.2

100

1042.2

Ksp = s × s = s2 = (1.04 × 10

-5)2 = 1.08 × 10

-10 at 25

oC

In a saturated solution the product of the molar concentrations of Ba2+

and SO42-

cannot exceed

1.08 × 10-10

at 25oC. If the ion product [Ba

2+][SO4

2-] exceeds 1.08 × 10

-10 at 25

oC, precipitation of

BaSO4 would occur until this product is reduced to the value of Ksp or if, for example, a solution of

Na2SO4 is added to a solution of Ba(NO3)2, BaSO4 would precipitate if the ion product

[Ba2+

]×[SO42-

] is greater than Ksp. If we determined that the solubility of Ag2CO3 were 3.49mg/100mL, we could calculate the

solubility-product constant for Ag2CO3 as follows:

The solubility equilibrium involved is

Ag2CO3 (s) 2 Ag+ (aq) + CO32- (aq)

[4]

and the corresponding solubility-product expression is Ksp = [Ag+]

2 [CO3

2-].

The solubility of Ag2CO3 in moles per liter is:

Mmol

g

L

gs 4

3

3

1027.16.27610100

1049.3

[CO32-

] = 1.27 × 10-4

M ; [Ag+]= 2 × 1.27×10

-4 = 2.54 × 10

-4 M

Ksp= [Ag+]

2 [CO3

2-] = [2.54 × 10

-4]

2 [1.27 × 10

-4] = 8.19 × 10

-12 M

3

To determine the solubility-product constant for a slightly soluble substance, we need only to

determine the concentration of one of the ions since the concentration of the other ion is related to

the first ion's concentration by a simple stoichiometric relationship.

In this experiment you will determine the solubility-product constant for Ag2CrO4. This substance

contains the yellow chromate ion, CrO42-

that should be determined experimentally by

spectrophotometric measurement at 375 nm.

To determine the solubility of Ag2CrO4, you will first prepare it by the reaction of AgNO3 with

K2CrO4:

2 AgNO3 (aq) + K2CrO4 (aq) Ag2CrO4 (s)15 min

+ 2 KNO3 (aq)

15 min

2 Ag+ (aq) + CrO42- (aq) Ksp = ?

If a solution of AgNO3 is added to a solution of K2CrO4, precipitation will occur if the ion product

[Ag+]

2 [CrO4

2-] numerically exceeds the value of Ksp if not, no precipitation will occur. Then, the

Ag2CrO4 formed is isolated by simple decantation and a new equilibrium is established between

solid Ag2CrO4 and Ag+ and CrO4

2- ions.

Ksp = [Ag+]equ

2 [CrO4

2-] equ.

CHE 1402

Lab Manual 40

PROCEDURE

A. Preparation of a Calibration Curve

Using a buret, add 1, 5, 10, and 15 mL of standardized 0.0024 M K2CrO4 to each of four clean, dry

100 mL volumetric flasks and dilute to the 100 mL mark with 0.25 M NaNO3.

1. Calculate the CrO42-

concentration in each of these solutions.

2. Measure the absorbance of these solutions at 375 mn.

3. Plot the absorbance versus concentration to construct your calibration curve.

B. Determination of the Solubility-Product Constant

Accurately prepare three separate solutions in separate 150-mm test tubes by adding 5 mL of

0.004 M AgNO3 to 5 mL of 0.0024 M K2CrO4. Stopper each test tube and shake the solutions

thoroughly at periodic intervals for about 15 min to establish equilibrium between the solid phase

and the ions in solution. Transfer the contents of each test tube into 75-mm test tubes and centrifuge.

Discard the supernatant liquid and retain the precipitate. To each of the test tubes add 2 mL of

0.25 M NaNO3. Shake each test tube thoroughly for another 15 minutes to establish equilibrium

between the solid and the solution and centrifuge again. There must be some solid Ag2CrO4

remaining in these test tubes. If there is not, start over again. Transfer the clear, pale yellow

supernatant liquid from each of the three test tubes to a clean, dry cuvette. Measure and record the

absorbance of the three solutions.

Using your calibration curve, calculate the molar concentration of CrO42-

in each solution.

Note on Calculations:

You should note that at equilibrium [Ag+]eq = 2[CrO42-

]eq. Therefore, having determined the

concentration of chromate ions, you know the silver-ion concentration

CHE 1402

Lab Manual 41

REVIEW QUESTIONS

Before beginning this experiment in the laboratory, you should be able to answer the following questions: 1 Write the solubility equilibrium and the solubility-product constant expression for the slightly

soluble salt CaF2.

2 Calculate the number of moles of Ag+ in 5 mL of 0.004 M AgNO3 and the number of moles of

CrO42-

in 5 mL of 0,0024 M K2CrO4,

3 If 10 mL of 0,004 M AgNO3 is added to 10 mL of 0,0024 M K2CrO4' is either Ag+ or CrO4

2- in

stoichiometric excess? If so, which is in excess? 4. The Ksp for BaCrO4 is 1.2 × 10

-10. Will BaCrO4 precipitate upon mixing 10 mL of 1 × 10

-4 M

Ba(NO3)2 with 10 mL of 1 × 10-4

M K2CrO4?

5. The Ksp for BaCO3 is 5.1 × 10-9

. How many grams of BaCO3 will dissolve in 100 mL of water?

6. Distinguish between the equilibrium-constant expression and Ksp for the dissolution of a sparingly soluble salt.

7. List as many experimental techniques as you can that may be used to determine Ksp for a sparingly soluble salt.

8. Why must some solid remain in contact with a solution of a sparingly soluble salt in order to

ensure equilibrium?

9. In general, when will a sparingly soluble salt precipitate from solution?

10. Sparingly soluble bases and salts, such as Fe(OH)2 and FeCO3 are more soluble in acidic than in neutral solutions. Why?

CHE 1402

Lab Manual 42

Experiment 7 Determination of the Solubility-Product

Constant for a Slightly Soluble Salt

Name(s)


REPORT SHEET

A. Calibration Curve

Initial [CrO42-

] = 0.0024 M

Volume of

0.0024 M K2CrO4

Total volume of

CrO42-

Final [CrO42-

]

(M)

Absorbance Molar absorption

coefficient

(M-1

cm-1

)

1 1 mL 100 mL _________ _______________ ____________

2 5 mL 100 mL _________ _______________ ____________

3 10 mL 100 mL _________ _______________ ____________

4 15 mL 100 mL _________ _______________ ____________

εaverage = ____________

B. Determination of Ksp

Absorbance [CrO42-

] [Ag+] Ksp of Ag2CrO4

1 ____________ ____________ ____________ ____________

2 ____________ ____________ ____________ ____________

3 ____________ ____________ ____________ ____________

Ksp average = ____________

CHE 1402

Lab Manual 43

EXPERIMENT 8

Molar Solubility, Common-ion Effect

OBJECTIVES

To determine the molar solubility and the solubility constant of calcium hydroxide.

To study the effect of a common ion on the molar solubility of calcium hydroxide.


DISCUSSION

Salts that have a very limited solubility in water are called slightly soluble (or "insoluble") salts. A

saturated solution of a slightly soluble salt is a result of a dynamic equilibrium between the salt and

its ions in solution; however, because the salt is only slightly soluble, the concentrations of the ions

are low. For example, in a saturated silver sulfate, Ag2SO4, solution, the dynamic equilibrium

between solid Ag2SO4, the Ag+ and SO4

2- ions in solution lies far to the left because of the low

solubility of silver sulfate.

Ag2SO4 (s) 2 Ag+ (aq) + SO4

2- (aq) [1]

for which the solubility product Ksp is: Ksp = [Ag+]

2 [SO4

2-]

What happens to the molar solubility of a salt when an ion, common to the salt, is added to the

saturated solution? According to Le Chatelier's principle, the equilibrium for the salt shifts to

compensate for the added ions; that is, it shifts left to favor the formation of more of the solid salt.

This effect, the addition of an ion common to an existing equilibrium, is called the

common-ion effect. As a result of the common ion addition and the corresponding shift in the

equilibrium, fewer moles of the salt dissolve in solution, lowering the molar solubility of the salt.

The molar solubility of a salt is the number of moles of that salt that dissolves per liter of (aqueous)

solution.

In this experiment, you will determine the molar solubility and the solubility constant for calcium

hydroxide, Ca(OH)2.

A saturated Ca(OH)2 solution is prepared; after an equilibrium is established between the solid

Ca(OH)2 and the Ca 2+

and OH- ions in solution, the decanted solution is analyzed. The hydroxide

ion, OH-, in the solution is titrated with a standardized HCl solution to determine its molar

concentration.

According to the following equation: Ca(OH)2 (s) Ca2+ (aq) + 2 OH- (aq)

[2]

For each mole of Ca(OH)2 that dissolves, 1 mole of Ca2+

and 2 moles of OH- are present in solution.

Therefore, by determining the molar concentration of hydroxide ion, [Ca2+

], Ksp, and the molar

solubility s of Ca(OH)2(s) can be calculated.

s

OHCa

2

2 ; 22 OHCaKsp

CHE 1402

44

Likewise, we can use the same procedure to determine the molar solubility of Ca(OH)2 in the

presence of added calcium ion, an ion common to the slightly soluble salt equilibrium.

PROCEDURE

The decantate from a saturated calcium hydroxide solution is titrated with a standardized

hydrochloric acid solution to the methyl orange endpoint. An analysis of the data results in the

determination of the molar solubility and solubility constant of calcium hydroxide. The procedure is

repeated on a decantate from a saturated calcium hydroxide solution containing added calcium ion.

A. Molar Solubility and Solubility Constant of Calcium Hydroxide

Three analyses are to be completed.

To hasten the analyses in Parts A.4 and A.5, prepare three, clean labeled 125- or 250-mL Erlemneyer

flasks and pipet 25 mL of the saturated Ca(OH)2 solution into each flask before you begin any

titration.

1. Prepare the Stock Calcium Hydroxide Solution. Prepare a saturated Ca(OH)2 solution 1 week

before the experiment by adding approximately 3 g of Ca(OH)2 to 120 mL of boiled, deionized

water in a 125-mL Erlenmeyer flask. Stir the solution and stopper. The resulting saturated solution

of calcium hydroxide is called limewater. This solution may have been prepared for you. Ask your

instructor.

2. Set Up the Titration Apparatus. Prepare a clean, 50-mL buret for titration. Rinse the clean buret

and tip with three 5-mL portions of the standard 0.05 M HCl solution and discard. Fill the buret with

standardized 0.05 M HCl, remove the air bubbles in the buret tip, and, after 30 seconds, read and

record the initial volume (± 0.02 mL). Record the actual concentration of the 0.05 M HCl on the

Report Sheet. Place a sheet of white paper beneath the receiving flask.

3. Transfer the Saturated Calcium Hydroxide Solution. Allow the solid Ca(OH)2 to remain

settled on the bottom of the flask (in Part A. 1). Carefully (try not to disturb the finely divided

Ca(OH)2 solid) decant about 90 mL of the saturated Ca(OH)2 solution into a second 125-mL flask.

4. Prepare a Sample for Analysis. Rinse a 25-mL pipet twice with 1- to 2-mL portions of the

saturated Ca(OH)2 solution and discard. Pipet 25 mL of the saturated Ca(OH)2 solution into a 125-

mL flask and add 2 drops of methyl orange indicator.

5. Titrate. Titrate the Ca(OH)2 solution to the methyl orange endpoint, where the color changes

from yellow (=basic pH) to red (=acidic pH). Remember the addition of HCl should stop within one-

half drop of the endpoint. Read (± 0.02 mL) and record the final volume of standard HCl in the

buret.

6. Repeat. Titrate two additional samples of the saturated Ca(OH)2 solution until a 0.50 %

(maximum) reproducibility is achieved.

100

1

100

2

x

xxn

xilityreproducib

i

CHE 1402

45

Example:

After 3 trials, you get the following volumes of equivalence:

V1= 10.15 mL, V2= 10.05 mL, V3= 10.15 mL.

3n

mLx 12.10

%47.0100

12.10

12.1015.1012.1005.1012.1015.103

1 222

ilityreproducib

7. Do the Calculations. The reported values for the Ksp of Ca(OH)2 will vary from chemist to

chemist. Complete your calculations as outlined on the Report Sheet.

B. Molar Solubility of Calcium Hydroxide in the Presence of a Common Ion

Again, three analyses are to be completed. Clean and label three 125- or 250-mL Erlenmeyer flasks.

Prepare all three of the Ca(OH)2-CaCI2 samples at the same time.

1. Prepare the Stock Solution. Mix 3 g of Ca(OH)2 and 1 g of CaCl2.2H2O with 120 mL of

boiled, deionized water in a 125-mL flask 1 weeks before the experiment. Stir and stopper the flask.

2. Prepare a buret for analysis, prepare the sample, and titrate. Repeat Parts A.2-A.6.

Disposal: Discard all of the reaction mixtures in the sink, followed by a generous supply of water.

CLEANUP: Discard the HCl solution in the buret into the sink. Rinse the buret twice with tap

water and twice with deionized water.

REVIEW QUESTIONS 1. How did the addition of CaCl2 affect the molar solubility of Ca(OH)2?

2. a. In Part A.3, suppose that some solid Ca(OH)2 was accidentally transferred to the titrating flask.

What effect does this error have on the reported Ks value?

b. As a result of the inadvertent transfer, will the calculated molar solubility for Ca(OH)2 be too

high or too low? Explain.

3. If the endpoint in the titration is surpassed in Part A.5, will the reported Ks value be too high or

too low? Explain.

4. Does adding boiled, deionized water to the titrating flask to wash the wall of the flask and the

buret tip affect the Ks value of the Ca(OH)2? Explain.

5. How will tap water instead of boiled, deionized water affect the Ks value of Ca(OH)2 in Part A?

Hint: How will the minerals in the water affect the solubility of Ca(OH)2 ?

CHE 1402

46

Experiment 8

Molar Solubility,

Common-ion Effect

Name(s)


REPORT SHEET

A. Molar Solubility and Solubility Constant of Calcium Hydroxide

Trial 1 Trial 2 Trial 3

1 Concentration of standard HCl solution (M) _______________

2 Buret reading, initial (mL) 0 0 0

3 Buret reading,final (mL) ________ ________ ________

4 Volume of HCl used (mL) ________ ________ ________

5 Amount of HCl added (mol) ________ ________ ________

6 Amount of OH- in satd. solution (mol) ________ ________ ________

7 Volume of satd. Ca(OH)2 solution (mL) 25.0 25.0 25.0

8 [OH-], equilibrium (M) ________ ________ ________

9 [Ca2+

] equilibrium (M) ________ ________ ________

10 Molar solubility of Ca(OH)2 (M) , s ________ ________ ________

11 Average molar solubility of Ca(OH)2 (M) _______________

12 Ksp of Ca(OH)2 ________ ________ ________

13 Average Ksp of Ca(OH)2 _______________

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B. Molar Solubility of Calcium Hydroxide in the Presence of a Common Ion

Trial 1 Trial 2 Trial 3

1 Concentration of standard HCl solution (M) _______________

2 Buret reading, initial (mL) 0 0 0

3 Buret reading,final (mL) ________ ________ ________

4 Volume of HCl used (mL) ________ ________ ________

5 Amount of HCl added (mol) ________ ________ ________

6 Amount of OH- in satd. solution (mol) ________ ________ ________

7 Volume of satd. Ca(OH)2/CaCl2 solution (mL) 25.0 25.0 25.0

8 [OH-], equilibrium (M) ________ ________ ________

9’ Molar solubility of Ca(OH)2 (M) , s’ ________ ________ ________

10’ Average molar solubility of Ca(OH)2/CaCl2 (M) _______________

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EXPERIMENT 9

Determination of Orthophosphate in Water

OBJECTIVE To gain some familiarity with the techniques of spectrophotometric analysis by analysing a water solution for

its phosphate content.

APPARUTUS & CHEMICALS water sample spectrophotometer

KH2PO4 (oven-dried) 1-L volumetric flask

conc. H2SO4 l-, 5-, and 10-mL pipets

ammonium vanadomolybdate solution 100-mL volumetric flasks (6)

cuvets graduated 5-mL pipet

balance CHCl3

DISCUSSION

Tripolyphosphates have been found to be extremely effective in enhancing the cleansing ability of detergents;

they also are very inexpensive. Their aid in cleaning is probably due, in part, to the stable complexes that they

form with Ca2+

and Mg2+

, thus softening the water. Their extensive use, however, has very serious side effects

on nature.

The accelerated eutrophication, or overfertilization, of our lakes has aroused a great deal of ecological

concern. Nutrient enrichment enhances the growth of algae and other microscopic organisms. This produces

the green scum of an algal bloom on the water surface, masses of waterweeds, and a depletion of dissolved

oxygen; it also kills fish and other aquatic organisms and produces malodorous water systems.

When the photosynthetically active algae population near a lake's surface rapidly expands, most of the oxygen

produced escapes to the atmosphere. After the algae die, they sink to the lake bottom, where they are

biochemically oxidized. This depletes the dissolved oxygen needed to support aquatic life. When oxygen is

removed, anaerobic decomposition of the algae continues, producing foul odors.

Although many factors affect algae growth, the only one that is readily subject to preventive control is the

supply of nutrients. The many nutrients important to the growth of algae include phosphorus, carbon,

nitrogen, sulfur, potassium, calcium, and magnesium. Many environmentalists have accepted the idea that

phosphorus is generally the key nutrient that limits the plant growth that a body of water can support.

There are at least four major sources of phosphorus associated with human activity: human and food wastes,

fertilizers, industrial wastes, and detergents. Although detergent products contribute only about one-third of

the phosphates entering our water systems, curtailing this particular source is a logical place to begin to

combat eutrophication.

The phosphate found in natural waters is present as orthophosphate, PO43-

, as well as the polyphosphates

P2O7-4

and P3O105-.

The species present, PO43-

, HPO42-

, H2PO4-, or H3PO4, depend on the pH. Trace amounts

are also present as organophosphorus compounds. Detergents usually contain triphosphate, P3O105-

which

slowly hydrolyses to produce orthophosphate, PO43-

, according to the following reaction:

P3O105- + 2 H2O 3 PO4

3- + 4 H+

In this experiment you will determine the amount of orthophosphate present in a sample of water of

unknown concentration.

Analytical Method

In dilute phosphate solutions ammonium metavanadate, (NH4VO3), molybdate (MoO42-

), and phosphate

(PO43-

) condense to form an intensely yellow colored compound called a heteropoly-acid, whose formula

CHE 1402

49

is thought to be (NH4)3PO4 NH4VO316MoO3. The intensity of the yellow color is directly proportional to

the concentration of phosphate. The relative amount of color developed is measured with a

spectrophotometer.

The amount of light absorbed by the spectrophotometer is directly proportional to the concentration of the

colored substance. This is stated by Beer’s law, also known as Beer-Lambert law:

clA

A is the absorbance, ε (the molar absorptivity coefficient, in M-1

cm-1

) constant at any given wavelength for a

thickness l (in cm), and c is the molar concentration of the absorbing substance.

The colored solutions that you study in this experiment have been found to obey the Beer-Lambert law in the

region of wavelengths ranging from 350 nm to 410 nm. It is convenient to run this experiment at

400 nm, wavelength at which maximum absorption is detected for you solute. The amount of phosphate in

the unknown sample of interest is determined by comparison with a calibration curve constructed by using a

distilled water reference solution and solutions of known phosphate concentrations. The minimum detectable

concentration of phosphate is about 0.01 mg/L (10 ppb). The usual experimental precision will lie within

about ±l percent of the result obtained by an experienced analyst.

Comparative Phosphate Levels in Water Systems

Limiting nutrients and their critical concentrations are likely to differ in different bodies of water. Analysis of

the waters of 17 Wisconsin lakes has led to the suggestion that an annual average concentration of 0.015 mg/L

of inorganic phosphorus (0.05 mg phosphate/L) is the critical level above which algal blooms can be expected

if other nutrients, such as nitrogen, are in sufficient supply. During the 1968-69 period, Lake Tahoe in Nevada

had an average phosphate level of 0.006 mg/L, while its tributaries averaged 0.08mg/L. Lake Tahoe is one of

the two purest lakes in the world. The other is Lake Baikal in Russia. This figure thus represents the lowest

natural-water value of phosphate one is likely to find.

In July of 1969 the phosphate level of Lahontan Reservoir (about 55 km east of Reno, Nevada) was 0.52

mg/L. By comparison, Lake Erie's phosphate level increased from 0.014mg/L in 1942 to 0.40 mg/L in 1967-

68. The U.S. Public Health Service has set 0.1mg/L of phosphorus (0.3 mg phosphate/L) as the maximum

value allowable for drinking water. Raw sewage contains an average of about 30 mg/L of orthophosphate, of

which about 25 percent is removed by most secondary sewage-treatment plants.

PROCEDURE

A. Preparation of Calibration Curve

Dissolve about 136 mg of oven-dried KH2PO4 (weigh accurately) in about 500 mL of water. Quantitatively

transfer this solution to a 1-L volumetric flask, add 0.5 mL of 98% H2SO4, and dilute to the mark with

distilled water. This yields a stock solution that is about 1×10-3

M in various phosphate species. From this

stock solution prepare a series of six solutions with phosphate concentrations 2×10-5

, 5×10-5

, 1×10-4

, 2×10

-4,

5×10-4

, and 7.5×10-4

M by appropriate dilution of the stock solution. You must know the precise

concentrations of these solutions.

Each point on the calibration curve is obtained by mixing 10 mL of the phosphate solution with 5 mL of the

ammonium vanadomolybdate solution (see note below) and measuring the absorbance on the

spectrophotometer at 400 nm. Your curve is constructed by plotting absorbance as the ordinate versus

concentrations of phosphate as the abscissa. A straight line passing through the origin should be obtained.

Your calibration curve should be handed in with your report sheet.

B. Analysis of Water Sample

The unknown samples (A, B, and C) to be analyzed are stored in three test tubes and tightly stoppered.

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50

Add 5 mL of the ammonium vanadomolybdate solution to 10 mL of the unknown and measure the

absorbances at 400 nm.

Note: Preparation of Ammonium vanadomolybdate solution:

Dissolve 40 g of ammonium molybdate (molybdic acid, 85% MoO3) in about 400 mL of distilled water.

Dissolve 1 g of ammonium metavanadate, NH4VO2 in about 300 mL of distilled water and add 200 mL of

concentrated nitric acid. Mix the two solutions and dilute to 1 L.

This solution is stable for about 90 days and will be provided for your analysis.

REVIEW QUESTIONS

Before beginning this experiment in the laboratory, you should be able to answer the following questions: 1. What volume of 1×10

-3 M solution is required to make 50 mL of solution with the following

concentrations: 2×10-5

, 5×10-5

, 1×10-4

, 2×10-4

, 5×10-4

and 7.5×10-4

M?

2. What species is thought to be the light-absorbing species in this experiment?

3. Write a balanced chemical equation for the formation of the light- absorbing species that results from

reaction of phosphate and ammonium vanadomolybdate (AVM).

4. State the Beer-Lambert law and define all terms in it.

5. What are the five fundamental components of a spectrophotometer?

6. Why is a calibration curve constructed? How?

7. How do you know whether to measure the absorbance of a more dilute or more concentrated solution if

the absorbance of your unknown solution is not within the limits of your calibration curve?

8. A 0.0750 M sample of CO(NO3)2 gave an absorbance of 0.38 at 505 nm in a l-cm cell. What is the cobalt concentration of a solution giving an absorbance of 0.52 in the same cell at the same wavelength?

9. A 17.28-ppm (1 ppm = 1 mg/L) a solution of FeSCN2+

has a transmittance of 0.59 when measured in a 1.00-cm cell at 580 nm. Calculate the extinction coefficient for FeSCN

2+ at this wavelength.

10. Define eutrophication.

11. If raw sewage contains 30 mg/L of phosphate and a secondary sewage treatment plant removes 25% of the phosphate, would a secondary treatment plant provide potable water if 0.3 mg/L is the maximum phosphate concentration allowable in drinking water?

12. Write a balanced chemical equation for the hydrolysis of triphosphate, P3O105-

, to orthophosphate, PO43-

.

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Experiment 9

Determination of Orthophosphate

in Water

Name(s)


REPORT SHEET

A. Preparation of Calibration Curve

[PO43-

] (M)

in the flask

[PO43-

] (M)

in the cuvette

Absorbance of PO43-

in the cuvette

Absorbance of PO43-

in the flask

Soln. 1 2×10-5

__________ __________ __________

Soln. 2 5×10-5

__________ __________ __________

Soln. 3 1×10-4

__________ __________ __________

Soln. 4 2×10-4

__________ __________ __________

Soln. 5 5×10-4

__________ __________ __________

Soln. 6 7.5×10-4

__________ __________ __________

B. Determination of the concentration of PO43-

in the unknown

Absorbance of

PO43-

in the cuvette

[PO43-

] (M) in the

cuvette

[PO43-

] (M) in the

sample test tube

unknown A __________ __________ __________

unknown B __________ __________ __________

unknown C __________ __________ __________

÷ 1.5 × 1.5

× 1.5

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EXPERIMENT 10

Galvanic Cells

OBJECTIVES To create a Galvanic Cell and measure the current it produces.

To observe several redox reactions and note the changes of substances from their elemental form to their

ionic form and vice-versa


BACKGROUND Oxidation-reduction reactions are reactions in which electrons are transferred from one element to

another. Many everyday reactions are based on oxidation-reduction (redox) reactions. Even rust is

produced by a redox reaction:

4 Fe + 3 O2 2 Fe2O3

(rust)

One common misconception is that water causes rust. Although objects in water will rust more quickly,

water only provides the medium for the ions to travel more quickly.

Copper ions and aluminium metal Redox reactions occur because one element has a stronger attraction for electrons than another element.

As an example, the copper 2+ ion will take electrons from aluminum metal, producing copper metal and

aluminum ions.

3 Cu2+ + 2 Al (s) 2 Al3+ + 3 Cu (s)

Due to the presence of a thin and dense layer of aluminium oxide, Al2O3, on the surface of any piece of

aluminium exposed to the oxygen of the air, Cu2+

ions cannot react directly with Al (s) as described in the

equation above.

The solution to this problem is to add Cl- ions to the solution by adding a few drops of NaCl. Chloride

ions react easily with aluminium to form AlCl3. In presence of Cl- ions, the reaction is:

3 CuCl2 (aq) + 2 Al (s) 2 AlCl3 (aq) + 3 Cu (s)

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53

Iodine and zinc metal Another redox reaction occurs between zinc and iodine in solution.

Zn (s) + I2 Zn2+ + 2 I-

(in alcohol)

purple colorless Chlorine will react with the iodide ion.

2 I- + Cl2 2 Cl- + I2

colorless purple

H+

Galvanic Cell (battery) A galvanic cell is an electrochemical cell that produces electrical energy from chemical reactions taking

place within the cell.

It consists of two compartments called half-cells. Each half-cell is made of a metal electrode immersed in

an aqueous solution of the same metal ion. The two metal electrodes are connected with wires to a

voltmeter, which ensures the flow of electrons. The two solutions are also connected through either a salt

bridge or a porous membrane which ensures the flow of ions.

The electrode at which oxidation occurs is called the anode, and the electrode at which reduction occurs is

called the cathode. The electrons flow spontaneously from the negative anode to the positive cathode.

Anions in solution move toward the anode and cations move toward the cathode, through the salt bridge.

Procedure 1. Make a Galvanic cell using two coins made of different metals. Cut several pieces of paper towel to

the size of the coins. Wet these pieces with 1M sodium chloride NaCl.

Place one coin on the bottom, then the pieces of wet paper towel and the other on top.

Measure the voltage produced by this Galvanic cell.

2. Place several pieces of zinc metal in a small beaker. Cover the metal with tincture of iodine solution

and set it aside.

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54

3. Place 25 ml of 1M copper sulfate (CuSO4) in a 50-ml beaker. Roll a 4 « by 4 » piece of aluminum foil

into a roll. Place the strip in the beaker. It does not have to be completely immersed.

4. Place 5 ml of 1M copper sulfate in a second 50-ml beaker and set it aside If the purple color of the

solution (prepared in step 2) over the zinc has faded, pour the liquid into a second beaker leaving the

metal behind. Add several drops of bleach to the solution .Add a few drops of 1-M acetic acid.

5. Remove the aluminum foil from the beaker of copper sulfate (step 3). The black solid is copper metal.

Observe the difference in the aluminum foil. Compare the color of the two beakers of copper sulfate.

6. Another piece of aluminum foil should be placed in other copper solution (step 4) so that the copper is

removed. Then the solutions can be flushed down the drain with plenty of water.

The aluminum foil may be thrown away in the trash.

The zinc solutions from steps 1 and 5 should be placed on a metal pan to dry.

The zinc oxide formed can then be buried in a landfill.

CHE 1402

55

Experiment 10

Galvanic Cells

Name(s)


REPORT SHEET

I. Galvanic cell voltage ________________V

1. Would other metal strips work?

2. Would other ionic compounds like KBr work in the salt bridge instead of NaCl?

II. Zinc, Iodine and chlorine

1. In what form is the zinc (Zn(s) or Zn2+

(aq)) when the purple color is gone from the solution?

2. In what form is the zinc when the purple color is gone from the solution?

3. Does bleach change the zinc ions back to the metal?

4. In what form is the chlorine (Cl2(aq) or Cl-(aq)) when the purple color appears in the solution?

5. What is the color of the solution after adding the bleach?

III. Copper and aluminium

1. What does the difference in color in the copper solution indicate?

2. What happened to the aluminium foil?

3. In what form is the copper at the end of the reaction?

4. Could copper be recycled by collecting it on aluminium foil?

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APPENDIX

laboratory manual for general chemistry...

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