Chemistry lab manual

School of Biotechnology, International University-HCMCDepartment: Applied Chemistry Laboratory ManualGeneral Chemistry

LABORATORY MANUAL

FOR

GENERAL CHEMISTRY FOR ENGINEERINGSTUDENTS


This laboratory is designed for internal use only. Its content is collected and composed by DrHoang Le Son; edited by Dr Huynh Kim Lam for Laboratory Module.

*Some of its contents are extracted from references as indicated.

References:

[1] Experiments in General Chemistry: Inquiry and Skill buildingby ickie Williamson (Author), LarryPeck, (Brooks/Cole Laboratory Series for General Chemistry), 2008.[2] Introduction to Chemical Principles: A Laboratory Approach by Susan A. Weiner, Blaine Harrison(Brooks/Cole Laboratory Series for Introductory Chemistry) 7Edition, 2009.[3] Laboratory Experiments for General, Organic and Biochemistryby Frederick A. Bettelheim(Brooks/Cole Laboratory Series for General Chemistry) 7th Edition, 2009.[4] Experiments in General Chemistry by Bobby Stanton, Lin Zhu, Charles H. Atwood (FeaturingMeasureNet), 2nd Edition, 2009.[5] Laboratory Manual for Principles of General Chemistry by Jo Allan Beran, 9th Edition, 2010.

Common LaboratoryGlassware and Equipment

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electronic balance buret clamp graduated thermometercylinder

beaker Bunsen burner clay triangle

crucible tongs crucible and lid Erlenmeyer flask

evaporating dish funnel glass rod with rubber policeman BobbyStanton/W

adsworth/CengageLearning

x Common Laboratory Glassware and Equipment

ring stand and iron ring scoopula test tube rack

test tube clamp test tube utility clamp

watch glass wire gauze buret

pipets volumetric

flask BobbyStanton/W

adsworth/CengageLearning

Practice Safe Laboratory

A few precautions can make the laboratory experience relatively hazardfree and safe. These experiments are on a small scale and thus many of thedangers found in the chemistry laboratory have been minimized. Inaddition to specic regulations that you may have for your laboratory, thefollowing DO and DONT RULES should be observed at all times.

D O R U L E Sq Do wear approved safety glasses or goggles at all times.

The rst thing you should do after you enter the laboratory is to put onyour safety eyewear. The last thing you should do before you leave thelaboratory is to remove them. Contact lens wearers must wearadditional safety goggles. Prescription glasses are not suitable safetyglasses; you must wear safety goggles over them.

q Do wear protective clothing.Wear sensible clothing in the laboratory (i.e., no shorts, no tank tops, nosandals). Be covered from the neck to the feet. Laboratory coats oraprons are recommended. Tie back long hair, out of the way of ames.

q Do know the location and use of all safety equipment.This includes eyewash facilities, re extinguishers, re showers, andre blankets. In case of re, do not panic, clear out of the immediatearea, and call your instructor for help.

q Do use proper techniques and procedures.Closely follow the instructions given in this laboratory manual. Theseexperiments have been student tested; however, accidents do occur butcan be avoided if the steps for an experiment are followed. Pay heed tothe Caution! signs in a procedure.

q Do discard waste material properly.Organic chemical waste should be collected in appropriate waste contain-ers and not ushed down sink drains. Dilute, nontoxic solutions may bewashed down the sink with plenty of water. Insoluble and toxic wastechemicals should be collected in properly labeledwaste containers. Followthe directions of your instructor for alternative or special procedures.

q Do be alert, serious, and responsible.The best way you can prepare for an experiment is to read theprocedure carefully and be aware of the hazards before stepping footinto the laboratory.

xi

COPYRIGHT

2010Brooks/Cole,CengageLearning

D O N T R U L E Sq Do not eat or drink in the laboratory.

Consume any food or drink before entering the laboratory. Chemicalscould get into food or drinks, causing illness. If you must take a break,wash your hands thoroughly before leaving.

q Do not smoke in the laboratory.Smoke only in designated smoking areas outside the laboratory.Flammable gases and volatile ammable reagents could easily explode.

q Do not taste any chemicals or breathe any vapors given off by areaction.If there is a need to smell a chemical, you will be shown how to do itsafely.

q Do not get any chemicals on your skin.Wash off the exposed area with plenty of water should this happen.Notify your instructor at once. Wear gloves as indicated by yourinstructor.

q Do not clutter your work area.Your laboratory manual and the necessary chemicals, glassware, andhardware are all that should be on your benchtop. This will avoidspilling chemicals and breaking glassware.

q Do not enter the chemical storage area or remove chemicals from thesupply area.Everyone must have access to the chemicals for the days experiment.Removal of a chemical from the storage or supply area only compli-cates the proper execution of the experiment for the other students.

q Do not perform unauthorized experiments.Any experiment not authorized presents a hazard to any person in theimmediate area.

q Do not take unnecessary risks.

These DO and DONT RULES for a safe laboratory are not an exhaustivelist, but are a minimum list of precautions that will make the laboratory asafe and fun activity. Should you have any questions about a hazard, askyour instructor rstnot your laboratory partner. Finally, if you wish toknow about the dangers of any chemical you work with, read the MaterialSafety Data Sheet (MSDS). These sheets should be on le in the chemistrydepartment ofce. A sample sheet is included here so you know what onelooks like. This is the MSDS for glucose. Read it and see the kind of dataincluded in there. Imagine all the additional cautions and precautions thatthe sheets would contain were you dealing with a chemical that is toxic orcarcinogenic.

xii Practice Safe Laboratory

Safety Quiz

Indicate whether each of the following statements is true or false by writing the word TRUE or FALSE inthe space provided.

_____ 1. If chemicals come into contact with your skin, immediately wash the affected area withcopious quantities of water.

_____ 2. Fume hoods are used in the chemical laboratory when using volatile or poisonouschemicals.

_____ 3. It is permitted to leave a lit Bunsen burner unattended.

_____ 4. Always return unused chemicals to a reagent bottle to avoid wasting chemicals, you willnot contaminate the entire reagent bottle.

_____ 5. When heating a liquid in a test tube, always point the test tube in a direction away fromany other person in the laboratory.

_____ 6. Always add boiling chips to a hot solution.

_____ 7. The wearing of shorts, tank tops, mid-riffs and sandals is permitted in the laboratory.

_____ 8. Drinking soda in the lab is permitted as long as the soda can is at least 10 feet away fromall chemicals.

_____ 9. I am not required to wear safety goggles while in the laboratory unless I am actuallyperforming an experiment.

_____ 10. It is a violation of Federal Law to leave a Waste Container uncapped.

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Name . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . Section . . . . . . . . . . . . . . . Date . . . . . . . . . . . . . . .

Instructor . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . .


Revision: October 1, 2013 Page 1

EXPERIMENT 1: CHEMICAL REACTIONS

1. OBJECTIVES To perform different types of chemical reactions including acid-base, precipitation, gas

forming, complex compound forming and oxidation-reduction reactions. To identify the products in these reactions and describe the chemical changes. To write and balance the chemical equations for the reactions observed.

2. INTRODUCTIONMatter can undergo both physical and chemical changes. Chemical changes result in theformation of new substances. When a chemical reaction occurs, substances called reactantsare transformed into different substances called products that often have differentappearances and different properties. In this experiment, you will perform nd observe anumber of chemical reactions. Observable signs of chemical reactions can be a change in color,the formation of a solid, the release of gas, and the production of heat and light. You will alsolearn how to classify the chemical reactions. One classification system involves five generaltypes of reactions: synthesis, decomposition, single displacement, double displacement, andcombustion.

3. PROCEDURE

1. REACTIONS OF Cu2+

Step1: Put 10 drops of 0.5M CuSO4 into each of three test tubes.Step 2:

o Test tube #1: add 10 drops of 2M NaOHo Test tube #2: add 10 drops of 2M NH4OHo Test tube #3: add 10 drops of 0.5M K4[Fe(CN)6].

Step 3: Record your observations (Remember to shake the test tubes gently).

2. REACTIONS OF SILVER HALIDES

Section 1: Reactions of KClStep 1: Prepare 03 test tubes each containing 10 drops of KCl.Step 2: Add 10 drops of 0.1M AgNO3 to each of test tube above.Step 3:

o Add nothing to test tube #1 for control .o Add 5 drops of 2M NH4OH to test tube #2.o Add 5 drops of 2M KCN to test tube #3.

Step 4: Record your observations. Remember to shake the test tubes well and wait for atleast 02 minutes.



Section 2: Reactions of KBrRepeat Section 1 with KBr instead of KCl at Step 1. Record your observations.

Section 3: Reactions of KIRepeat Section 1 with KI instead of KCl at Step 1. Record your observations.

3. REACTIONS OF H2O2

Section 1: Place 5 drops of 0.1M KMnO4 solution into a test tube. Acidify this solution with 5drops of 2M H2SO4 and then add 5 drops of 3% H2O2 solution. Record your observations.Remember to shake the test tubes well and wait for at least 02 minutes.

Section 2: Place 5 drops of 0.1 M KI solution into a test tube. Acidify this solution with 5drops of 2M H2SO4 and then add 5 drops of 3% H2O2 solution. Record your observations.

Section 3: Place 10 drops of 3% H2O2 solution into a test tube, then, add a pinch of solidMnO2. Record your observations.

4. REACTIONS OF NITRATE BROWN RING TEST

Section 1 Step 1: Prepare a test tube containing 10 drops of 0.1M NaNO3. Step 2: Add 10 drops of saturated FeSO4 to test tube above. Step 3: Pour concentrated sulfuric acid, H2SO4, (96%) carefully down the inside wall of

the test tube. Wait for a few seconds and record the change of color at the interfacebetween the nitrate solution and the concentrated sulfuric acid.

Note: The nitric acid is reduced to nitrogen monoxide by the iron (II) ion and the brownish violetnitroso complex compound is then found by nitrogen monoxide with the excess iron (II) ions.

Section 2Repeat Section 1 with NaNO2 instead of NaNO3 at Step 2. Record and compare theobservations.

Section 3Repeat Section 1 with acetic acid (CH3COOH) instead of H2SO4 at Step 3. Record andcompare the observations.

5. REACTIONS OF KMnO4 Step 1: Obtain 3 clean test tubes

o Test tube #1: 10 drops of 0.5M Na2SO3 and 5 drops of 2M H2SO4.



o Test tube #2: 10 drops of 0.5M Na2SO3 and 5 drops of 6M NaOH.o Test tube #3: 10 drops of 0.5M Na2SO3 and 5 drops distilled water.

Step 2: Add 5 drops of 0.1M KMnO4 to each of test tubes. Record your observations.

6. REACTION OF POTASSIUM DICHROMATE (K2Cr2O7)

Step 1: Place 10 drops of 0.5M K2Cr2O7 into a test tube. Step 2: Add 10 drops of 6M H2SO4. Step 3: Add 5 drops of C2H5OH. Record your observations.

7. REACTIONS OF Fe2+ and Fe3+

Section 1 Step 1: Place 10 drops of 0.5M FeCl3 solution in each of five test tubes. Step 2:

o Test tube #1: Add 5 drops of 0.5M KCN.o Test tube #2: Add 5 drops of 0.1M KSCN.o Test tube #3: Add 5 drops of 2M KOH.o Test tube #4: Add 5 drops of 0.5 M K4[Fe(CN)6].o Test tube #5: Add 5 drops of 2M NH4OH.

Step 3: Record your observations Section 2

Repeat Section 1 using FeSO4 instead of FeCl3 at Step 1. Record and compare the results.

8. REACTIONS OF Al3+

Step 1: Place 10 drops of 0.5M Al2(SO4) to each of 2 test tubes Step 2: Add 5 drops of 2M NaOH to each test tube above.

o Test tube #1: Add 20 drops of 2M HClo Test tube #2: Add 20 drops of 2M NaOH

Step 3: Record your observations.

9. FLAME TEST

Step 1: Light the alcohol lamp. Step 2: Dip a looped wire into one of the five solutions supplied (LiCl, NaCl, KCl, CaCl2

and BaCl2), and then hold it in the flame. Step 3: Record the dominant flame color observed. Step 4: Using the wavelengths shown below, calculate the frequency and energy of

the photons emitted during the flame tests.

Dominant color Wavelength (nm)Red 701Red-orange 622



Orange 609Orange-yellow 597Yellow 587Yellow-green 577Green 535Green-blue 492Blue 474Blue-violet 455Violet 423

Wavelength values are given for mid-range of the color indicated. relationship between thewavelength, frequency and speed of an electromagnetic wave is given by theequation:

C =

And the energy per photon (in Joules) is given by the equation:Ephoton = h

Where h is Plancks constant, which has a value of 6.626 x 10-34 J.s.

4. REPORT

Prepare the report in the following order:

Summary of theory Describe your experiments and observation Name the reaction type for each experiment

5. CHEMICALS AND EQUIPMENTS



EXPERIMENT 1: LAB STRUCTURE



EXPERIMENT 2: pH AND BUFFERS

1. OBJECTIVES

To distinguish between strong and weak acids To learn how to calculate and prepare a buffer solution and test its buffering ability.

2. INTRODUCTIONAcids are proton donors and bases are proton acceptors. Acids primarily serve as sources ofhydrogenions (H+ ) or hydronium ions (H3O+ ) while bases mainly provide hydroxide ions (OH-).

Water is amphoteric because it can play a role as either an acid or a base. In other words,water can donate and accept protons. Water undergoes auto-ionization to form H3O+ and OH-

2H2O H3O+ (aq) + OH-

The extent of dissociation of water is very small; therefore pure water has no electricalconductivity. At the equilibrium, the ion product of water is only 1 10-14 at 25o C.

Kw = [H3O+][ OH-]= 10-14

In pure water, the concentration of the [H3O+] and [ OH-] are equal:

[H3O+] = [ OH-] = 1 10-7 M

In acidic solutions, H3O+ ions predominate: [H3O+] > [ OH-]

In basic solution, OH- ions predominate: [H3O+] < [OH-]

For convenience the negative value of the log [H+] is used to express the concentration of H+.Therefore the pH can be defined as:

pH=-log[H3O+]

In neutral solutions, pH = 7

In acidic solutions, pH < 7

In basic solutions, pH > 7

As a consequence, pH denotes the strength of acids or bases. The lower pH, the more acidicthe solution whereas the higher pH the more basic the solution. Strong acids and strong basesare completely dissociated in water to produce hydrogen ions or hydroxide ions respectively.



Weak acids dissociate only partially and form little or very little H+. This means that anequilibrium isestablished between the dissociated and un-dissociated forms:HA(aq) H + A

= [H][A][HA]Where HA is the weak acid and A- is its conjugate weak base, of HA. The equilibrium constant(Ka) is called the acid dissociation constant or acid ionization constant. pKa is defined in a waycomparable to pH, i.e. pKa = - lgKa.

A buffer is a solution of a weak acid and its conjugate weak base. Buffers have the functionthat resists large change in pH on the addition of H+ or OH-. This is because the weak base, A-,will react with added H+ and the weak acid, HA, will react with added OH. Changes in pH ofbuffer solutions can be determined using the Henderson-Hasselbach equation:

pH = pK + log [A][HA]A pH meter can be used to measure the pH of prepared solutions. Different classes ofchemicals behave differently when dissolved in water. By doing this experiment, you will gain abetter understanding of strong acids and strong bases, weak acids and weak bases, salts andbuffers.

3. PROCEDURE

Note: Please make sure you rinse the pH meter after each measurement.

3.1 DEIONIZED WATER

Pour about 50 mL of the room temperature deionized water into a 150 mL beaker.Stirthe water. Assemble pH meter. Record the pH value.

Stir the water for about 20 seconds. Record the pH again Repeat the stirring and measurement process at 20 second intervals, recording each

time and pH value, until there is no appreciable change in the pH. Repeat the experiments at least 02 times.

3.2 STRONG ACID

Pipet 10 mL of 0.1M HCl into a 100 mL beaker. Measure the pH. Add 90 mL of distilled water. Measure the pH.



Add 10 mL of 0.10M NaOH. Record the pH. Add 90 mL of 0.01M NaOH. Record the pH.Repeat the experiments at least 02 times

3.3 WEAK ACID

Prepare three 250-mL beakero Beaker 1: Place 20 mL of 0.1M acetic acid (CH3COOH). Measure the pH and

calculate Ka.o Beaker 2: Place 20 mL 0.01M CH3COOH. Measure the pH and calculate Ka.o Beaker 3: Place 20 mL 0.001M CH3COOH Measure the pH and calculate Ka.

Repeat the experiments at least 02 times

3.4 SALTS : prepare 3 beakers

Place 50 mL of 0.1M NaCl into a 250 mL beaker. Measure the pH. Place 50 mL of 0.01M CH3COONa into a 250 mL beaker. Measure the pH. Place 50 mL of 0.1M NH4Cl into a 250 mL beaker. Measure the pH.Repeat the experiments at least 02 times

Note: in this experiment, you should prepare CH3COONA and NH4Cl solution with thecorrect concentrationy by yourself.

3.5 BUFFERS

Prepare 02 150 mL beakers each containing ~100 mL of 0.1M CH3COOH and 100 mL of0.1M CH3COONa, respectively.

Section 1: Prepare 50 mL buffer A in a 100 mL beaker by mixing 10 mL of 0.1M CH3COOH and 40

mL of 0.1M CH3COONa Measure the pH of the buffer solution.

Repeat the experiments at least 02 times.

Section 2: Divide buffer A into 2 equal parts (25 mL of each). Label them.

For part 1:o Add 10 drops of 0.1M HCl. Measure the pH.o Add enough to change the pH by one unit from the start. Record the

volume.Note: roughly 10 drops ~ 1 mL

For part 2:o Add 10 drops of 0.1 M NaOH. Measure the pH.o Add enough to change the pH by one unit from the start. Record the

volume.



Section 3: Prepare 50 mL buffer B in a 100 mL beaker by mixing 40 mL of 0.1M CH3COOH and 10


Repeat Section 2 with buffer B instead of buffer A.

Section 4: Prepare 50 m buffer C in a 100 mL beaker by mixing 25 mL of 0.1M CH3COOH and 25


Repeat Section 2 with buffer C instead of buffer A.


5. LAB REPORT Present your observation and obtained data for each part Present your explanation and conclusions



EXPERIMENT 3: OXIDATION/REDUCTIONTITRATION WITH KMnO4

1. OBJECTIVES

Learn about the term of gram equivalent weight. Review of oxidation-reduction reactions. Standardize the concentration of KMnO4 solution and determine the oxalic acid

normality.

2. INTRODUCTION

In an oxidation-reduction or redox reaction, there is an exchange of electrons between toreactants, resulting in the changes of oxidation number. The substance that gains electrons issaid to be reduced; therefore, it is called the oxidizing agent. The substance that loseselectrons is said to be oxidized; thus, it is called the reducing agent.

One gram equivalent weight (GEW) of oxidizing agent is the weight that gains 6.02x1023electrons and one gram equivalent weight of reducing agent is the weight that loses 6.02x1023electrons. According to the definition of gram equivalent weight, one GEW of oxidizing agentreacts with one GEW of reducing agent:

GEWox = GEWred

Consider the reaction of potassium permanganate (KMnO4) with oxalic acid (H2C2O4) in thepresence of excess sulfuric acid (H2SO4). The balanced molecular and net ionic equations are asfollows, respectively.

2KMnO4 + 5H2C2O4 + 3H2SO4 10CO2 + K2SO4 + 2MnSO4 + 8H2O2MnO4 + 5H2C2O4 + 6H+ 10CO2 + 2 Mn2+ + 8H2O

The oxidation number of Mn in MnO4- is +7 while it is +2 in Mn2+ . Hence, each Mn undergoes achange in oxidation number of five. Since each formula unit of KMnO4 contains one Mn, andeach Mn gains five electrons, one mole of KMnO4 is five gram equivalent weights in thisreaction. As a result, KMnO4 produces 5 moles of electrons per mole of KMnO4 or has fiveequivalents per mole of KMnO4. Thus, the gram equivalent weight of KMnO4 in this reaction is31.60 grams.

= . = .



The oxidation number of carbon in H2C2O4 is +3 while it is +4 in CO2. Thus each carbonundergoes a change in oxidation number of one. However, each formula unit of H2C2O4contains two carbons, and since each carbon loses one electron, one mole of oxalic acid is twogram equivalent weights in this reaction.

Consequently, H2C2O4 produces 2 moles of electrons per mole of oxalic acid or has twoequivalents per mole of oxalic acid. The gram equivalent weight of H2C2O4 is 45.0 grams

= . = .In this experiment, you will prepare an approximately 0.05N KMnO4 solution and standardizethis solution by titrating against a standard solution of H2C2O4 (primary standard). Then thestandardized KMnO4 solution (secondary standard) will be used to determine theconcentration of unknown oxalic acid solution and unknown Fe2+ solution. For redox titrations,the number of equivalents of oxidizing agent must be equal to the number of equivalents ofreducing agent. For the reaction of KMnO4 with H2C2O4:

Eq. of KMnO4 = Eq. of H2C2O4

Alternatively, this relationship can be expressed as follows:

Voxidizing Noxidizing=Vreducing Nreducing

where V is the volume of oxidizing or reducing agents used in titrations and N is the normalityof oxidizing or reducing agents.

At the end of a titration, three of the four variables will be known and the unknown variablecan be determined.

3. PROCEDURE

3.1. HANDLING WITH BURET:

Clean the buret with distilled water Rinse it three times with ~5 mL of the prepared KMnO4 solution. Discard the

rinse solution. Fill the buret with the KMnO4 solution and allow it to drain through the buret tip

until no air bubbles remain in the tip. Record the buret reading before beginning the titration.

Note: as the KMnO4 solution is dark color, read the buret at the top of the meniscus.



3.2. STANDARDIZATION OF PREPARED KMNO4 SOLUTION:

Pipette 10 mL of standard oxalic acid solution into each of two 250 mL erlenmeyerflasks.

Use cylinder to add approximately 40 mL of distilled water to each flask. Use cylinder to add approximately 20 mL of 6N H2SO4 solution to each flask (This step

must be done in the fume hood). Warm the flasks in the water bath 85oC 90oC (Use the thermometer to check the

temperature) Titrate the hot solution against the KMnO4 solution.

Note: the KMnO4 solution should be added very slowly initially. Endpoint for this experimentrefers to the titrate volume needed to keep the faint pink color throughout the stirredsolution for at least twenty seconds. Record the buret reading and calculate the normality ofthe KMnO4 solution.

3.3. DETERMINATION OF UNKNOWN CONCENTRATION H2C2O4 SOLUTION: Prepare 10 mL of the unknown concentration solution of H2C2O4 into each of

two 250 mL Erlenmeyer flasks. Use cylinder to add ~40 mL of distilled water to each flask. Use cylinder to add ~20 mL of 6N H2SO4 solution to each flask (fume hood). Warm the flasks in water bath 85oC 90oC. (Use the thermometer to check the

temperature) Titrate the hot solutions agains the KMnO4 solution. Calculate the normality of the unknown concentration H2C2O4 solution; determine the

average and the standard deviation.

3.4. DETERMINATION OF UNKNOWN CONCENTRATION FESO4 SOLUTION: Prepare 10 mL of unknown concentration solution of FeSO4 solution into each of

three 250 mL Erlenmeyer flasks. Add 40 mL of distilled water to each flask. Add 20 mL of 6N H2SO4 solution to each flask (fume hood). Warm the flasks in water bath 85oC 90oC. Titrate the hot solutions. Calculate the normality of the unknown concentration FeSO4 solution; determine the

average and the standard deviation.




7. LAB REPORT Present your observation and obtained data for each part Present your explanation and conclusions



EXPERIMENT 4: CHEMICAL EQUILIBRIUM1. OBJECTIVES

To observe the effect of applying stresses on chemical systems at equilibrium. To apply Le Chateliers Principle to explain the changes in the system

2. INTRODUCTION

A reversible reaction is at equilibrium when the rate of the forward reaction becomes equal tothe rate of the backward reaction. Reversible reaction:

A reversible reaction at equilibrium can be disturbed if a stress is applied to it. Stresses can bechanges in concentration, temperature or pressure. The composition of the reaction mixturewill shift until equilibrium has been reestablished. This is known as Le Chateliers Principle. Inthis experiment, the effect of applying stresses to a variety of chemical systems at equilibriumwill be observed and we also see if the results are consistent with Le Chateliers Principle.

3.PROCEDURE

1. SYSTEM 1: ACID/BASE EQUILIBRIA

Place 10 drops of 0.5M K2CrO4 to a clean test tube. Add 05 drops of concentrated HCl. Observe the change of color. And then add 10 drops of 6N NaOH. Record your observations.

Equilibrium System:2CrO42- + 2H+ (aq) Cr2O72- + H2O(l)

2. SYSTEM 2: EQUILIBRIA OF ACID/BASE INDICATORS

Place 2 drops of methyl violet to a clean test tube. Add 20 mL of distilled water, mixwell. Divide the solution evenly into two test tubes. Save one as a reference. Note thecolor. Test tube #1 (reference): add nothing Test tube #2:

o Addition #1: add the 6M HCl solution drop wise until further additionresults in no significant change. Observe the change.

o Addition #2: add the 6M NaOH solution drop wise until further additionresults in no color change. Observe the change.



o Addition #3: again add the 6M HCl solution drop wise until further additionresults in no significant change. Observe the change.

Equilibrium System:H(MV)(aq) + H2O(l) H3O+ (aq) + MV(aq)

3. SYSTEM 3: COMPLEX ION FORMATION

Preparation of iron (III) thiocyanate solution: in a 250 mL beaker, place 10 mL of 0.01M FeCl3 and 10 mL of 0.01 M KSCN, and then add 50 mL of distilled water, mix well.Use a pipet, divide the solution evenly among 07 similarly-sized test tubes (#1-7). Test tube #1 (control): add nothing Test tube #2: add 2 mL of 0.01 M FeCl3 to the solution. Observe the change.

Note that 1 mL 1012

Test tube #3: add 2 mL of 0.01 M KSCN to the solution. Observe the change. Test tube #4: add 10 drops of 6 M NaOH to the solution. Describe the change in

the solution. Test tube #5: cool the test tube in an ice bath. Test tube #6: warm the test tube in a hot water bath. Compare the intensity of

the color in test tubes #1(control room temperature), #5 (cold), and #6 (hot). Test tube #7: add 0.1 M AgNO3 solution drop by drop until all the color

disappears. Record your observations

Equilibrium System:Fe3+ (aq) + SCN-(aq) [Fe(SCN)]2+(aq)Pale yellow Clear Red

4. SYSTEM 4: EQUILIBRIA OF PRECIPITATION REACTIONS Place 5 mL of 0.05 M CaCl2 into each of the two test tubes labeled #1 and #2

Test tube #1: add 1 mL of 0.1M Na2C2O4 solution. Observe the change. Test tube #2:

o Addition #1: add 1 mL of 0.1M H2C2O4. Observe the change, comparing totest tube #1.

o Addition #2: add 10 drops of 6M HCl. Observe the change.o Addition #3: add 10 drops of 6M NH4OH. Observe the change.

Equilibrium System:Ca2+ (aq) + C2O42- (aq) CaC2O4(s)

5. SYSTEM 5: TEMPERATURE EFFECTS ON EQUILIBRIA

Place 3 mL (~30 drops) of 0.1 M CoCl2 into a test tube. Add concentrated HCl dropwise until the solution turns a purple-violet color. If the system turns a deep



blue, indicating too much chloride, discard the solution and start again. Note:this practice should be performed under the fume hood.

Divide the solution equally into three test tubes labeled #1-3. Test tube #1(control): keep at room temperature. Test tube #2: place in a hot water bath. Observe the change. Test tube #3: place in an ice-water bath. Observe the change.

Switch test tubes 2 and 3. Observe the change. Allow them both to cool to roomtemperature. Compare to the control.

Equilibrium System:[Co(H2O)6]2+(aq) + 4Cl-(aq) [CoCl4]2-(aq) + 6H2O(l)

4. CHEMICAL AND EQUIPMENTS

8. LAB REPORT Present your observation and obtained data for each part. Present your explanation and conclusions.



EXPERIMENT 5: REACTION RATE1. OBJECTIVES

To examine the effect of concentration, temperature, and catalysts on reaction rates.

2. INTRODUCTION

The rate of a chemical reaction describes how fast the reaction occurs. The rate of a chemicalreaction is affected by a number of factors including temperature of the reaction, the nature ofthe reactants, concentration of the reactants, the surface area of the reactants, the presenceof a catalyst and the pressure the reaction is under. The greater the rate of a chemicalreaction, the less time is needed for a specific amount of reactants to be converted toproducts. The rate of a reaction can be determined one of two ways; either measure the timeit takes for one or more of the reactants are used up, or for the products to be formed.

3. PROCEDURE

PART 1: EFFECT OF CONCENTRATION ON REACTION TIME

The solutions to be used are as follows:

Preparation of Solution A: 0.20M potassium iodide (KI) Preparation of Solution B: 0.005M sodium thiosulfate (Na2S2O3). This solution also

contains starch that will act as an indicator to detect the presence of iodine. Preparation of Solution C: 0.10M ammonium peroxydisulfate ((NH4)2S2O8)

In this reaction, solution B will be the limiting reagent. The reactions involved are these:

Reaction 1: 2I- + S2O82- I2 + 2SO42

Iodide ions + peroxydisulfate ions iodine + sulfate ions

Reaction 2: I2 + 2S2O32 2I + S4O62 Iodine + thiosulfate ion iodide ion + tetrathionate ion

Reaction 1 is relatively slow. As the iodine is formed it is quickly used in reaction 2, which isrelatively fast. The limiting reaction (solution B) is a source of the thiosulfate ions. Whensolution B is used up, the excess iodine formed will react with starch to form a deep bluesolution.

In this experiment, you will vary the concentrations of solutions A and C. The temperature willremain constant at room temperature.

Combine the solution in 11 different combinations. The procedure for each of the reactions isthe same



o Step 1: label 11 test tubes #111 with the corresponding amount of solution A (see the table below).

o Step 2: place 5.0 mL of solution B in each test tube and add 12 drops of starch o Step 3: label another 11 test tubes with the corresponding amount of solution C

(see the table below).o Step 4: add solution C into test tube containing solution A+B with the volume as

shown in the table below. Begin timing using stopwatch. Sir the solution with aclean stirring rod. At the first sign of color, stop timing. Record the results on thedata table.

o Step 5: CalculationsCalculate the initial concentrations of iodide and peroxydisulfate ion for each of themixtures.

For example: mixture 1

Iodide ion:( )(. /)

= 0.080 mol/L

Peroxydisulfate:( )(. /)

= 0.040 mol/L

o Step 6: Built the graphs Plot the concentration of iodide ion versus time for mixtures # 16. Time should

be on the X axis and the concentrations should be on the Y axis. Plot the concentration of peroxydisulfate ion versus time for mixtures # 1, 7, 8, 9,

10, and 11. Again, time should be on the X axis and the concentrations should beon the Y axis.

Number Solution A Solution B Solution C1 10.0 5.0 10.02 8.5 + 1.5 distilled water 5.0 10.03 7.0 + 3.0 distilled water 5.0 10.04 5.5 + 4.5 distilled water 5.0 10.05 4.0 + 6.0 distilled water 5.0 10.06 2.5 + 7.5 distilled water 5.0 10.07 10.0 5.0 8.5 + 1.5 distilled water8 10.0 5.0 7.0 + 3.0 distilled water9 10.0 5.0 5.5 + 4.5 distilled water

10 10.0 5.0 4.0 + 6.0 distilled water11 10.0 5.0 2.5 + 7.5 distilled water



PART 2: EFFECT OF TEMPERATURE ON THE REACTION RATE

The reaction rate for the oxidationreduction reaction between potassium permanganate,

KMnO4, and oxalic acid, H2C2O4, can be measured by observing the time elapsed for the purplecolor of the permanganate ion, MnO4 , to disappear.

5H2C2O4(aq) + 2KMnO4(aq) + 3H2SO4 2MnSO4(aq) + K2SO4(aq) + 10CO2(g) + 8H2O

Prepare the reaction system (use cylinder to get the chemicals):

o Place 1 mL of 0.01M KMnO4 and 5 mL of 3M H2SO4 into a clean test tube (3tubes)

o Place 5 mL of 0.33M H2C2O4 into a second, clean test tube. (3 tubes)o

Observe the reaction at room temperature:Pour the H2C2O4 solution into the KMnO4 solution. Observe and record the time for thepurple color of the permanganate ion to disappear.

Observe the reaction at high temperature:o Place a second KMnO4H2C2O4 pair of test tubes in warm water (500 C) bath

until thermal equilibrium is established. Pour the H2C2O4 solution into theKMnO4 solution, mix well and return the reaction system to the warm waterbath. Record the time for the purple color to disappear.

o Repeat the same procedure, but increase the temperature of the water bath toabout 900 C. Record the change

PART 3: EFFECT OF A CATALYST ON THE REACTION RATE

Hydrogen peroxide is relatively, but readily decomposes in the presence of a catalyst. In thispart, you will observe which reagent(s) act as a catalyst for the decomposition of hydrogenperoxide.

2H2O2 2H2O + O2

Label 7 test tubes # 17

Place 5 mL of the 3% H2O2 solution into each of the 8 test tubes. Add a pinch of each of the following reagents to separate test tubes:

MnCl2 MnO2 NaCl CaCl2 Zn KNO3 Fe(NO3)3 Mix well and observe the change with the production of gas bubbles. Record each reaction rate as fast, slow, very slow, or none in your data table.




5. LAB REPORT Present your observation and obtained data for each part Plot the graph in excel and prepare your report in Microsoft Word Present your conclusions




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