AMBO UNIVERSITY FACULTY OF NATURAL AND COMPUTATIONAL SCIENCE DEPARTMENT OF CHEMISTRY BASIC PRACTICAL CHEMISTRY Chem.203 LABORATORY MANUAL

1

EXPERIMENTAL BASIC PRACTICAL CHEMISTRYCourse code: Chem.203

2

Experiment TitleChemical laboratory rules and policies Safety rules in the chemical laboratory

Page 4 5 7 14 24 32 38 44 49 55 59 64 68 72 76 81

1 2 3 4 5 6 7 8 9 10 11 12 13 14

Bunsen Burner Mass and volume measurement Oxidation- reduction reaction Equivalent weight of metal Diffusion of gases: molecular motion Molar volume of gases Colligative properties - freezing point depression Heat of reaction Acid- base titration Effect of concentration of reaction rates Reaction speed and temperature Chemical equilibrium and Le Chateliers principle- I Chemical equilibrium and Le Chateliers principle - II Determination of solubility of salt

CHEMICAL LABORATORY RULES & POLICIES 1. Students can only work in the lab section, which they have registered 3

2. Students must come to the lab on time. Anyone who comes to the lab 5 minutes late will not be

allowed to work in the lab.3. Always read the upcoming experiments carefully and thoroughly, being sure to understand all of the

directions before entering the lab.4. All students must come to the lab with their lab coats, lab notebooks and lab manuals. No one will be

allowed to work in the lab without a lab coat and a lab manual5. Students cannot leave the lab without permission. 6. Students must attend every experiment. Those who miss the lab without a valid excuse will get zero

for that experiment. Students who miss more than two labs will fail in the lab course for the semester.7. The name, date, purpose and equipment parts of the assignment as well as your experimental data

should be written to the notebooks during the laboratory session.8. It is forbidden for students to be involved in activities other than those related to that days

experiment. (No eating, drinking, smoking, doing assignments related to other courses etc.). No playback devices, cell phones and personal computers are allowed in the lab.9. For safety and health reasons, students must strictly follow lab rules and regulations and should work

carefully in the lab. Playing with chemicals are dangerous and forbidden.10.

Students will be asked to compensate for the items they have broken in the lab. Otherwise the student will not be allowed to take any exams No unauthorized experiments are to be performed in the laboratory Do not take the reagent bottles away from their places A laboratory apron or nonflammable coat should always be worn

11. 12. 13. 14.

Students are required to use only the items in their own working desk. Please inform your lab assistant about any missing and broken item so that it can be replaced. Taking from other working desk is forbidden. Students should clean the items they have used during the experiment and at the end of each lab session. Sandals, open toed and high heels are not permitted in the lab

15.

SAFETY RULES IN THE CHEMICAL LABORATORY There are many potential sources of hazard in a chemical laboratory .Therefore you should be aware of this fact and should at all times be on the alert to avoid or minimize accidents. Pay attention to the4

frequent warnings made by instructors regarding safe working procedures in the laboratory. You must realize that you are working in the laboratory at your own risk and you should therefore always strive to avoid and minimize accidents. There are certain ways of tacking laboratory accidents .you should be aware of this right from the beginning of the course A few instances are listed below:1. Smoking, eating or drinking are absolutely forbidden in the laboratory 2. Never taste any laboratory chemical. 3. Never inhale gaseous fumes. 4. Add acid to water, but never add water to acid. This is to prevent splashing from the acid due to the

generation of excessive heat as the two substances mix5. Never return unused chemicals to stock bottles. 6. Do not insert your pipette or dropper into the reagent bottles. 7. Always label your solution-containers. 8. Do not place a bottle stopper down on any surface. 9. Take the amount of reagent indicated. 10. 11. 12. 13. 14. 15. 16.

Never use flammable liquids near a flame. In case of fire, notify the instructor immediately Never throw matches, litmus paper, or any solid waste into the sink. Put any broken glass in container labeled "broken glassware". Chemicals should be assumed toxic unless known to be otherwise. Never pick up hot objects with your bare hands Report any accident immediately to your instructor Before using a burner be sure nobody else on the bench has any organic solvent.

17.

Be familiar with the following terms and their effects:5

Flammable------------------------- They are easily ignited (burn).

Irritants -------------------------------They irritate the eyes, lungs, and skin. Toxic ---------------------------------They are poisonous in either the short or the long term. Carcinogenic----------------------- -They cause cancer. Teratogenic----------------------------They cause defects in the unborn fetus. Mutagenic---------------------------- They cause genetic mutations. Explosive----------------------------- They explode, usually on being mixed with air. Corrosive---------------------------------- They burn the eyes, lungs, and skin

EXPERIMENT- 1 BUNSEN BURNER1. OBJECTIVE: 6

To learn handling the Bunsen burner and study properties of its flame 2. THEORY The principle of work of a Bunsen burner is the combustion of a gas composed of lower hydrocarbons. That may be natural gas (which consists mainly of methane, CH4) or butane, C4HIO, the one used in our laboratory (and sold locally in metal cylinders for domestic needs). The complete combustion of hydrocarbons produces carbon dioxide and water with a considerable release of heat energy. Butane, for example, it is completely burnt in sufficient supply of oxygen (in the ratio of butane to oxygen 2: 13 molecules) according to the equation 2C4H10 + 1302 8 CO2 + 10 H2O + heat The complete combustion of butane gas produces a blue non-luminous flame. If the supply of oxygen were limited, the hydrocarbon would not be completely combusted, one of the products being carbon. 2C4H10 + 802 3CO2 + 10H2O + 5C + heat In that case, the gas bums with a yellow luminous flame whose color is due to the hot particles of carbon. These move to the air and settle on surfaces of surrounding objects forming a layer of carbon (commonly known as soot). The Bunsen burner serves as a laboratory source of thermal energy. Many chemical processes are accelerated at high temperatures. Moreover, many processes are just impossible at room temperatures. That is why heating devices are necessary in chemical laboratories and one of these devices, widely used especially in teaching laboratories is the Bunsen burner. In this experiment, you are expected to learn the design of the Bunsen burner and to be able to operate it confidently and safely.

3. APPARATUS

Bunsen burner, tongs, evaporating dish, and lighter. 4. CHEMICAL Copper turnings (metallic Cu) 5. DESIGN OF THE BUNSEN BURNER7

Study the diagram of the Bunsen burner (Fig 1.1). The functions of its parts are as follows: The gas enters the burner through the gas inlet in the base and passes through the jet along the needle valve into the barrel. When lighted, it burns at the top of the barrel. Oxygen required is supplied from the surrounding air, which enters the barrel through the air inlet holes. These may be opened or closed by the collar. The needle valve is used to regulate the gas flow, and its working principle is shown on Fig. 1.2. The screw when screwed in prevents striking back of the flame, i.e. burning inside the barrel instead of above it. Now you must examine the "live" Bunsen burner, prepare all the additional equipment and may start the experiment 6. PROCEDURE 6.1. Operating the Bunsen burner6.1.1. Examine the burner. Find the parts denoted on Fig. 1.1. Unscrew the barrel and find the jet

and the needle valve. 6.1.2. Assemble the burner and make preparations for lighting itA. Make sure the bench gas tap is closed.

B. Open the gate valve (under your bench). C. Check the barrel is not screwed in fully D. Open the gas tap for a brief while (not more than 3 seconds ); the gas comes out with a hissing sound E. On hearing, this sound immediately closes the tap. The gas is poisonous and, besides, when in large quantities may explode. All the necessary measures must be taken to prevent escaping of unburnt gas to the air. F. Ask for a lighter

The most common infringement of these by students is to open the tap and start looking for a lighter. Do not repeat that mistake! You will not be allowed to light the burner, if you have managed to saturate the surrounding air with the gas. It is strictly forbidden to use paper to light the burner 6.1.3. Close the air inlet with the collar, open the gas tap and light the gas. 6.1.4. 6.1.5. Partially close and open the gas tap. Partially close and open the needle valve using the barrel.8

6.1.6. Open and close the air inlet using the collar. Note the effect of operations 6.1.4, 6.1.5 and 6.1.6 on the size and color of the flame6.1.7. Adjust controls so that the burner gives a quiet blue flame 10-15 cm in height. That is the type

of flame mainly used in laboratory experiments. 6.2. Properties of the flame 6.2.1. Examine the flame obtained in 6.1.7. Two different zones are seen. The inner one is known as a reducing flame, the outer zone, paler in color, is an oxidizing flame (see Fig 1.1). Investigate the two zones in the following way:6.2.2. Take a copper turning (with tongs) and hold it in the reducing flame for about 1min. making

sure it is really in the inner region. 6.2.3. Now hold it in the oxidizing flame

9

Note down your observations You could see that the flame is not homogeneous. Its two zones not only look different but are also of different temperatures. For future practical applications, you should know that the hottest part of the flame is its upper part, slightly below the tip. How investigate another interesting property of the oxidizing flame, the one that confirms its being named oxidizing.6.2.4. First, obtain a luminous flame. (What changes do you .expect to observe if a non-flammable

and high melting object, like evaporating dish, is held in that flame ?)6.2.5. With the help of the tongs, hold a porcelain-evaporating dish in a luminous flame.

Do not hold the dish in a luminous flame for a long time, but stop it as soon as the change is observed otherwise the next stage will take you a plenty of time.

Hold the evaporating dish firmly with the tongs; otherwise, it may easily slip and get broken. Note your observations. Was your hypothesis made on the previous step correct?6.2.6. Make the flame non-luminous and hold the affected part of the evaporating dish from 6.2.5

in the oxidizing zone. If you cannot obtain non-luminous flame entirely with the air inlet fully open, manipulating the screw may be helpful Do not place the hot evaporating dish onto the bench surface. Cool it in air for at least 2 min. Note down observations and explain them. 6.2.7. Turn off the gas tap and close the gate valve10

Learning goals On performing this experiment you must be able to light and operate the Bunsen burner and to explain the following term: air inlet, barrel, collar, combustion, gas inlet, jet, luminous flame, needle valve, nonluminous flame, oxidizing flame, reducing flame, screw

RABORATORY REPORT FORMAT EXPERIMENT - 1 BUNSEN BURNER Date: _____________ Group: ____________ Sub group: _________

Name: ______________________11

ID. ______________________ Partner(s): ____________________ Objective: ____________________________________________________________________________ ____________________________________________________________________________ Theory: In six lines describe combustion of butane; include chemical equations. Draw the diagram of the Bunsen burner label it Apparatus: ____________________________________________________________________ Chemicals: _____________________________________________________________________ Observations and conclusion Copy the tables below and fill in their right columns Operation The gas tap partially The needle valve partially The air inlet opened closed opened closed opened closed Observation (change of flame) color size

Operation Copper turning in a reducing flame Copper turning in an oxidizing flame Evaporating dish in a luminous flame Affected evaporating dish in an oxidizing flame

Observation

Conclusion

QUESTIONS 1. How many grams of butane can be oxidized to carbon dioxide and water by 13 g of oxygen? 2. How many grams of soot are produced in case of incomplete combustion of 2 g of butane? tap open?4. Which two controls are operated to regulate the height of the burner flame? 5. When the burner is functioning, the air is being steadily sucked through the inlets into the barrel.

3. Name two controls, which are to be checked if the gas does not come out of the barrel with the gas

Explain this phenomenon in a three lines' passage.6. Since the burner is generally used with air inlet open the question may arise if the collar is really 12

necessary. State your opinion on the subject and support it by at least two reasons (3 lines).7. Four students are each using a Bunsen burner to boil 100 mL of water in a beaker. All the four

burners are regulated to the same gas flow. The first two students put the beakers with water at the top part of the flame; the third one places it in the middle of the flame where the flame is wider; the fourth student puts it just above the top of the barrel. All the students except the first one work with a non-luminous flame. Which student will get his water boiled faster?8. Which of the students in the experiment described in Q. 7 will be the last to get his water boiled. 9. State three general disadvantages of producing and consuming energy by burning organic material as

compared to those of electrical energy10.

State three general advantages of producing and consuming energy by burning organic material as compared to those of electrical energy

Experiment 2 MASS AND VOLUME MEASUREMENTS1. OBJECTIVE:

To attain skill in the use of a three-beam balance, graduated cylinder, burette and pipette 2. THEORY Undoubtedly, you have come across measuring characteristics of objects such as mass, length, etc. As a ruler, you always try to make your measurements with the best possible means to obtain high accuracy. However, it has not probably occurred to you that it is impossible to measure the exact value of any quantity. An error is introduced in every measurement no matter how precise the instrument and how skilled the person is. This does not of course; mean that the measurement should be considered invalid, it13

only means, and that one should clearly indicate to what extent the measurement is certain. So, before starting this experiment study of experimental errors, accuracy, precision and significant digits 2.1. Experimental errors An error defined as deviation of a measured value (or average of measured values) from the true or accepted value (1) Where, E is error, xi value from an individual measurement, T true value, Errors classified into two groups; systematic errors and random errors. 2.1.1. Systematic errors A systematic error is the one that is constant in a set of readings taken for the same measurement. Systematic errors do not depend on how much the experimenter is skilled but are due to the inaccuracy of instruments or to the design of an experiment. They make all the readings taken either higher or lower than the true value. 2.1.2. Random errors A random error is the one, which varies in a set of reading taken for the same measurement. Random errors may be positive or negative. Their magnitude depends on the skill or care of an experimenter. To reduce the effect of random errors on a measurement several measurements of the same quantity are made and their arithmetic mean is calculated as mean of set measurements.

(2)

The larger the number of readings (n), the closer will errors. 2.2. Accuracy

be to the true value in the absence of systematic

Accuracy may be defined as the correctness of a measurement. When Xi approaches T (or T), the measurement becomes more accurate, i.e., E approaches O.

approaches

2.3. Precision (reproducibility) Precision maybe defined' as the closeness or agreement between the measured values of the same quantity . The mean deviation is one way of evaluating precision. A smaller mean deviation indicates higher precision of measurements (but is not guarantee to get accurate results. See Qn 8 and 9)2.4. significant digits ( significant figures ) 14

A significant digit is any numerical figure of a reported measurement that denotes an actual physical amount. Thus in 125 g there are three significant digits and in 125.2 g four, because every digit here indicates an actual physical amount. In 125.0 g, there are also four significant digits, for 0 indicates that there are no tenths of a gram In addition to 125 g in the mass reported. To report the measurement data with appropriate number of significant digits follow the rule: the last (right end) significant digit must indicate the estimated number of tenth parts of the smallest division of the scale (see Fig. 2.1). Significant digits are taken into account when calculations are made with experimental data. The result can never be of more significant digits than the least accurate number involved in a calculation. 3. DESCRIPTION OF APPARATUS 3.1. The three-beam balance The diagram (Fig. 2.1) shows the device, its principal parts labeled. Objects to be weighed are placed on the weighing pan and then balanced by moving weights (0.1 g, 1 g, or/and 10 g) along the beams until the pointer is at the centre (0) of the scale. The beams are graduated, the central one measuring tens of grams, the farthest one grams and the nearest tenths of a gram. The mass is thus the sum of readings from the three beams. The balance is also complete with adjustment devices (leveling screw, and leveling bubble) which are used to adjust the balance before weighing.

15

3.2.

Apparatus for measuring volumes of liquids These all require reading of a liquid level against a scale. To avoid a parallax error when reading have

your eye on a level with the surface of the liquid. The surface of a liquid in a container is not flat but curved and forms a meniscus. Liquids which wet glass such as water and aqueous solutions form a concave meniscus. For such liquids readings are taken at the bottom of the meniscus. Therefore, on Fig. 2.3 the volume is 7.0 mL which corresponds to the bottom of the meniscus, but not 7.1 mL corresponding to its top edge. For liquids which do not wet glass (mercury) the meniscus is convex and readings are taken from the top. 3.2.1. The measuring cylinder To measure out the, volume of the liquid required the measuring cylinder is filled with that liquid to the appropriate mark on the scale. Measuring cylinders are of different capacities tram 10 to 1000 mL. Small size cylinders are more accurate but even those do not provide high precision and are used for rough measurement only. More precise measurement is made by a burette and a pipette

3.2.2. The burette This is a graduated glass tube fitted with a stop-cork. For measurement it is fixed vertically in a clamps and filled with a liquid to be measured through the funnel placed at the top. Opening the stop-cork the liquid is allowed to flow out of the burette to a container. The volume the liquid flown out is found as the difference between the .final and initial readings. The burette is16

of

graduated to tenths of a milliliter and thus allows to measure volumes up to 0.01 mL.3.2.3.

The pipette This is also the accuracy of 0.01 mL but it can measure only the fixed volume which is indicated on

its bulb. Therefore, with 10 mL pipette we cannot measure 25 mL of a liquid but must take a 25 mL pipette. The volume indicated is measured out by filling the pipette to the mark, which is seen at some distance above the bulb of the pipette. 4. APPARATUS Three beam balance, measuring cylinder, burette, funnel, pipette, beakers, and two control objects for weighing. 5. CHEMICAL Tap water

6. PROCEDURE First train in handing measuring devices 6.1. Measuring mass a. Examine the three beam balance and find all the parts shown in Fig. 2. 1 b. Check the pan is empty, weights are at zero positions. c. Place the balance so that it will be convenient to operate. d. Level the balance by screwing its movable feet so that the leveling bubble in the indicator is at the centre of the circle. (Fig. 2.1). e. Release the beam by opening the beam lock and see if the pointer is at "0" of the scale. If not, screw the nuts to the right or left until the pointer comes to "0" Adjustments d and e must be done every time the balance is moved to a new place. f. Get from your instructor two labeled objects to be weighed and place one of them onto the17

pane g. Move the 10g weight to one notch and release the beam. If the mass of the object is more than10 g (pointer above "0"), arrest the beam and move the 10g weight to one more notch. If the mass is less than 10g (pointer below "0") arrest the beam, return the weight to zero position on the beam and move the 1 g weight. Manipulate the weights until the beam is balanced. There is no need to wait until the pointer settles at "0. When it swings an equal number of divisions from the two sides of rest point, the balancing is complete h. Calculate the mass of the object as the sum of all the weights. The mass is to be reported up to0.01 g (see 2.4). i. Record the label of the object and its mass. j. Repeat g - i for the second object Weighing chemicals never put them directly to the pan but use a container (beaker, dish, paper, etc.). The mass of the container is to be found first and then subtracted from the mass of the container with the chemical.

6.2. Measuring volume . For training purposes tap water is used Use of measuring cylinder a. To practice in how to use the measuring cylinder measure out volumes of 10,25,45, 50 and 100 mL of water Use of a buretteb. Check the burette is clean, the stop-cork moves freely and the holes in the stop-cork and the jet are

not blocked. If not, the stop-cork should be greased with a very small amount of grease. Blockages are cleared with wire. Rinse the burette with water. c. Clamp the burette vertically in the stand and fill it from the top with water. d. To fill the jet of the burette with water, open the stop-cork and allow some water to flow out. Make sure there are no air bubbles, close the stop-cork and wait for about 1 min. to see that there are no leakages. e. Read the liquid level on the scale.18

f. Open the stop-cork and, after some water flows out, close it. g. Read the level on the scale. The difference between the two readings is the volume delivered.h. Practice in using the burette measuring out successively 2, 3, 5, 7, 4.5 and 6.3mL

Use of a pipette i. Immerse the jet of the pipette in water in a beaker. Filling the pipette be sure its tip is always well dipped in the liquid to be sucked Otherwise air enters the pipette and the liquid will be sucked much easier and you will unexpectedly find some of it in your mouth. This may be a fun when water is measured, but with acids, bases and other chemicals this is seriousj. With your mouth, suck air out allowing the liquid to rise to a little above the mark.

k. Remove mouth and quickly put your index finger firmly over the top.l. Release the finger slightly to allow the liquid to fall to the mark and again firmly press down. m. Transfer the pipette to another beaker and release finger to allow the water to run out. Hold the tip

gently against the side of the container.n. Repeat i - m three or more times for better practice.

Every student must do operations 6.1 and 6.2 in turn

6.3.

Comparison of precision of a measuring cylinder, a burette and a pipette In this experiment, you will measure out the same volumes of water several times using a

measuring cylinder and then a more precise instrument (a burette or a pipette) and then determine masses of each volume of water. This will enable you to see which volume-measuring device gives a smaller mean deviation and therefore, which is more precise. a. Weigh a dry beaker and record its mass mb b. Measure 10 mL of water with the measuring cylinder, pour it to the beaker and weigh the beaker with 10 mL of water. Record the mass m1, c. Add another 10 mL of water to first 10 mL of it in the beaker and reweigh. Record the mass m2. d. Repeat c for the third and the fourth times and record mass m3 and m4.Mass of every 10 ml is calculated as the difference between two successive masses. e. Dry the beaker and repeat b - d this time measuring volumes with either a burette or a pipette (you do not need to use both) and recording mass ml', m2', m3' and m'4 In 6.3, every student performs at least one trial with a cylinder and one with a burette or a19

pipette. Learning goals This experiment helps yon attain the skill: Measurement of mass with a laboratory three beam-balance and measurement of volumes of liquids with a measuring cylinder, a burette and a pipette. You should also he able to understand the following terms. Burette, deviation, experimental error, mass, volume, measuring cylinder, meniscus, pipette, Precision, random error, significant digit, systematic error, three-beam balance

DO NOT WRITE ON THIS FORM! LABORATORY REPORT FORMAT EXPERIMENT -2 (Name of Institution)__________________________ Department________________________________ Student Department:_________________________ MASS AND VOLUME MEASUREMENTS

Date: ______________________ Section: ___________________ Group: ____________________ Sub group: _________________ Name:_____________________ ID.:____________________ Partner(s):__________________20

Objective:_____________________________________________________________________ ______________________________________________________________________ Theory: State definitions of a random error, systematic error, accuracy, precision, significant Apparatus: ______________________________________________________________________ Chemicals:______________________________________________________________________ Observations: Copy the tables below and fill in blank spaces. 1. . Weighing control objects Label of object digit

Mass

2. Comparison of precision of a measuring cylinder and a burette ( a pipette)

Mass of the beaker and the beaker with water m b= m 1= Measuring m 2= cylinder m 3= m 4= m 1'= m Burette or pipette 2'= m (specify ) 3'= m 4'= Conclusion

Mass of 10mL of waterm 1 m 2 m 3 m 4 m 1 m 2 m 3 m 4

Average mass of 10mL of water

Deviation

Mean deviation

-mb -m1 -m2 -m3

'- mb= '- m1'= '- m2'= '- m3'=

(Fill the blanks) The more precise is a ___________________________________Volumes measured in a cylinder have an uncertainty of _______________mL. Volumes measured in a __________ have an uncertainty of _______ mL. The volume of 10 mL if measured by a cylinder should only be quoted to_________________21

significant digits, and if measured by a ______________ should only be quoted to _____________ significant digits.

QUESTIONS1. If the pointer of a three beam balance swings five divisions above the rest point and two

divisions below is the object heavier of lighter than the weights on the balance.2. Using a three beam balance to weighs 10 g sodium chloride a student d:d the following

3.

4.

5. 6.

operations. List the mistakes he made Gust indicate the numbers). i. Leveled balance ii. Found rest point. iii. Moved balance nearer to NaCl bottle. iv. Moved weights on bean to 10g v. Put NaCl on pan gradually until rest point was reached. vi. Tipped NaCl into beaker How many significant digits are there in each of the following numbers? 20.8; 20.80; 2.08 x 103; 0.00208; 2.08 x 10-3 Two students wished to determine the density of an unknown liquid. Student A used 10 mL of liquid from a cylinder and found the mass to be 8.375 g. Student Bused 10.00 mL of liquid from a pipette and found the mass to be 8.81 g. Calculate the density of the liquid according to the two measurements. Give each result to the correct number of significant digits. Determine the number of significant digits in each of the following: 204; 3,000,000; 2.0120; 2361310 What is the number of significant digits in each of the following:22

500; 3.7 x 105; 1.873484 X 102; 0.0022746 7. Three students weighed the same object of mass 22.30 g on two different balances and made five trials each. The second and the third student used the same balance. The results were: 1st 2nd 3rd 22.41 g 21.97 g 22.55 g 22.39 g 21.98 g 22.39 g 22.38 g 21.97 g 21.80 g 22.39 g 22.00 g 21.85 g 22.43 g 21.97 g 21.75 g What type of experimental error (random or systematic) prevails in the measurements of the first, second and third student? 8. Calculate the mean deviation for measurements by each student in Qn 7. Which student was most precise? Least precise? 9. Calculate the mean error for measurements by each student in Qn 7. Which student was most accurate? Least accurate? Note from Qn 8 and 9 that the terms "accuracy" and "precision" are not synonymous. 10. You are given: m = 3.2; n = 2.1424; l= 0.085. In each of these, how many significant digits are there? Calculate mxn/ l.

Experiment 3 OXIDATION-REDUCTION REACTIONS1. OBJECTIVE:

To study some typical oxidation-reduction reactions and to determine the relative strength of oxidizing and reducing agents 2. THEORY Oxidation-reduction (redox) is a type of chemical reaction that involves transfer of electron(s) from one reactant to another. Oxidation is loss of electron(s) while reduction is gain of electron(s) by an atom, molecule or ion. The substance, which is, oxidized losses electrons to another substance involved in the reaction. Hence, the oxidized substance is a reducing agent and the reduced substance acts as an oxidizing agent. Oxidation and reduction processes.a1ways take place simultaneously. Redox reaction may be easily examined, if the individual processes are separately presented by half-equations: Oxidation: Red1 Ox1 + neReduction: Ox2 + ne- Red2 Net reaction: Red1 + Ox2 Ox1 + Red2 (3) (1) (2)

Where Red1 is the reduced form of Ox1 and Ox2 is the oxidized form of Red2 (or Red1) is a reducing agent23

and Ox2 is an oxidizing agent). The strength of oxidizing and reducing agents may be predicted from a series known as the ELECTROCHEMICAL series. In this series, the elements are arranged in a decreasing order of ease of oxidation (i.e. in an increasing order of being reduced) from left to right. The series represented below contains elements selected for our purpose and is not complete. Li K Ca Mg Al Zn Fe Ni Pb H Cu Ag I Br CI F

The series is useful to predict the possibility of redox reactions based on the following facts. An element at the left is a stronger reducing agent than any at the right and vice versa. ii. A metal (including hydrogen) at the left reduces the oxidized form of the element at the right in the series. For example 2AlO + 6H+Cl- 2AI3+Cl-3 + 3H20 (4)

MgO + Cu2+SO42-

Mg2+S042- + Cuo

(5)

iii. A halogen in its reduced form (halide) at the left can be oxidized by the stronger oxidizing

elemental the right in the series. For example 2I- + F2O 2F- + I2O (6)

Because of the above reason, the electrochemical series is sometimes known as REPLACEMENT series. Oxidation-reduction reactions also occur between atoms, molecules, and ions other than the ones in the replacement series above. Depending on their strength of reducing or oxidizing properties, they may be included in the electrochemical series. The following equations serve as examples 2K+Mn7+O4 2- + 16HCl 2Mn2+Cl22-+ 2K+Cl- + 5Cl2O + 8 H2O (7) (8)

Cr26+O72- + 6Fe2+ + 14H+ 2Cr3+ + 6Fe3+ + 7H2O

Examination of the oxidation number changes tells us what is reduced and what is oxidized in each redox reaction.3. APPARATUS

Test tubes, spatula, and lighter24

4. MATERIAL Wooden splinter 5. Chemicals Solutions of copper(II) nitrate, iron(II)nitrate, zinc nitrate, magnesium nitrate, potassium iodide, sodium bromide, sodium chloride; dilute hydrochloric acid, Chlorine water, bromine water, iodine, carbon tetrachloride; zinc strip, iron filings, copper wire or strip and magnesium ribbon

6. PROCEDURE A Use reagents economically, DO NOT take anything more than the prescribed amount. Save the metal strips if there is no (or no more) reaction with the solution in the test tube. Wash them with tap water and return to the lab. Technician after drying 6.1. For each step below add 5 mL of copper (II) nitrate solution into a test tube anda. Add place a strip of clean zinc metal. b. Add one spatula of iron filings. c. Add a strip of clean magnesium ribbon.

6.2. Add 5 mL of iron (II) nitrate solution into a test tube and place a. A piece of copper wire or strip b. A strip of clean zinc metal c. A small magnesium ribbon 6.3. Pour 5 mL of zinc nitrate solution into a test-tube anda. Place a piece of copper wire or strip. b. Add a spatula of iron filings. c. Place a clean magnesium ribbon. 6.4. Add 5 mL of magnesium nitrate solution into a test tube and 25

a. b. c.

Place a piece of copper wire or strip. Place a strip of clean zinc metal. Add spatula of iron fillings.6.5. For each step below, add 5 mL of dilute hydrochloric acid into a test tube and place each metal into

it. Test the gas produced, if any, by bringing a burning wooden splinter into the mouth of the test tube. Do you get a popping sound? If yes, what is it? Shake the test tube very gently and direct its mouth away from you and your colleagues.a. A piece of copper wire or strip. b. Spatula of iron filings (warm, do not boil), [Read caution). c. A strip of clean zinc metal

d. A clean magnesium ribbon

The following procedure may involve the formation of a halogen. The halogen produced, if any is soluble in carbon tetrachloride why? 7. PROCEDURE B 7.1. Mix the following in a test-tube and add 1 mL of carbon tetrachloride and shake wella. 3 mL of potassium iodide (KI) solution and an equal volume of chlorine water (add more

drops, if necessary, of chlorine water until a violet product appears in the CCl4 layer).b. 3 mL of sodium chloride (NaCl) solution and 2 mL of iodine water c. 3 mL of sodium bromide (NaBr) and equal volume of chlorine water. If red solution appears

in the CCl4 layer, add more drops of chlorine water.d. 3 mL of sodium chloride and an equal volume of bromine water e. 3 mL of potassium iodide and an equal volume of bromine water (apply the same procedure

as in step c if red color appears in the CCl4 layer).f. 3 mL of sodium bromide and 2 mL of iodine in water

Learning goals

From this experiment and your previous background, you are expected to have grasped the meaning of the26

following: Bromine water, chlorine water, electrochemical series, oxidation, oxidation half equation, oxidationreduction (redox) reaction, oxidizing agent, oxidized and reduced forms of a substance (with examples), reducing agent, reduction, reduction half-equation

LABORATORY REPORT FORMAT EXPERIMENT- 3 OXIDATION-REDUCTION REACTIONS

Date: ________________ Section: _____________ Sub group: _____________________ Name: ________________________________________ ID.:____________________ Partners: __________________________________ Objective: ________________________________________________________________ ________________________________________________________________ Theory: Explain redox reactions and the terms associated with it in less than eight lines. _______________________________________________________________ Apparatus: _________________________________________________________________ Materials: __________________________________________________________________27

Chemical: __________________________________________________________________ Observations Fill in the results and the inferences in the table below 1. PROCEDURE A a. Addition of metals indicated below into copper (ii) solution (operation 6.1). Metal Zinc Iron filings Magnesium ribbon Observation Inference

b. Addition of metals indicated below into iron (ii) solution (operation 6.2). Metal Zinc Copper Magnesium ribbon Observation Inference

c. Addition of metals indicated below into zinc (ii) solution (operation 6.3). Metal Copper Iron filings Magnesium ribbon Observation Inference

d. Addition of metals indicated below into magnesium nitrate solution (operation 6.4). Metal Zinc28

Observation

Inference

Iron filings Copper

e. Placing metals below into dilute hydrochloric acid (operation 6.5) Metal Zinc Iron filings Magnesium Copper Observation Inference

2. EQUATIONS Write net ionic equations for each of the above, if any. Example: When Zn metal is placed into copper (II) solution ZnO + Cu2+ 3. PROCEDURE B Prepare a similar table for procedure B and report the results Zn2+ + CuO

29

QUESTIONS

1. Which halide is the strongest reducing agent? 2. Which halogen is the strongest reducing agent? 3. Which halogen is the strongest oxidizing agent? 4. Balance each equation below and state the oxidation state change(s) a. Fe + HN03 Fe (NO3)3 + NO + H2O b. Mn2+ + BiO3 + H+ MnO4- +/Bi3+ + H2O c. ZnS + N03 - + H+ Zn2+ + S + NH4+ + H2O d. MnO4 - + H+ + C Mn2+ + CO2 + H2O 5. Indicate the oxidizing and reducing agents in each of the following: a. FeS2 + N2O2 Fe2O3 + N2SO4 + N2O b. As2S5 + Cl2 + H2O H3AsO4 + H2SO4 + HCl c. CN- + MnO4- + H2O CNO- + MnO2 + OH6. Balance the equations in question 5. 7. For the underlined elements in each set, which one is in a more oxidized form? a. MnO4-, MnO42d. AsH3, H3AsO4 b. CN-, CNOa. H2O, H2O2 c. PH3 , P4

8. . Which one of the two elements (underlined) in each equation is at the left side of the electrochemical series? a. MgO + Cu2+ Mg2+ + CuO b. 2Ag+ + ZnO 2Ag + Zn2+ c. 3Br- + Cl2 Br2 + 2Cl30

9. What change, if any; would you expect in the following? Write equation where necessary a. Bubbling fluorine gas into iron (II) chloride solution b. Adding iron fillings to magnesium sulphate solution c. Mixing bromine water with sodium chloride or fluoride solution

EXPERIMENT - 4 EQUIVALENT WEIGHT OF METAL - I1. OBJECTIVE

To study- a method for determining the equivalent weight for a metal. 2. THEORY One of the laws of chemical change is that exactly equivalent weights of elements combine to form compounds. The equivalent weight of an element may be defined as the weight in grams of the element, which combines with or displaces 8 g of oxygen or 1.008 g of hydrogen or its equivalent: The equivalent weight of zinc for example, when it combines with oxygen can be calculated using the equation below. Zn 65 g 32.5 g + 1/2O2 16 g 8g ZnO (1)

In any reaction, the number of equivalents of the reactants and products are equal. Hence, one equivalent of oxygen reacts with one equivalent of zinc to form one equivalent of zinc oxide. X equivalent of any substance reacts with X equivalents of another substance. The following equation will serve additiona1example. Mg 2 equivalents + 2H + 2 equivalents Mg2+ + H2 2equivalent (2)

2 equivalents

For reaction of the above type (redox), the number of moles of electrons gained or lost per mole of the substance reduced or oxidized is the number of equivalents of the substances. The equation below illustrates this 2AlO + 3Cu2+ 6 equivalents 2Al3+ 6equivalent + CuO 6 equivalent31

6equivalents

From the equation, the equivalent weight of aluminum is 2 x at. wt. /6 or more simple at. Wt/3 and that of copper is 3 x at. wt /6 or at Wt/2 From a given weight of a reactant and the quantity of a product, the equivalent weight of the reactant can be calculated by direct proportion, if the equivalent weight of the product is known or vice-versa.

3. APPARATUS

150 mL beaker, glass stirrer, funnel, wash bottle, balance, and watch glass. 4. MATERIAL Filter paper 5. CHEMICALS 0.25 g of zinc sheet or-granulated zinc, 0.3 M copper (II) sulphate solution, 0.1 M barium chloride solution, and alcohol. 6. PROCEDURE6.1. Drop about 0.25 gram of zinc metal into 25 mL of a warm solution of 0.3 M copper sulphate.

Occasionally stir and heat the, mixture to dissolve the zinc metal and precipitate copper powder.6.2. Filter the solution through a weighed filter paper (the diagram (Fig 4.1) shows how you fold the

filter paper and place it on a funnel)..6.3. Wash the copper powder, with water from a wash bottle (the instructor will show you how to l1se a

wash bottle), to free it from copper sulphate solution. Complete removal of copper sulphate solution is known by testing the filtrate (drops from the funnel) with barium chloride solution, which gives a white precipitate if there is copper sulphate on the metal. Finally pour 5 mL of alcohol over the filter paper. Why?6.4. Remove the filter paper carefully from the funnel, unfold it and spread it on a watch glass. Dry the

filter paper in sun light until constant mass is obtained.

32

Fig.4.1 Preparation of a fluted filter paper

Learning goals

From this experiment and your previous background, you are expected to understand the concepts and the use of the apparatus and skills below: Filtration, filter paper, equivalent weight, equivalent weight of an oxidizing agent, equivalent weight of a reducing agent, wash bottle.

33

LABORATORY REPORT FORMAT EXPERIMENT -4 EQUIVALENT WEIGHT OF METAL Date: _______________________ Group: _____________________ Sub group: ____________________ Name: __________________________________________ ID.:____________________ Partner(s):________________________________________ Objective: _____________________________________________________________________________ _____________________________________________________________________________ Theory: Define equivalent weight Write equation of the reaction of Zn with CU2+ ions and equivalent weight Apparatus: ____________________________________________________________________ Material:_______________________________________________________________________ Chemicals: ____________________________________________________________________ Observations: Record: i. if there is a color change or fading of the copper sulphate solution; ii. the color of the metal precipitated; iii. observations of the barium chloride test34

deduce the

DATA Mass of zinc =______________grams Mass of copper and filter paper =_____________grams Mass of copper =____________grams Equation of reaction _____________ + ______________ ___________ + _______________ (Calculate the equivalent weight of zinc from the data and the equivalent weight of copper, 31.75) ____________________________. The equivalent weight of zinc is___________

Error calculations The correct equivalent weight of zinc is 32.7. Calculate the percentage error from your determined value. Absolute error = correct value - experimental value Percentage error = Correct value experimental value X 100 Correct value Percentage error = ______________%

Possible source of error: ________________________________________________________________________________ ________________________________________________________________________________

35

QUESTIONS

1. What would be the error if the copper sulphate is not completely washed out? Is it Negative error or positive error? 2. If one has to prepare copper (II) sulphate solution in water for this experiment, the salt must be ground to powder. Why? 3. What is the relationship between the equivalent weight and the atomic weight of zinc? For any element that is oxidized or reduced. 4. Why was alcohol poured over the filter paper? 5. Why is the reaction mixture in the experiment stirred and heated? 6. 5.4 g of silver combines with 4.0 g of bromine, and 0.4 g of hydrogen unites with 32 g of bromine. What is the equivalent weight of silver (At. wt. of H = 1.0)? 7. 1.0 g copper is displaced from CuSO4 solution by 0.38 g of an unknown metal. Compute the equivalent weight of the metal. The equivalent weight of copper is 31.75. 8. 0.60g of a metal whose equivalent weight is 9 displaced 7.2g of silver from silver nitrate solution. Find the equivalent weight silver from the data only. 9. Why does the color of the copper sulphate solution change in the experiment? 10. Write the equation of the reaction of barium chloride test.

36

EXPERIMENT - 5

DIFFUSION OF GASES: MOLECULAR MOTION

1. OBJECTIVE To verify Graham's la w by measuring the relative diffusion rates of two gases; and to observe the motion of gas molecules 2. THEORY On the account of the molecular motion and the great distances between molecules of a gas, one gas will diffuse or mix with another, even against the force of gravity. When a gas of characteristic odor is released into a room, the odor will "move" through the room. This proves that gaseous molecules have sufficient energy to move through air. We call this process diffusion. From a quantitative point of view, the kinetic theory of gases demands that the average kinetic energy of molecules of two-different gases (A and B) to be equal at the same temperature

(1)

Where MA and MB are molecular masses of gases A and B, VA and VB are average speeds of the gas molecules of A and B, respectively. Rearranging and cancellation in equation one obtains

Heavier gases diffuse at lower rate than lighter gases at a given temperature.

37

This mathematical relationship, between the rates of motion and the densities (or molecular weights), of gases is expressed in Graham's law. This law states the rate of diffusion of gases vary inversely as the square roots of their densities (or molecular weights). An example of this ratio may be shown by comparing hydrogen (H2) with oxygen (O2),

Graham's Law of Diffusion may be used for the determination of molecular mass of an unknown gas by comparison of its rate of diffusion with another gas of known molecular mass at a given condition. 3. APPARATUS A glass tubing about 15 mm in diameter and about 80 cm long with two corks that can fit the tube, cotton (tissue paper), metal stand, clamps, ruler, medicinal droppers, test tubes, and stoppers 4. CHEMICALS Conc NH3, Conc.HCl, distilled water and phenolphthalein 5. PROCEDURES 5.1. Diffusion rates Clamp the glass tubing in a horizontal position on a table (Fig 5.1 )

Fig.5.1 Insert pieces of cotton at the two ends of the tube. c. Carefully add 5 drops of ammonia on the cotton pad (tissue paper) at one end while another student adds 5 drops of hydrochloric acid at other end exactly at the same time and immediately close the two ends with corks simultaneously and record the time.38

Do not push the corks with force! Do not drop hydrochloric acid and ammonia solutions on your skin and clothes. Avoid inhaling ammonia and hydrogen chloride gasses. d. Closely watch till a white" ring is formed and record the time at which the white ring is formed. e. Measure the distances between the white ring and the two ends.

f. Clean the tube by pushing a clean cotton pad (tissue paper) with the help of a glass rod. g. Repeat the experiment twice more

5.2 The Motion of gas moleculesa. Arrange two test tubes as shown in Fig.5.2. b.

Pour 2.5 mL of conc. NH3 into test tube A, and 5 mL of water Phenolphthalein, before inserting the stoppers to connect the tubes

c. d. e.

Insert the stoppers. Allow the apparatus to stand for several minutes. Observe the change and record results.

Fig.5.2 Learning goals Having done this experiment makes sure you understood the following: Graham's Law of Diffusion, composition of the white ring in the experiment, kinetic energy of molecules, and molecular velocity (motion)

39

LABORATORY REPORT FORMAT EXPERIMENT - 5 DIFFUSION OF GASES: MOLECULAR MOTION

Date: ______________ Group: ____________ Sub group:_________ Name:_______________________ ID.:____________________ Partner(s) ____________________ Objective: _____________________________________________________________________ _____________________________________________________________________ Theory: Not in more than 4 lines write on the relationships between molecular mass and velocity of molecules including Graham's equation for diffusion. Apparatus: ________________________________________________________________ Chemicals: ______________________________________________________________ Observations The white ring was due to the formation of Equation + _____________________ Equation _______________ + __________________________ The white ring expands to the direction opposite the ______________ end.

gas

40

DATA 5.1 Diffusion rates Trial 1 2 3 Mean

1. 2. 3. 4. 5.

Time diffusion started (tl) Time white ring formed (t2) Time taken for white ring to be formed t=t2- t1 Distance moved by NH3 (y) (cm) Distance moved by HCl (x) (cm)

6. 7. Error calculation Practical result Theoretical result Possible source of errors________________________________________________________ ____________________________________________________________________________

Conclusion (Is the law verified or not?)____________________________________________ _____________________________________________________________________________

5.2 The motion of gas molecules Write your observation and discuss the results

QUESTIONS41

1. Do the following question based on the instruction

Calculate the ratio of the distance moved by ammonia to that moved by hydrogen Chloride b. In this particular experiment, do you think it is essential to note the time taken for the diffusion? c. Suggest an alternative expression for Graham's Law of Diffusion in terms of distance moved by the gases. d. In this experiment, which gas took longer time to form the white ring? e. Which gas moved shorter distance? 2. How would the experiment be affected if one of the corks were pushed by strong force after addition of the NH3 and HCl solutions? 3. ______________gas diffuse at the same rate as one isotope of hydrogen called ________________ under the same condition. 4. Arrange the following in their increasing rates of diffusion under the same conditions. H2, CO, CO2, N2, He, O2, SO25. In a laboratory experiment gas "X" was found to diffuse at a rate three times that of gas "Y"

a.

Calculate the molecular mass of Y in terms of gas X. Molecular mass of B = __________ x molecular mass of A 6. The rate at which a gas diffuses __________ when the temperature is increased. 7. How did an increase in temperature affect each diffusion distances? Did they change? Explain your results.

EXPERIMENT- 6 MOLAR VOLUME OF A GAS42

1. OBJECTIVETo determine the volume of 1 mol of a gas at standard pressure and temperature, i.e. at 1 atm and 0 C

2. THEORY At standard temperature and pressure (STP), one mole of an ideal gas occupies 22.4 L. This can done by measuring the volume of a given mass of a gas at conditions different from STP conditions since this volume can be easily converted to STP conditions using the relationships obtained in the combined gas law equations derived from Charles' Law and Boyle's Law P1V1 = POVO (1) T1 TO Where V1, P1and T1are the experimental volume, pressure and temperature respectively. VO represents the volume of the same mass of gas at the standard pressure, PO, and temperature, TO. The volume of the gas at standard conditions will therefore be Vo = P1V1 X TO (2) T1 PO If the mass of the gas whose volume is determined at STP is m in g, and then the volume of 1 mol of the gas will be is (M/m) x V0 ) where M is the molecular weight (glmol). Hence, if molecular weight of the gas is known, the volume of one mol of the gas (molar volume of gas) is easily calculated. Experimentally, when a gas is collected over water, the contribution of the partial pressure of water, Pw should be corrected from the total atmospheric pressure (Patm) to obtain the partial pressure of the gas, P1 ;by the following relationship P1=Patm -Pw (3) 3. CHEMICALS 3% solution (w/w) of hydrogen peroxide, 0.1 M K2Cr2O7 4. MATERIALS A burette, a 'Y' tube, rubber tubing, funnel, metal stand and clamps

5. PROCEDURE5.1. 5.2.

Set up the apparatus as shown in Fig. 6.1 Disconnect the 'Y' tube and using a pipette carefully, add 2 mL of K2CrO7 solution into one of its arm.43

Carefully add 5 mL of3 % H2O2 (w/w, d = 1 g/mL) into the other arm of the 'Y' tube. Using another clean pipette (no mixing of the two solutions at this stage! 5.4. Adjust the water level in the burette to a zero reading. 5.5. Reconnect the 'Y' tube to the system.5.3.

Fig. 6.1 Assembly for determination of molar volume of a gas (Important!) Level the funnel at that position for sometime and if the water level in the altered for one or two minutes then the system is tight. 5.7. Move the funnel up and down so the water level in it is equal to the water in the burette and fix the position of the funnel when this is so. 5.8. Tilt the Y' tube to add the K2CrO7 solution into the H2O2 solution (but not vice versa).Wait until the decomposition of H2O2 according to the following equation:5.6.

12H2O2 + 2CrO4 2-+ 16H+5.9.

2Cr3+ + 20 H2O (l) + 9O2 (g)

(4)

If some bubbles remain carefully, shake the 'Y' tube until they escape. Repeat this until all the oxygen bubbles escape 5.10.When the level of the water in the burette remains unchanged for a few minutes, move the funnel up and down to equalize the water levels in the funnel and the burette. 5.11. Record the reading of the water level on the burette. 5.12. Wash the 'Y' repeat the experiment once again. 5.13. Note thermometer and the barometer readings.

6.

Calculations

Calculate the molar volume of oxygen in L using the equation44

Where V1 is the volume in mL, m = mass of oxygen from the 3% (w/w) H2O2 and M = molecular weight of O2 .Use stoichiometry to calculate mass of oxygen from H2O2

Learning Goals After completing this experiment, the student is expected to use the combined gas law equations in such experiments and be able to understand the physical principles behind this.

LABORATORY REPORT FORMAT EXPERIMENT- 6 MOLAR VOLUME OF A GAS Date: ___________________45

Group: __________________ Subgroup ________________ Name: ____________________________________________ ID.:_____________________________________________ Partner(s):__________________________________________ Objective: __________________________________________________________________

Theory :( State the definitions of molar volume of a gas, derive the formula for the calculation of molar volume of a gas in this experiment) Apparatus: __________________________________________ Chemicals: ______________________________________________ Observation: (Write your observations in three lines) Data 1. Mass of oxygen from the given volume of 3% (w/w, d= l g/mL) H2O2 _____________g 2. Initial reading of water level in the burette ______________mL 3. Final reading of water level in the burette _____________mL 4. Volume of oxygen collected. __________mL 5. Temperature of the room____________ oK 6. Atmospheric pressure ( Patm ) ____________ mmHg 7. Standard temperature ___________oK 8. Standard pressure (Po) ______________ mmHg 9. Vapor pressure of water at T1 __________________mmHg 10. Volume of O2 collected at STP ________________mL show calculation below ___________________________________________________________________________ __________________________________________________________________________ 11. Volume of oxygen at STP ________________________________________________ 12. Theoretical molar volume of O2 ______________________________________L Error calculations %E= practical value theoretical value Practical value

the

Therefore, %E__________________ Possible source of error: _________________________________________________________________ ____________________________________________________________________________________ Conclusions: __________________________________________________________________________

QUESTIONS46

1. Do you think the solubility of O2 in water significantly affects the results of this experiment? If yes,

how? 2. Explain why is it necessary to equalize the levels of water in the funnel and in the burette before taking the reading?3. In this kind of experiment, the Vapor pressure of water is taken care of, but the volume of water vapor is

not taken into account. Do you expect this effect to be significant?4. What is the maximum volume of CO2 obtained by thermal decomposition of 10 g of CaC03 after

adjusting the gas to 1 atm and 10OC?5. At 20OC and 580 mmHg, 30 mL of N2 (g) is collected over water. What would be the volume of the dry

gas at STP?6. What is the mass of 10 L of H2 at 380 mmHg and 10OC

EXPERIMENT- 7 COLLIGATIVE PROPERTIES - FREEZING POINT DEPRESSION 1. OBJECTIVE To determine how a change in a colligative property depends on the amount of solute present in a solution. 2. THEORY47

The properties of solution are different from those of pure solvents. Water, the most common solvent, has a boiling point (TB) of 100OC under a pressure of one atmosphere. At 100OC water molecules have, enough energy to escape from the liquid state by overcoming the pressure of air molecules.

Figure 7.1 Liquid and gaseous water molecules Gaseous molecules in equilibrium with molecules in the liquid state exert a definite pressure above the liquid at a definite temperature. This pressure, called the vapor pressure, increases with temperature until when equal to the atmospheric pressure the liquid boils.

Fig.7.2 The vapor pressure of water

Each liquid has a definite boiling point (TB) depending on how much strong intermolecular attraction exists between molecules of the liquid. Since water molecules attract each other, greatly high temperatures are required for water to boil. Dissolving solute molecules or ions in a solvent such as water serves to decrease the number of water molecules on the liquid of the solution resulting in a lowering of the escape tendency of water molecules into the gaseous state. This lowering of the vapor pressure is an example of a change in a colligative property. A colligative property is dependent on the ratio of solute to solvent particles in solution. Examples of such properties are the boiling point and the freezing point. Adding solutes to solvents also serves to reduce the number of solvent molecules at the surface of the solution again reducing the escape tendency of solvent molecules.48

As liquid cool molecules loose, enough energy to fall into a solid crystal pattern at a temperature called the freezing point (Tf). Reversing the process by warming a solid result in a solid to liquid change at the same temperature known as the melting point. Dissolving a solute in a solid or a liquid results in molecules or ions of solute reducing the ease with which solvent molecules can crystallize. This effect of lowering the freezing point is another example of a change in a colligative property.

Fig.7.3 Vapor pressure lowering for solutions

Figure 7.3 illustrates the raising of the boiling point to a temperature TB' and the lowering of the freezing point to a temperature (Tf,) by the addition of a solute to a solvent. These observations are true when the solute molecules exert insignificant vapor pressure of their own. Such solutes, usually solids, are known as non-volatile solutes. Solutes are added to water to make coolant solutions, which prevent car radiator from overheating in the summer and freezing in the winter. While lakes and streams may, freeze in winter oceans may remain as solutions due to high concentrations of dissolved salts. Sidewalks and highways are kept from freezing by the addition of salt solutes, which lower the temperature at which an ice solution will freeze.

Figure 7.4 illustrates how the temperature drops for a liquid that is cooling. At the melting point (freezing point), the temperature remains constant for a period until the liquid phase is converted entirely to a solid phase.

49

Figure.7.4 Change of state diagram In this experiment, you will be asked to observe how the change in melting point of a solvent depends upon the addition of a non-volatile solute. This observation will be made using several solutes. 3. APPARATUS Timer, (stop watch), ring stand, clamp and ring, 8 inch test tube fitted two-hole rubber stop ,400 mL beaker, thermometer ( 100 C ) graduated to 0.1 C, and wire stirrer (copper or nichrome). 4. CHEMICALS Naphthalene, sulphur, urea, bromocamphor, and diphenyl 5. PROCEDURE Your instructor will arrange the class into four groups: A, B, C, and D. Each group will use the same solvent, naphthalene, but a different solute. The change in freezing points, Tf recorded in Table 7.1 will be averages reported by each group.5.1.

Set up the apparatus as shown in Figure 7.5.

5.2. Add 8.0 g of naphthalene to an 8 inch test tube.5.3.

Heat the naphthalene in a water bath to a temperature of 90C and then take temperature

reading every 15 seconds until the temperature drop to 70C. Record the date in Table 7.1

50

Fig.7.55.4.

Depending upon the group to which you are assigned add the amount of solute shown in Table 7.2 to the naphthalene. Again, heat the contents until a temperature of 90C is reached and record temperature-time data every 15 seconds in Table 7.1 until the temperature of the solution drops to 70C. Members of group A will need to heat the test tube containing the sulphurnaphthalene mixture directly to dissolve the sulphur. Be certain to remove the thermometer from the test tube before heating the mixture.

Heat the mixture until clear solution results and then return the burner off. After waiting for a few minutes carefully, check the solution temperature making sure that the mercury in the thermometer does not rise above 100OC. When the temperature drops to 90OC proceed as described earlier to make temperature-time data.5.5. 5.6.

Collect the change in freezing point (Tf) from the students in your group, compute the average and enter into Table 7.1. Your instructor will announce the results from each group. Complete Table 7.2

Leaning goals After doing the experiment you should be able to define: Colligative properties and discuss the different colligative properties

LABORATORY REPORT FORMAT EXPERIMENT -7 COLLIGATIVE PROPERTY - FREEZING POINT DEPRESSION

Date: ___________ Group: _________Sub group: __________

Name:___________________________________ ID.:____________________________________ Partner(s):_______________________________51

Objective: __________________________________________________________________ ___________________________________________________________________ Theory: Discuss the different colligative properties in short Apparatus: ___________________________________________________________ Chemicals:__________________________________________________________Observations and conclusion Copy the table is below and fills in their right columns. Table 7.1 Temperature-time data Naphthalene(8.00g) Temperature Time Naphthalene and solute Temperature Time

Freezing point: ______________C (Tf) Freezing point: ___________ C (Tf) Change in freezing point: __________Tf = Tf + Tf = _________

,

Table 7.2 Results Weight of Solute Molecular Weight of Added Solute 1.28 g 0.67 g 1.16 g 0.77 g 256 g 135 g 231 g 154 g I Number of Moles of Solute Added Number of Molecules Added Tf*

Group A B C D..

Solute sulphur urea Bromocamphor diphepyl

* Record the averages for each group here.52

QUESTIONS1. What conclusion can you draw about the number of moles solute added in each solution? 2. What conclusion can you draw about the number of molecules of solute added in each solution? 3. What conclusion can you draw about the approximate Tf for each solution? 4. What would be the effect on the freezing points of the solutions in this experiment if the number of moles

of solute were doubled?5. Chemical reference table list freezing point constants for many solvents

The constant kf is the ratio of Tf /m where m is defined as moles of solute/ kg of solvent. Use your datato find this ratio for all four solutes in naphthalene.

EXPERIMENT - 8 HEAT OF REACTION 1. OBJECTIVE To estimate the heat of reactions in the dissolution of NaOH and the reaction between NaOH and hydrochloric acid . 2. THEORY A chemical reaction may be exothermic .or endothermic, depending on the heat evolved or absorbed in the reaction. In an exothermic reaction, heat is liberated or lost by the reacting system and this is absorbed by the surrounding, their by increasing the temperature of the latter. When a reaction takes place endothermically, heat is gained by the reacting system due to loss of heat from the surrounding, resulting in a decrease in the temperature of the latter. The heat of a reaction is denoted- by H and, - in53

the case of exothermic reactions, the quantity of heat liberated is by convention assigned a negative value to designate the heat loss, while a positive- value is assigned for endothermic reactions: The change in the temperature of the aqueous solution and the glass container may be used to estimate the heat of reaction. Hence, other small losses to the surroundings are neglected. From the definition of specific heat of water, 1.0 calorie changes the temperature of one gram of water by 1 OC . It takes 0.2 calorie to change the temperature of one gram of glass by 1C. The quantity of heat (q) expressed absorbed by a substance, in joules (SI unit), is calculated as follows q = Specific heat (joule deg-l g-l) X g substance x T (deg) (1)

In this experiment, the quantity of the heat evolved in three reactions will be estimated from the heat absorbed by the aqueous solution and an Erlenmeyer flask used as a simple calorimeter. Reaction 1 Solid sodium hydroxide dissolves in water to form an aqueous solution of ions: NaOH(s) Na+(aq) + OH-(aq) H1 = -X1 joule (2)

Reaction 2 Solid sodium hydroxide reacts with an aqueous solution of hydrogen chloride to form water and an aqueous solution of sodium chloride: NaOH(s) + H+(aq) + Cl-(aq) H2O + Na+(aq) + Cl-(aq) H2 = -X2 joule (3)

Reaction 3 An aqueous solution of sodium hydroxide reacts with an aqueous solution of hydrogen chloride to for water and an aqueous solution of sodium chloride: Na+(aq) + OH-(aq) + H+(aq) + Cl-(aq) H2O + Na+(aq) + Cl-(aq) H3 = -X3 joule 3. APPARATUS 250 mL Erlenmeyer flask, 150 mL beaker, balance and thermometer 4. CHEMICALS NaOH (s), 0.25 M hydrochloric acid and distilled water 5. PROCEDURES 5.1. Determination of Heat of Reaction 1 a. Weigh a clean, dry 250 mL Erlenmeyer flask to the nearest 0.1 g.54

(4)

b. Put 200 mL of distilled water (record mass 200 g) into the flask. Stir carefully with the thermometer

until a constant temperature is reached. Record this temperature.c. Weigh about 2 g (1.9- 2.1 g) of pellets of solid sodium hydroxide, NaOH. Since sodium hydroxide

becomes moist as it is being weighed in the open air, weigh out rapidly. d. Put the weighed NaOH(s) into the water in the Erlenmeyer flask. Swirl the flask until the sodium hydroxide is dissolved. Place the thermometer into the flask and record the highest temperature reached. Before proceeding to Reaction 2, rinse the 250 mL flask thoroughly with water. Do not dry. Rinse and dry the thermometer before transferring it from one solution to another 5.2. Determination of Heat of Reaction .2 Repeat steps b, c, and d used in the determination of the heat for Reaction 1 BUT in step b replace the water by 200 mL of 0.25 M HCl. Rinse the 250 mL flask again and proceed to Reaction 3 5.3. Determination of Heat of Reaction 3 a. Measure 100 mL of 0.50 M HCI into the 250 mL flask and 100 mL of 0.50 M NaOH into a 250 mL beaker. Record their temperatures. b. Add the sodium hydroxide solution to the hydrochloric acid solution, mix quickly and record the highest temperature reached. Learning Goals After studying this experiment, the student is expected to: Appreciate that chemical reactions generate measurable quantities of heat energy.

LABORATORY REPORT FORMAT EXPERIMENT - 8 HEAT OF REACTION Date: _____________________ Group: ___________________ Sub group: ________________ Name :______________________________ ID.:_______________________________ Partner(s):___________________________ Objective:______________________________________________________________________ ______________________________________________________________________55

Theory: ________________________________________________________________________ ________________________________________________________________________ Calculations: (Specific heat of water 4.2 J deg-l g-l, specific heat of glass 0.84 J deg-l g-) Record the masses of water and flask, mol of NaOH, and T in the table below. Calculate the heat absorbed by the water and the flask using equation 1 (3 lines) Calculate the heat H1, H2, H3, lost (show sign) per mol of NaOH (3 lines) Record calculation results in the table below. Qty measured or calculated Mass water or aqueous solution, g Mass of flask, g T(OC) Heat absorbed by solution, J Heat absorbed by flask, J Total amount of heat absorbed, J Amount of NaOH, mol Heat liberated mor-1 NaOH H1= __ J/mol NaOH H2=____ J/mol H3=___J/mol NaOH NaOH Reaction 1 Reaction 2 Reaction 3

Conclusion(a) Theoretically, H2 should be equal to H1 + H3.Why ? (b) Calculate the percent difference (% error) between H1 + H3 and H2, assuming H2 is correct.

QUESTIONS Write the net ionic equation for Reactions 2 and 3 (equations 3 & 4) and observe their relation with reaction 1 (equation 2).1. 2. 3.

In Reaction 1, H1 represents the heat of solution of NaOH(s). Look at the net ionic equations for reactions 2 and 3 and make a statement concerning the significance of H2 and H3. Suppose you used 4 g of NaOH(s) in Reaction 1.56

(a) What would be the number of calories evolved?

(b) What effect would this have on your calculation of H1, the heat evolved per mol of NaOH?

In this experiment, the glass beaker is used as a "calorimeter', assuming it would attain a uniform body temperature and no loss in heat. What do you expect, negative or positive error?4.

EXPERIMENT- 9 ACID-BASE TITRATIONS1. OBJECTIVE:

To find the normality of hydrochloric acid by titration against a standard sodium hydroxide solution 2. THEORY Titration is a method of chemical analysis made by volume measurement of a standard solution (titrant) required to react with an equivalent amount of a sample solution (titrand). Very often, titration is used to determine an acid or a base. Acidimetry is a titration when the acid is determined by a standard base while alkalimetry is the vice versa. In acid-base titrations, equivalent amounts of an acid (H3O+) and a base (OH-) are brought together and the net reaction may be represented by H3O+ + OH- 2H2O (1) The point at which equivalent quantities of the reactants are brought together is called equivalent point. At the equivalent point of an acid-base titration, the number of milliequivalents (1 milliequivalents (meq)= one thousand of the equivalent) of the acid and the base are equal.57

meq acid = meqbase The number of milliequivalents of a substance in a solution is obtained by N x V = (meq/mL) x mL Where; N is normality and V volume in mL. Hence, at the equivalent point

(2)

(3)

Nacid Vacid = Nbase Vbase

(4)

Fig. 9.1 represents a titration curve for 50 mL of 0.1 N HCI against 0.1 N NaOH. A suitable acid-base indicator changes colour sharply at or near the equivalent point. The point of the titration at which color change of the indicator occurs is known as end-point. The relative indicator error introduced by the variation from the equivalent point is evaluated from the volume difference between the indicator determined end-point (Vind) and the stoichiometric equivalent point (Veq) is simply calculated as % E= (Vind Veq )/ Veq x 100 (5)

Here are three indicators with color-transition pH ranges (pH ranges) that may be used: bromthymol-blue (pH 6.8 - 8.0, yellow to blue), phenolphthalein (pH 8.0 - 9.6, colorless to pink) and methyl orange (pH 3.1 4.4 red to yellow). Try to see which indicator can best be used in the titration of 0.1N HCl verses O.1 N NaOH.

Fig. 9.1 Titration curve of 0.1 N HCl verses 0.1 N NaOH

58

3. APPARATUS 10 mL pipette, burette, 150 mL Erlenmeyer flask, beaker, funnel, burette clamp and metal stand. 4. CHEMICALS Hydrochloric acid of unknown normality, 0.10 N sodium hydroxide, phenolphthalein indicator 5. PROCEDURE5.1. Obtain a burette and clean it with distilled water. If it is not yet cleaned, use a cleansing solution and

rinse with distilled water. Make sure the burette does not leak and the stopcock moves freely. Otherwise, apply some grease to the stopcock. Rinse the burette with the 0.10 N sodium hydroxide and fix the burette on the burette clamp in a vertical position (see set-up in Fig.9.2) 5.2. Using a funnel, introduce 0.10N sodium hydroxide into the burette. Allow some of the solution to flow out and make sure there are no air bubbles in the solution (Why?). Record the level of the solution, corresponding to the bottom of the meniscus, to the nearest 0.1 mL 5.3 Measure out exactly 10 mL of hydrochloric acid (sample), from a 10 mL pipette, into a clean 150 mL Erlenmeyer flask and add two or three drops of phenolphthalein indicator.

Caution! When you suck hydrochloric acid (or any reagent solution) into a pipette, have the maximum caution not to bring into your mouth Titration: First, hold the neck of the Erlenmeyer flask with one hand and the stopcock with the other. As you add the sodium hydroxide solution, swirl the contents of the flask gently and continuously. Add the base until the first faint pink color comes which disappears upon swirling. Drop by drop; add more sodium hydroxide gently until the pink color persists for a few seconds. Record the volume of sodium hydroxide from the difference between the initial level and the level reaching after the end point. 5.5 Discard the titration mixture and rinse the flask with water. Make two more titrations following 5.3 5.4. Refill the burette with the same base, in the course of the titrations, before the level of the solution is below the lowest mark of the burette.5.4

59

Fig. 9.2 Assembly of a typical titration set-up

Learning goals From this experiment and your previous background, you are expected to understand the terms and the use of apparatus underlined: Acid, acid-base indicator, Acidimetry, alkalimetry, base, burette, equivalent and Milliequivalents, equivalent point, equivalent weight, end-point, meniscus, neutralization, Normality, pH range, pipette, standard solution, titrand, titrant, titration, volumetric flask

LABORATORY REPORT FORMAT EXPERIMENT- 9 ACID-BASE TITRATIONS Date: ___________ Group: _________ Sub group: _________ Name: _______________________________ ID.:_______________________________ Partner(s):___________________________ Objective: ____________________________

Theory: (Write the principles of titration as applied to acids and bases in less than seven lines, including a chemical equation). Apparatus:_____________________________________________________________ Chemicals:_____________________________________________________________ Observation: Color change at the end-point is from________________ to___________________60

Data: Titration of hydrochloric acid by standard sodium hydroxide solution1. 2. 3.

Volume of HCl solution, (Va)= _________________mL Volume of NaOH solution, (Vb) = ___________: 2)_________3)_______ mL Normality of NaOH solution, (Nb)= ___________________meq/mL

Calculations1. Average volume of n titrations for NaOH solution, Vb =____________ 2. Normality of HCI solution, Na =meq/mL ( use NaVa = NbVb)

Show your calculations below Possible sources of error:________________________________________________________________

Conclusion:_______________________________________________________________________

QUESTIONS1. Show that the unit of normality, N, can be expressed in terms of equivalent/L or milliequivalents mL 2. 4.0 g of sodium hydroxide (NaOH) is dissolved to make a 125 mL solution. What are the normality

and the morality of the base? 3. 0.10 N hydrochloric acid is used to titrate (i) 10 mL of 0.10 N of sodium hydroxide and (ii) 10 mL of 0.1 N of sodium hydroxide after dilution to a larger volume, say to 40 mL with What is the volume of the acid required in each of the titration?4. The mass of a solute in a solution is calculated by the equation

water.

g(solute) = N(meq) mL

x V(mL) x( g ) meq

Find the mass, in gram, of sodium hydroxide required to prepare 100 mL of 0.2 N solution of the base in water.5. If phenolphthalein indicator is used in a titration of a given unknown normality of sodium hydroxide

solution against standard hydrochloric acid from a burette, the color change at the end-point would be from ____________________to ____________________________61

6. What are the similarity and the difference between equivalent point and end-point?

7. Which of the acid-base indicators discussed in the theory section can best be used for the 0.1 N HCl in this experiment?

EXPERIMENT- 10 THE EFFECT OF CONCENTRATION ON REACTION RATE 1. OBJECTIVE To observe how changing the concentration of a reactant will affect the speed of a reaction 2. THEORY The rate of a chemical reaction can be measured in many ways. In reactions where gases are involved the rate can be monitored by measuring the pressure change in a system, e.g., in mmHg per second. How fast the pressure increases in a system is a measure of how fast an increase in the number of molecules in the system is occurring. Measuring how fast the color of a solution changes is another method of measuring reaction rate. Colorimeters are capable of measuring the amount of light passing through a solution containing color absorbing species are used to measure reaction rates. In all cases where reaction rates are, changing the concentrations or number of reacting molecules, atoms or ions per unit volume influences the rate of a reaction. Molecular collisions between reacting molecules occur at grater frequency under conditions of higher concentrations thus increasing the speed of the reaction. In this experiment, hydrochloric acid reacts with sodium thiosulphate (Na2S203) to produce colloidal sulphur as one of the products. At some point in the reaction, the sulphur content of the solution results in an opaque solution. The time for each solution to become opaque will be a measure of reaction rate. The reaction is: 2HCI + Na2S2O3 SO2 (g) + S(s) + 2Na+ + 2Cl- + H2O

62

3. APPARATUS

Four 100 mL beakers, one 400 mL beaker, timer, and 100 mL graduated cylinder. 4. CHEMICALS0.1 M Na2S2O3 (sodium thiosulphate), and 1.0 M HCl

5. PROCEDURE5.1. Label four 100 mL beakers as solutions 1, 1/2, 1/4, and 1/8 5.2. Add 50 mL of 0.1 M Na2S2O3 to the beaker labeled solution 1. 5.3. Add 50 mL of water to 50 mL of 0.1 M Na2S2O3and mix (Obviously the concentration of this

diluted solution is 0.05 M N Na2S2O3Pour 50 mL of this solution into the beaker labeled solution 1/2. Save the rest of this solution for the next step.5.4. To the remaining 50 mL above (step 5.3) add 50 mL of water and mix. Add 50 mL this solution to

beaker labeled solution 1/4. Save the rest of this solution for the next step.5.5. Dilute the remaining solution above (step 5.4) with 50 mL of water; mix and add 50 mL of this

solution to the beaker labeled solution 1/85.6. Set your timer at 0.00 seconds. Quickly add 50 mL of 1 M HCI to the content of the beaker labeled

solution 1 and start the timer. When the solution become opaque note and record the time5.7. Repeat the above procedure for the remaining solutions of Na2S2O3

5.8. Tabulate your data. Learning goals After completing this experiment make sure, you have understood the following:63

The effect of concentration on reaction speed, molecular collisions, some common ways of measuring of rate of chemical reactions (pressure, color intensity, concentration etc.)

LABORATORY REPORT FORMAT EXPERIMENT - 10 THE EFFECT OF CONCENTRATION ON REACTION RATE

Date: ______________________ Group: ____________________ Subgroup: _________________ Name: ___________________________________ ID.:___________________________________ Partner(s):_______________________________ Objective: ____________________________________________________________________ _____________________________________________________________________ Theory: Discuss factors that affect rate of reactions. Apparatus: ____________________________________________________ Chemicals: ____________________________________________________ Data Solution 164

Reaction Time (min.)

1/2 1/8 Results and discussion

Using your data above plot initial concentration of Na2S2O3versus reaction time on a graph paper From your graph state, the general conclusions you can draw about the effect of concentration on reaction speed. Discuss possible sources of error.

QUESTIONS

1. From the curve you obtained in this experiment read the time necessary for reaction to be completed

at relative concentration of 1/5, 1/9 and 3/4, and calculated the ratio of time at relative concentration of 1/5 to time at relative concentration of 1/9 2. The Food and Drug Administration sets the permissible levels of toxic substances in foods. How do the results of this experiment relate to how these levels are set? Companies? Why the body weight of an individual is considered when medication is prescribed?4. Calculate the number of moles of Na2S2O3 in each milliliters of solution 1, 1/2, 1/4 and, 1/8.

3. How do the results of this experiment relate to drug dosages recommended by pharmaceutical

65

EXPERIMENT 11 REACTION SPEED AND TEMPERATURE 1. OBJECTIVE To study the effect of temperature on the rate of a chemical reaction 2. THEORY The speed of a chemical reaction is affected by the nature of the reactants, the presence or absence of a catalyst, temperature and concentration of the reactants. In this experiment, we will study the effect of temperature on the chemical reactivity of a chemical substance. The effect of temperature on reaction speed is quite noticeable. Generally, a 10C rise in the temperature of a reaction system approximately doubles or triples the rate of a chemical reaction. In the present experiment, the rate of disappearance of the color of potassium permanganate in acidic medium is taken as the measure of the oxidation of iron metal. The reaction steps are as follows 2H+ + Fe Fe2+ + H2 5Fe2+ + MnO4- + 8H+ 5Fe3+ + Mn2+ + 4H2O (1) (2)

Initially, metallic iron is oxidized to Fe2+ by the acid in' the solution which is further oxidized to Fe3+ by potassium permanganate. In experiments used to measure, rates of a reaction all factors are kept constant and one variable is studied at a time. In this experiment, the amount of iron, the concentrations of the acid and potassium permanganate are kept constant, while the temperature of the reaction is varied. Hence, the rate oxidation of iron is determined by the temperature.66

3. APPARATUS

Balance, six test tubes (3 cm diam and 15 cm length), 600 mL beaker, 100 mL measuring cylinder~, metal stand and ring, wire gauze, burner, thermometer, test tube holder and extension clamp. 4. CHEMICALS Potassium permanganate, (1 % (w/v) in 7% (v/v) sulphuric acid, iron filings and distilled water

5. PROCEDURE1.1. Assemble the apparatus as shown in Fig. 11.1 and light the burner. Heat the water to 85C. This will

serve as a water bath for the desired temperature. Make sure that the temperature does not change and this applies for the other temperatures. 1.2. Weigh out 0.5 g of iron filings six times on pieces of paper.

Fig.11.1 Assembly of apparatus for studying reaction with temperature 1.3. Measure out 12.5 mL of acidified potassium permanganate solution into one test tube. 1.4. When the temperature reaches 85 C, immerse the test tube with a holder into the heated water and keep it for 5-7 min. Record the temperature. 1.5. Add 0.5 g of the iron filings to one of the test tubes. Record the time, to 1.6. Swirl the contents of the test tube as it is in the water bath. Again, record the time at which the color of the permanganate disappears (tf). Do not take the test tubes out of the water bath until the color fades away. Caution! Do not point the mouth of the test tube at yourself or any body else 1.7. Cool the water in the beaker to 75 C and put another test tube with 12.5 mL of fresh potassium permanganate solution in the water bath. Wait until the temperature of the solution becomes equal to the temperature of the water in the beaker. Then repeat steps 5.5 and 5.6. 1.8. Cool the water in the beaker to 65 OC, and repeat steps 5.4, .5.5, and 5.6. 1.9. Repeat steps 5.4, 5.5 and 5.6 at temperatures 55 OC, 45 OC and 35 OC, with 12.5 mL each of fresh67

potassium permanganate solution. 1.10. Tabulate your data. Learning goals After completing this experiment, make sure you understand the following: Rate of a reaction, factors that affect the rate of reaction, the effect of temperature on reaction speed

LABORATORY REPORT FORMAT EXPERIMENT - 11 REACTION SPEED AND TEMPERATURE Date: __________________ Group: ________________ Sub group: _____________ Name:___________________________________ ID.:_________________________________ Partner(s):_______________________________ Objective:________________________________________________________________ ________________________________________________________________ Theory: Write the factors that affect the rate of reactions and the specific reaction for this experiment, about 6Iines. Apparatus:________________________________________________________________ Chemicals:________________________________________________________________ Observations: The color of the reaction solution is due to____________ ion and when it fades away it is changed to_____________________ In your experiment at what temperature does the color takes the shortest time to fade away? At __________________ OC Data (reproduce the Table below and fill in the data)

68

Results and discussion Using your data, plot the time taken for the color to disappear (y-axis) versus temperature (x-axis) on a graph paper. Join the points with a line. From your graph state, the general conclusions you can draw about the effect of temperature on reaction speed.

Discuss the possible sources of error: Data Table Trial mL of KMn04 solution Color Iron fading added (g) time (tf- tO) Temperature

1. 2. 3. 4.

QUESTIONS1. Indicate the oxidizing and the reducing agents in the above experiment, 2. Do you think the above experiment works if the permanganate solution was not initially acidified? 3. From the curve you obtained in this experiment (a) estimate time necessary for the reaction to be

completed (tf- tO) at 40C, 58 C, 70C, 76OC and 80C and (b) calculate the ratio of time intervals for 80C to that for 70 c.4. Does the graph, from your experiment above, demonstrate that the reaction speed doubles for every

increase of 10 C?5. If you double the concentration of the reactants and repeat an experiment at one of the temperatures

you applied, do you expect the rate would increase?6. A reaction at 30C requires 100 s for completion. If the speed is doubled for each 10C rise

temperature, what is the time necessary for completion at (a) 60C; (b) 10 C ?

69

EXPERIMENT- 12 CHEMICAL EQUILIBRIUM AND LE CHATELIER'S PRINCIPLE I 1. Objective To observe the shift of chemical equilibrium in response to concentration change or species at chemical equilibrium 2. Theory A chemical system is said to be in equilibrium when the rate of the forward reaction is equal to that of the reverse reaction. If reactant species A and B yield products C and D, Reversible arrows in a chemical equation represent a chemical equilibrium as foll