Chemistry lab manual

School of Biotechnology, International University-HCMC Department: Applied Department: Applied Chemistry

Laboratory Manual General Chemistry

LABORATORY MANUAL FOR “GENERALL CHEMI “GENERA CHEMISTRY STRY FOR ENGIN ENGINEERING EERING STUDENTS”

School of Biotechnology, International University-HCMC Department: Applied Department: Applied Chemistry


This laboratory is designed for internal use only. Its content is collected and composed by Dr Hoang 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 building”by ickie Williamson (Author), Larry Peck, (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] “Labor “Laborato atory ry Experi Experime ments nts for Gener General, al, Organi Organicc and Biochem Biochemist istry ry”by ”by Freder Frederick ick A. Bettel Bettelhei heim m (Brooks/Cole Laboratory Series for General Chemistry) 7th Edition, 2009. [4] “Experi “Experimen ments ts in Genera Generall Chemis Chemistry try”” by Bobby Bobby Stanto Stanton, n, Lin Zhu, Zhu, Charle Charless H. Atwood Atwood (Featu (Featurin ring g MeasureNet), 2nd Edition, 2009. [5] “Laboratory Manual for Principles of General Chemistry” by Jo Allan Beran, 9th Edition, 2010.

Common Laboratory Glassware and Equipment

electronic balance

buret clamp

beaker

Bunsen burner

crucible tongs

evaporating dish

crucible and lid

funnel

graduated cylinder

thermometer

clay triangle

Erlenmeyer flask

glass rod with rubber policeman

g n i n r a e L e g a g n e C / h t r o w s d a W / n o t n a t S y b b o B

ix

x

Common Laboratory Glassware and Equipment

ring stand and iron ring

scoopula

test tube clamp

test tube

watch glass

wire gauze

pipets

volumetric flask

test tube rack

utility clamp

buret

g n i n r a e L e g a g n e C / h t r o w s d a W / n o t n a t S y b b o B

Practice Safe Laboratory

A few precautions can make the laboratory experience relatively hazard free and safe. These experiments are on a small scale and thus many of the dangers found in the chemistry laboratory have been minimized. In addition to specific regulations that you may have for your laboratory, the following DO and DON’T RULES should be observed at all times. D O R U L E S

g n i n r a e L e g a g n e C , e l o C / s k o o r B 0 1 0 2 

T H G I R Y P O C

q

Do wear approved safety glasses or goggles at all times. The first thing you should do after you enter the laboratory is to put on your safety eyewear. The last thing you should do before you leave the laboratory is to remove them. Contact lens wearers must wear additional safety goggles. Prescription glasses are not suitable safety glasses; 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, no sandals). Be covered from the neck to the feet. Laboratory coats or aprons are recommended. Tie back long hair, out of the way of flames.

q

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

q

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

q

Do discard waste material properly. Organic chemical waste should be collected in appropriate waste containers and not flushed down sink drains. Dilute, nontoxic solutions may be washed down the sink with plenty of water. Insoluble and toxic waste chemicals should be collected in properly labeled waste containers. Follow the 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 the procedure carefully and be aware of the hazards before stepping foot into the laboratory.

xi

xii

Practice Safe Laboratory

D O N’ T R U L E S q

Do not eat or drink in the laboratory. Consume any food or drink before entering the laboratory. Chemicals could 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 flammable reagents could easily explode.

q

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

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 your instructor.

q

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

q

Do not enter the chemical storage area or remove chemicals from the supply area. Everyone must have access to the chemicals for the day’s experiment. Removal of a chemical from the storage or supply area only complicates 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 the immediate area.

q

Do not take unnecessary risks.

These DO and DON’T RULES for a safe laboratory are not an exhaustive list, but are a minimum list of precautions that will make the laboratory a safe and fun activity. Should you have any questions about a hazard, ask your instructor first—not your laboratory partner. Finally, if you wish to know about the dangers of any chemical you work with, read the Material Safety Data Sheet (MSDS). These sheets should be on file in the chemistry department office. A sample sheet is included here so you know what one looks like. This is the MSDS for glucose. Read it and see the kind of data included in there. Imagine all the additional cautions and precautions that the sheets would contain were you dealing with a chemical that is toxic or carcinogenic.

Name . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . Section . . . . . . . . . . . . . . .

Date . . . . . . . . . . . . . . .

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

Safety Quiz Indicate whether each of the following statements is true or false by writing the word TRUE or FALSE in the space provided. _____

1. If chemicals come into contact with your skin, immediately wash the affected area with

copious quantities of water. _____

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

chemicals. _____

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 will

not contaminate the entire reagent bottle. _____

5. When heating a liquid in a test tube, always point the test tube in a direction away from

any 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 from

all chemicals. _____ 9. I am not required to wear safety goggles while in the laboratory unless I am actually

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

xiii

School of Biotechnology, International University-HCMC Department: Applied Chemistry


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.

2.



To identify the products in these reactions and describe the chemical changes.



To write and balance the chemical equations for the reactions observed.

INTRODUCTION

Matter can undergo both physical and chemical changes. Chemical changes result in the formation of new substances. When a chemical reaction occurs, substances called reactants are

transformed

into different substances called products that often have different

appearances and different properties. In this experiment, you will perform nd observe a number 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 also learn how to classify the chemical reactions. One classification system involves five general types of reactions: synthesis, decomposition, single displacement, double displacement, and combustion.

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 NaOH

o

Test tube #2: add 10 drops of 2M NH4OH

o

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 KCl

Step 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 at least 02 minutes. Revision: October 1, 2013

Page 1



 Section 2: Reactions of KBr

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

Repeat 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 5

drops 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 5

drops of 2M H2SO4 and then add 5 drops of 3% H 2O2 solution. Record your observations.  Section 3: Place MnO2. Record

10 drops of 3% H2O2 solution into a test tube, then, add a “pinch” of solid

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 interface between 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 violet nitroso complex compound is then found by nitrogen monoxide with the excess iron (II) ions.  Section 2

Repeat Section 1 with NaNO2 instead of NaNO3 at Step 2. Record and compare the observations.  Section 3

Repeat Section 1 with acetic acid (CH 3COOH) instead of H2SO4 at Step 3. Record and compare 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.

Revision: October 1, 2013

Page 2





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 (K 2Cr2O7) 

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 FeSO 4 instead of FeCl 3 at Step 1. Record and compare the results.

8. REACTIONS OF Al3+ 

Step 1: Place 10 drops of 0.5M Al2(SO4)₃ 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 HCl

o

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 Red Red-orange


Wavelength (nm) 701 622 Page 3

School of Biotechnology, International University-HCMC Department: Applied Chemistry Orange Orange-yellow Yellow Yellow-green Green Green-blue Blue Blue-violet Violet


609 597 587 577 535 492 474 455 423

Wavelength values are given for mid-range of the color indicated. relationship between the wavelength, frequency and speed of an electromagnetic wave is given by the equation: C = λ  ν And the energy per photon (in Joules) is given by the equation: Ephoton = h  ν Where h is Planck’s 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


Page 4




Page 5



EXPERIMENT 1: LAB STRUCTURE


Page 6



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. INTRODUCTION Acids are proton donors and bases are proton acceptors. Acids primarily serve as sources of +

+

-

hydrogenions (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 H 3O and OH +

2H2O  H3O (aq) + OH

-

The extent of dissociation of water is very small; therefore pure water has no electrical conductivity. At the equilibrium, the ion product of water is only 1  10 +

-

Kw = [H 3O ][ OH ]= 10 +

-14

o

at 25 C.

-14

-

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

-

-7

[H3O ] = [ OH ] = 1 10 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 acidic the solution whereas the higher pH the more basic the solution. Strong acids and strong bases are completely dissociated in water to produce hydrogen ions or hydroxide ions respectively.


Page 7


Laboratory Manual General Chemistry +

Weak acids dissociate only partially and form little or very little H . This means that an equilibrium is established between the dissociated and un-dissociated forms:

HA(aq) ↔

௔=

H

ା + Aି

[Hା ][Aି ] [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 way comparable to pH, i.e. pKa = - lgKa. A buffer is a solution of a weak acid and its conjugate weak base. Buffers have the function +

-

-

that 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 of buffer solutions can be determined using the Henderson-Hasselbach equation:

[Aି ] pH = pK ୟ + log [HA]

൬

൰

A pH meter can be used to measure the pH of prepared solutions. Different classes of chemicals behave differently when dissolved in water. By doing this experiment, you will gain a better understanding of strong acids and strong bases, weak acids and weak bases, salts and buffers.

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.Stir the 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.


Page 8



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 beaker o 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 CH 3COONA and NH4Cl solution with the correct concentrationy by yourself.

3.5 BUFFERS Prepare 02 150 mL beakers each containing ~100 mL of 0.1M CH3COOH and 100 mL of 0.1M 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: Add 10 drops of 0.1M HCl. Measure the pH. Add enough to change the pH by one unit from the start. Record the o volume. Note: roughly 10 drops ~ 1 mL o



For part 2: o o

Add 10 drops of 0.1 M NaOH. Measure the pH. Add enough to change the pH by one unit from the start. Record the volume.


Page 9



 Section 3: 



Prepare 50 mL buffer B in a 100 mL beaker by mixing 40 mL of 0.1M CH3COOH and 10 mL of 0.1M CH3COONa Measure the pH of the buffer solution.

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

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


Page 10





Page 11


EXPERIMENT


3: OXIDATION/REDUCTION

TITRATION 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 to reactants, resulting in the changes of oxidation number. The substance that gains electrons is said to be reduced; therefore, it is called the oxidizing agent. The substance that loses electrons 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 6.02x1023 electrons and one gram equivalent weight of reducing agent is the weight that loses 6.02x1023 electrons. According to the definition of gram equivalent weight, one GEW of oxidizing agent reacts with one GEW of reducing agent:

GEWox = GEWred Consider the reaction of potassium permanganate (KMnO 4) with oxalic acid (H 2C2O4) in the presence of excess sulfuric acid (H 2SO4). The balanced molecular and net ionic equations are as follows, respectively.

2KMnO4 + 5H2C2O4 + 3H2SO4 → 10CO2 + K2SO4 + 2MnSO4 + 8H2O + 2+ 2MnO4 + 5H2C2O4 + 6H → 10CO2 + 2 Mn + 8H2O -

The oxidation number of Mn in MnO4 is +7 while it is +2 in Mn

2+

. Hence, each Mn undergoes a

change in oxidation number of five. Since each formula unit of KMnO4 contains one Mn, and each Mn gains five electrons, one mole of KMnO4 is five gram equivalent weights in this reaction. As a result, KMnO 4 produces 5 moles of electrons per mole of KMnO4 or has five equivalents per mole of KMnO4. Thus, the gram equivalent weight of KMnO4 in this reaction is

31.60 grams.

૚૞ૠ.૙ ૚ ૝ ૚ ૞ =


૜૚.૟૙

=

Page 12



The oxidation number of carbon in H2C2O4 is +3 while it is +4 in CO2. Thus each carbon undergoes a change in oxidation number of one. However, each formula unit of H2C2O4 contains two carbons, and since each carbon loses one electron, one mole of oxalic acid is two gram equivalent weights in this reaction. Consequently, H2C2O4 produces 2 moles of electrons per mole of oxalic acid or has two equivalents per mole of oxalic acid. The gram equivalent weight of H 2C2O4 is 45.0 grams

ૢ૙.૙ ૚ ૛ ૛ ૝ ૚ ૛ =

૝૞.૙

=

In this experiment, you will prepare an approximately 0.05N KMnO 4 solution and standardize this solution by titrating against a standard solution of H 2C2O4 (primary standard). Then the standardized KMnO4 solution (secondary standard) will be used to determine the concentration of unknown oxalic acid solution and unknown Fe

2+

solution. For redox titrations,

the number of equivalents of oxidizing agent must be equal to the number of equivalents of reducing agent. For the reaction of KMnO 4 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 normality of oxidizing or reducing agents. At the end of a titration, three of the four variables will be known and the unknown variable can 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 KMnO 4 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 KMnO 4 solution is dark color, read the buret at the top of the meniscus.


Page 13



3.2. STANDARDIZATION OF PREPARED KMNO4 SOLUTION: 

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



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).



o

o

Warm the flasks in the water bath 85 C – 90 C (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 experiment refers to the titrate volume needed to keep the faint pink color throughout the stirred solution for at least twenty seconds. Record the buret reading and calculate the normality of the 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). o o Warm the flasks in water bath 85 C ‐ 90 C. (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). o o Warm the flasks in water bath 85 C ‐ 90 C. Titrate the hot solutions. Calculate the normality of the unknown concentration FeSO4 solution; determine the average and the standard deviation.



Page 14


7.


LAB REPORT 




Present your explanation and conclusions


Page 15





Page 16


EXPERIMENT


4: CHEMICAL EQUILIBRIUM

1. OBJECTIVES To observe the effect of applying stresses on chemical systems at equilibrium. To apply Le Chatelier’s 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 to the rate of the backward reaction. Reversible reaction:

A reversible reaction at equilibrium can be disturbed if a stress is applied to it. Stresses can be changes in concentration, temperature or pressure. The composition of the reaction mixture will shift until equilibrium has been reestablished. This is known as Le Chatelier’s Principle. In this experiment, the effect of applying stresses to a variety of chemical systems at equilibrium will be observed and we also see if the results are consistent with Le Chatelier’s 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:

2CrO4

2-

+

2-

+ 2H (aq) ↔ Cr2O7

+ 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, mix well. Divide the solution evenly into two test tubes. Save one as a reference. Note the color .  Test tube #1 (reference): add nothing 

Test tube #2: o

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

o

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


Page 17



Addition #3: again add the 6M HCl solution drop wise until further addition

o

results 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 0.01 M 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 ≈ 10‐12



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: 3+

-

2+

Fe (aq) + SCN (aq) ↔ [Fe(SCN)] (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 to test 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.


Ca (aq)

+

2-

C2O4 (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 drop wise until the solution turns a purple-violet color. If the system turns a deep


Page 18



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 room temperature. Compare to the control.


[Co(H2O)6] (aq)

-

+ 4Cl (aq)

4.

CHEMICAL AND EQUIPMENTS

8.

LAB REPORT

2-

↔ [CoCl4] (aq)



Present your observation and obtained data for each part.



Present your explanation and conclusions.


+ 6H2O(l)

Page 19





Page 20



EXPERIMENT 5: REACTION RATE 1.

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 chemical reaction is affected by a number of factors including temperature of the reaction, the nature of the reactants, concentration of the reactants, the surface area of the reactants, the presence of a catalyst and the pressure the reaction is under. The greater the rate of a chemical reaction, the less time is needed for a specific amount of reactants to be converted to products. The rate of a reaction can be determined one of two ways; either measure the time it 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: -

2-

2I + S2O8

Reaction 1:

→ I2 +

2‐

2SO4

Iodide ions + peroxydisulfate ions → iodine + sulfate ions

I2

Reaction 2:

+

2S2O3

Iodine + thiosulfate ion

2‐

→

‐

2

2I + S4O6 ‐

→ iodide ion + tetrathionate ion

Reaction 1 is relatively slow. As the iodine is formed it is quickly used in reaction 2, which is relatively fast. The limiting reaction (solution B) is a source of the thiosulfate ions. When solution B is used up, the excess iodine formed will react with starch to form a deep blue solution. In this experiment, you will vary the concentrations of solutions A and C. The temperature will remain constant at room temperature. Combine the solution in 11 different combinations. The procedure for each of the reactions is the same


Page 21

School of Biotechnology, International University-HCMC Department: Applied Chemistry o


Step 1: label 11 test tubes #1‐11 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 1‐2 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 a clean stirring rod. At the first sign of color, stop timing. Record the results on the data table.

o

Step 5: Calculations Calculate the initial concentrations of iodide and peroxydisulfate ion for each of the mixtures. For example: mixture 1

ଵ଴୫୐)×(଴.ଶ଴୫୭୪/୐) = 0.080 mol/L ଶହ୫୐ (ଵ଴ ୫୐)×(଴.ଵ଴୫୭୪/୐) = 0.040 mol/L Peroxydisulfate: ଶହ୫୐ Iodide ion:

o

(

Step 6: Built the graphs ‐ Plot the concentration of iodide ion versus time for mixtures # 1‐6. 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 be on the Y – axis. Number

Solution A

Solution B

Solution C

1

10.0

5.0

10.0

2

8.5 + 1.5 distilled water

5.0

10.0

3


5.0

10.0

4


5.0

10.0

5


5.0

10.0

6


5.0

10.0

7

10.0

5.0


8

10.0

5.0


9

10.0

5.0


10

10.0

5.0


11

10.0

5.0



Page 22



PART 2: EFFECT OF TEMPERATURE ON THE REACTION RATE The reaction rate for the oxidation‐reduction reaction between potassium permanganate, KMnO4, and oxalic acid, H2C2O4, can be measured by observing the time elapsed for the purple ‐

color of the permanganate ion, MnO 4 , 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 (3 tubes)

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 the purple color of the permanganate ion to disappear. 

Observe the reaction at high temperature: o

Place a second KMnO4‐H2C2O4 pair of test tubes in warm water (500 C) bath until thermal equilibrium is established. Pour the H 2C2O4 solution into the KMnO4 solution, mix well and return the reaction system to the warm water bath. Record the time for the purple color to disappear.

o

Repeat the same procedure, but increase the temperature of the water bath to about 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 this part, you will observe which reagent(s) act as a catalyst for the decomposition of hydrogen peroxide.

2H2O2 → 2H2O + O2 

Label 7 test tubes # 1‐7



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.


Page 23


4.

5.


CHEMICALS AND EQUIPMENTS

LAB REPORT 




Plot the graph in excel and prepare your report in Microsoft Word



Present your conclusions


Page 24

Chemistry lab manual

Recommend Documents