STK1084 Lab Manual

STK1084 INORGANIC CHEMISTRY I LABORATORY MANUAL SESSION 2011/2012

CHEMISTRY DEPARTMENT FACULTY OF RESOURCE SCIENCE & TECHNOLOGY UNIMAS

STK1084 Inorganic Chemistry I, Laboratory manual

IMPORTANT LAB SAFETY The chemistry laboratory can be a place of discovery and learning. However, by the very nature of laboratory work, it can be a place of danger if proper common-sense precautions aren't taken. While every effort has been made to eliminate the use of explosive, highly toxic, and carcinogenic substances from the experiments which you will perform, there is a certain unavoidable hazard associated with the use of a variety of chemicals and glassware. You are expected to learn and adhere to the following general safety guidelines to ensure a safe laboratory environment for both yourself and the people you may be working near. Additional safety precautions will be announced in class prior to experiments where a potential danger exists. Students who fail to follow all safety rules may be asked to leave the lab or suffer grading penalties. Attire 1. Safety goggles must be worn at all times while in the laboratory. This rule must be followed whether you are actually working on an experiment or simply writing in your lab notebook. You must wear safety goggles provided by the chemistry department. 2. Contact lenses are not allowed. Even when worn under safety goggles, various fumes may accumulate under the lens and cause serious injuries or blindness. 3. Closed toe shoes and long pants must be worn in the lab. Sandals and shorts are not allowed. 4. Long hair must be tied back when using open flames. Conduct 4. Eating, drinking, and smoking are strictly prohibited in the laboratory. 5. No unauthorized experiments are to be performed. If you are curious about trying a procedure not covered in the experimental procedure, consult with your laboratory instructor. 6. Never taste anything. Never directly smell the source of any vapor or gas; instead by means of your cupped hand, waft a small sample to your nose. Do not inhale these vapors but take in only enough to detect an odor if one exists. 7. Coats, backpacks, etc., should not be left on the lab benches and stools. There is a hook rack along the back wall at either end of the lab. There are coat racks just inside the each entrance to the balance room at the back of the lab. Beware that lab chemicals can destroy personal possessions. 8. Always wash your hands before leaving lab. 9. Learn where the safety and first-aid equipment is located. This includes fire extinguishers, fire blankets, and eye-wash stations. 10. Notify the instructor immediately in case of an accident.

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Proper Handling of Chemicals and Equipment 11. Consider all chemicals to be hazardous unless you are instructed otherwise. Material Safety Data Sheets (MSDS) are available in lab for all chemicals in use. These will inform you of any hazards and precautions of which you should be aware. 12. Know what chemicals you are using. Carefully read the label twice before taking anything from a bottle. 13. Excess reagents are never to be returned to stock bottles. If you take too much, dispose of the excess. 14. Many common reagents, for example, alcohols and acetone, are highly flammable. Do not use them anywhere near open flames. 15. Always pour acids into water. If you pour water into acid, the heat of reaction will cause the water to explode into steam, sometimes violently, and the acid will splatter. 16. If chemicals come into contact with your skin or eyes, flush immediately with copious amounts of water and consult with your instructor. 17. Never point a test tube or any vessel that you are heating at yourself or your neighbor--it may erupt like a geyser. 18. Dispose of chemicals properly. Waste containers will be provided and their use will be explained by your demonstrators. Unless you are explicitly told otherwise, assume that only water may be put in the lab sinks. 19. Clean up all broken glassware immediately and dispose of the broken glass properly. 20. Contact the stockroom for clean-up of mercury spills. 21. Never leave burners unattended. Turn them off whenever you leave your workstation. Be sure that the gas is shut off at the bench rack when you leave the lab. 22. Beware of hot glass as it looks exactly like cold glass.

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EXPERIMENT 1: INTRODUCTION TO ELEMENTS INTRODUCTION Periodic table is the only table to arrange the elements available in the world according to their own atomic number. Chemist can use the data contained in this table to predict the properties of each element. Elements in the world can be classified through three (3) methods: 1. Periodic table which classified the elements following the periodic law 2. Metallic, half metallic or non metallic and 3. Following the reactivity in electrochemistry series Based on the periodic table, we know that: a) Similar elements will be categorized in a same group b) The properties of each element will change slowly if its electronegativity changes from one element to another element c) The unsimilarity of each elements can be easily to differentiate Generally, elements also can be categorized based on their electronic configuration, namely: a) Noble gases (Group 0) b) Block s and p elements (Group IA, IIA, IIIB, IVB, VB, VIB and VIIB) c) Transition metals or block d elements (Group IIIA, IVA,VA, VIA, VIIA, VIII, IIB and IB) d) Lanthanides and actinides or block f elements

REAGENTS AND METHOD Reagents Na, K, Li, Mg, Ca, Al metals, CaCl2, NaCl, KCl, LiCl, MgCl2, BaCl2, BeSO4, MgSO4, H3BO3, AlCl3, SrCl2, bromothymol blue indicator, phenolphthalein, litmus paper universal, HCI, HNO3, H2SO4 acid, NaOH, ammonia solution, toluene. Apparatus Test tube, burette, dropper, Bunsen burner, pin, 500 mL conical flask, filter paper, watch glass, water bath, taper, platinum wire.

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Methods A. Group I elements Cut a small piece of Na (near to pea size) and dry it on two pieces of filter paper. Add the Na to the test tube, which contents 10 mL of distilled water and few drops of phenolphthalein. Close the test tube with its lid. Check the gas release from the reaction by doing a flame test with a piece of coconut taper. Repeat the same procedures for K and Li. Prepare two test tubes, which contain 5 mL of toluene. Add a small piece of K and Li into the respective test tubes. Place both test tubes into the water bath and observe the reactivity of both elements. Run the flame test on the NaCl, LiCl and KCl using the platinum wire. Observe the flame colour of each element. B. Group II elements Place a piece of Na, Mg and Ca on a filter paper, respectively. Test the hardness of each element using a pin. Run the flame test on the MgCl2, SrCl2 and BaCl2 using the platinum wire. Observe the flame colour of each element. Fill 1 mL of BeSO4 solution into two respective test tubes and drop 0.1 M NaOH solution into the test tube until precipitate formed. Fill in the first test tube with excess 0.1 M HCl solution and excess 0.1 M NaOH solution for the second one. Repeat the same procedures for the MgSO4 solution. Observe the solubility of each compound. Prepare 3 large test tubes and add 10 mL distilled water into each of them. Prepare 3x10 cm shining magnesium ribbons. Use sand papers to remove the oxidized surface on the magnesium ribbons. Add in the magnesium ribbons into each test tube. Test tube 1 works as a control. Burn test tube 2 hardly. Add few drops of 0.1 M HCl into test tube 3. Compare the rate of reaction for each test tube. Add 5 mL of distilled water and 5 drops of phenolphthalein into a test tube. Add a small piece of Mg (pea size) into the test tube. Close the test tube with its lid.

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Check the gas releases from the reaction in the test tube through flame test. Repeat the same test for Ca. Prepare four test tubes and add 5 mL, 0.1 M of MgCl2, SrCl2 and BaCl2 solution, respectively. Drop in 0.1 M NaOH solution into each test tube using burette until the precipitate is formed. Record the volume of NaOH used for each test tube. C. Group III elements Add 1 g of boric acid (H3BO3) into a test tube, which contains 5 mL of distilled water and test it by universal indicator. Add 5 mL of aluminium chloride (AlCl3) solution in the test tube and drop in concentrated NH3 until precipitate formed. Divide the solution into two portions in two different test tubes. Drop in 0.1 M HCl into first test tube and 0.1 M NaOH into second test tube. Observe the solubility of the precipitate. Place a few pieces of Al into a test tube, which contain 0.1 M HCl acid. Repeat the same thing but replace HCl with 0.1 M NaOH. Observe the differences between HCl and NaOH. Add few pieces of Al into distilled water. Observe the changes happen. Take few pieces of Al and wipe it with cotton which has been soaked in HgCl2 solution. Put this Al into the distilled water. Observe the changes happen. Precaution: 1) Wear protective goggle when conducting this experiment 2) Do not throw the waste from Na reaction into the sink. Throw it into the beaker filled with paraffin oil. RESULTS A) GROUP I ELEMENTS Test Observation and reaction Reactivity

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Chloride flame test

B) GROUP II ELEMENTS Test Observation and reaction Hardness

Solubility of chloride precipitate

Reactivity

Chloride flame test

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C) GROUP III ELEMENTS Test Observation and reaction Reactivity with water

Precipitate test

POST-LAB QUESTIONS 1) What is the apparatus that you use to cut Na? Why do we have to use this apparatus to cut Na? 2) Metallic oxide like Na2O and BaO can be dissolved in the water and cause the litmus paper change to red. What is the property of these substances that cause the changes? 3) Estimate the reactivity of Si in the water if compared to Na, Mg and Al. Explain.

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EXPERIMENT 2: QUALITATIVE ANALYSIS OF CATIONS

INTRODUCTION Qualitative analysis is defined as the determination of elements, which presence in a substance, mixture of solution. Qualitative analysis can be said as a pre-step for qualitative analysis in the determination of elements in a substance. Therefore, qualitative analysis is so important in chemistry and it requires a specific knowledge to understand the substances characteristics and its reactions in solution react with various types of reagents. The scheme of classical cation qualitative analysis can be divided into five groups. Group I contains the cations where their chloride salts do not dissolve in diluted HCl acid. Group II and III will form sulfide precipitate when they are react with H2S acid and Group III named as sulfide acid group. Group IV include those cations, where their carbonate salts do not dissolve in water. Finally, Group V will be for those cations where their salts do not dissolve in water. In this experiment, a solution may contain one or many types of metallic ion. The first step will be the separation of metallic ion into different group and each of them will show different characteristics, which are the basic of the separation. REAGENTS AND METHODS Youll be given 6 known solution and 2 unknown samples in this experiment. Reagents AgNO3 0.2 M, CuNO3 0.2 M, Fe(NO3)3 0.2 M, Ni(NO3)2 0.2 M, BaCl2 0.2 M, Ag(NH3)2Cl 0.2 M, concentrated HCl acid, HCl acid 1 M, H2SO4 acid 3 M, KSCN solution 0.2 M, NH3 solution 4 M, DMG, concentrated HNO3 acid, litmus paper. Apparatus Test tube, platinum wire, Bunsen burner, conical flask, beaker, petri dish, measuring cylinder (500 mL), burette, centrifuge.

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Method A) Cation Test 1) Barium (Ba2+) Add 5 drops of Ba2+ (BaCl2 0.2 M) followed by 5 drops of distilled water and 1 drop of HCl into a test tube. Observe the reaction occurs. After that, add 3 drops of H2SO4 acid 3 M and centrifuge the test tube for 1 minute. Add one more drop of H2SO4 to form the precipitate and continue to centrifuge. Record the colour of the precipitate. 2) Silver (Ag+) Add 10 drops of Ag+ solution (AgNO3 0.2 M) and followed by distilled water until it just enough 1 mL. Add 2 drops of HCl acid 6 M. centrifuge and separate the precipitate from the solution. Add 5 drops of NH3 solution 4 M and finally add HNO3 4 M (stir at the same time) until it become an acidic solution. 3) Copper (Cu2+) Add 5 drops of Cu2+ solution [Cu(NO3)2, 0.2 M] followed by 5 drops of distilled water and 1 drop of HCl acid 1 M and 1 drop of H2SO4 3 M into the test tube. Subsequently, add NaOH solution 2 M until the solution change to base solution. Shake the test tube and add one more drop of NaOH solution to complete the precipitation. Add 1 mL of NH3 3 M into the test tube and shake it (if necessary, add more NH3 solution to complete the precipitation) Add acetic acid 3 M drop by drop until the solution become colourless. 4) Iron (Fe3+) Add 5 drops of Fe3+ [Fe(NO3)3, 0.2 M] and followed by 5 drops of distilled water, 1 drop of HCl acid 1 M and 1 drop of H2SO4 3 M into a test tube. After that, add NaOH 2 M drop by drop until the solution change to basic solution (test with litmus paper). Shake the test tube. Add one more drop of NaOH to complete the precipitation. Dissolve the precipitation into 1 mL of NH3 and shake it. Add 1 mL of HNO3 6 M and finally add 5 drops of KSCN 0.2 M. 5) Nickel (Ni2+) Add 5 drops of Ni2+ [Ni(NO3)2 0.2 M] and followed by 5 drops of distilled water and followed by 1 drop of HCl acid 1 M and H2SO4 acid 3 M into a test tube. Add NaOH 2 M drop by drop until the solution change to basic solution (test by litmus paper). Centrifuge the test tube and add one more drop of NaOH to complete the precipitation. Add 1 mL of NH3 3 M solution and stir it. Add 1 more drop of NH3 solution to complete the precipitation. Add aceticSemester I, 2011 - 2012 Page 9


acid 3 M solution drop by drop until the solution become colourless. Finally add 2 pieces of dimethylglyoxime (DMG). 6) Lead (Pb2+) Add 10 drops of Pb2+ solution [Pb(NO3)2 0.2 M] into a test tube and add distilled water until it just enough to 1 mL. Add 2 drops of HCl acid 6 M. Centrifuge and separate the precipitate from the solution. Dissolve the precipitate with 10 drops of boil water and add 2 drops of K2Cr2O4 1 M into the solution. B) Cations Test in Unknown Solution For unknown solution, you need to determine the cations that contain in the solution. Show the result for each test and detect the cations that presence in the unknown solution. RESULTS A) CATION TEST Cation Observation 2+ Ba

Chemical equations

Ag+

Cu2+

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Cation Fe3+

Observation

Chemical equations

Ni2+

Pb2+

B) CATION TEST FOR UNKNOWN SOLUTION Test Observation Conclusion

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POST-LAB QUESTIONS 1) A solution containing Ag+ or Pb2+ is mixed with a chloride aqueous solution and a white precipitate is obtained. Why HCl acid is selected as the precipitation agent but neither CuCl2 nor KCl? 2) What kind of reagent that can be used to separate Cu2+ from a Bi3+ solution? 3) Explain the problems that may occur if we use flame test to determine cation from a solution that contains mixtures of cations.

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EXPERIMENT 3: QUALITATIVE ANALYSIS OF ANIONS INTRODUCTION In the previous experiment, the qualitative analysis of cations has been carried out. The cations can be divided into several groups based on their chemical characteristics. By using the continuous separation technique, all the cations content in the mixture can be determined one by one. Same procedures will be used in the determination of anions in a mixture in this experiment. One complete set of analysis scheme for anions will be consuming a lot of time. Therefore, usually the scientist will run the test on the salt where we know its anion. The method that always be used include acid-base reaction, precipitation, complexation and oxidation-reduction reaction.

REAGENTS AND METHOD Reagents Na2SO4 0.5 M, BaCl2 1 M, H2SO4 4 M, NaNO3 0.5 M, FeSO4 0.1 M, NaCl 0.5 M, NaI 0.5 M, AgNO3 0.05 M, NH3 solution 0.2 M, Fe(NO3)3 0.1 M, Na2CO3 0.5 M. Apparatus Test tube, conical flask, beaker, petri dish, measuring cylinder (500 mL), burette. Method Run the test for each anion solution following the procedures described as below. If there is no reaction during the testing, stir the mixture with a piece of rod glass. A. Anion Test 1) Sulfate (SO42-) Mix 5 drops of SO42- (Na2SO4 0.5 M) with 5 drops of HCl acid 6 M, add another 2 drops of SO42- until it excess. Heat the mixture in steam bath until excess CO2 release from it. Add few drops of BaCl2 1 M to complete the precipitation and finally leave the mixture for a while. 2) Nitrate (NO3-) Mix H2SO4 4 M into 5 drops of NO3- solution (NaNO3 0.5 M) until the solution become acidic (check with litmus paper). Add 5 drops of FeSO4 0.1 M and 5 drops of concentrated H2SO4. When adding in the concentratedSemester I, 2011 - 2012 Page 13


H2SO4, make sure the test tube in the position 45, so that the sulfuric acid can flow through the wall of test tube and form a different layer at the bottom part of the test tube. 3) Chloride (Cl-) Add 5 drops of Cl- (NaCl 0.5 M) and AgNO3 0.05 M into a test tube. Check the solubility of precipitate formed in NH3 2 M solution. 4) Iodide (I-) Dilute 6 drops of I- solution (NaI 0.5 M) by 12 drops of distilled water. Add HNO3 acid 4 M until the solution becomes acidic and add 2 drops of HNO3 for excess. Mix the solution with 1 mL of Fe(NO3)3 as well as 10 drops of dichloromethane and shake it. 5) Carbonate (CO32-) Add 1 mL of HCl acid 6 M (carefully) into a test tube, which contains 10 mL of CO32- solution (Na2CO3 0.5 M). Identify the gas releases (using flame test) from the test tube when adding the HCl. Finally, heat the test tube in the steam bath to increase the bubble release. B. Unknown Sample Test After finish all the tests above, get 2 unknown samples from your instructor and analyze the samples following the procedures above. Perhaps, the unknown samples contain more than one anion. Show all the results for each test and estimate which anions presence in the unknown solution. RESULTS A) ANION TEST Anion Observation 2Sulfate (SO4 ) Chemical equations

Nitrate (NO3-)

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Chloride (Cl-)

Iodide (I-)

Carbonate (CO32-)

B) ANION TEST FOR UNKNOWN SOLUTION Test Observation Conclusion

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POST-LAB QUESTIONS 1) A student has dissolved few grams of unknown salt into the distilled water. When this solution is added with NaOH, there is no precipitate formed and litmus paper change to blue in colour. This student finally concluded that the salt that he dissolved is ammonium salt. So the question is why the method that he used to determine the salt is not correct? 2) What are the reagents that can be used to differentiate CaCl2 and CaCO3 since both of them are white in colour?

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EXPERIMENT 4: DETERMINATION OF THE FORMULA OF A HYDRATED SALT COMPLEX INTRODUCTION Most of the hydrated salts can be determined by heating the salt for 5 - 10 minutes. The water contained in the salt will be removed during the heating process and remain a dehydrated salt. The formula of hydrated salt can be determined from the difference in the weight of hydrated salt with dehydrated salt. The right techniques applied in this experiment are very important to obtain the accurate result. All the weight should be recorded in at least 0.0001 g. Sulfate composition percentage in sulfate-hydrated salts can also be determined by changing the salt to barium sulfate. Barium sulfate forms an insoluble precipitate and can be easily filtered out. The percentage of sulfate in the salt can be determined by obtaining the net weight of dried BaSO4. REAGENTS AND METHOD Reagents Na2SO4, HCl acid 0.5 M, BaCl2, ethanol Apparatus Hotplate, crucible, analytical balance, filter paper Method A) Determination of water content in hydrated salt Wash two pieces of crucible thoroughly. Heat the crucible above a hotplate for 5 minutes. Cool it to the room temperature and weigh it. Record the crucibles weight. Heat it again for 5 minutes, cool it and weigh it again. Repeat this step until the weight difference not more than 0.0004 g. Add 1 - 2 g of Na2SO4.nH2O salt into the crucible and weigh it accurately. Heat the crucible with its content on the hotplate for 10 minutes. Cool down the crucible to room temperature and weigh it accurately. Repeat the same

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above processes again until the difference between the first and second reading not more than 0.0004 g. Calculate the percentage of the water and n value in the samples formula. B) Determination of sulfate content in sulfate-hydrated salt Weigh about 1 g of sample sufate-hydrated salt in a 400 mL beaker. Dissolved the salt into 50 mL of distilled water and add 5 drops of HCl acid 0.5 M. Heat this solution until it boils. Add 100 mL of BaCl2 0.1 M into this solution until the entire sulfate precipitated. Heat the precipitate slowly for 5 minutes. Filter out the precipitate using filter paper. Wash the precipitate few times with distilled water to remove the soluble particles. Place the precipitate with filter paper into a crucible (make sure this crucible has been weighed accurately before used). Heat the precipitate in the crucible on a hotplate until it turns to white in colour. Leave the precipitate to cool down to room temperature. Add 2 drops of the mixture of 1 mL of ethanol. Heat the precipitate again and cool it down to room temperature. Weigh the dry BaSO4. Repeat the heating, cooling and weighing processes few times until the weight of BaSO4 not much different (0.0004 g). Calculate the percentage of the sulfate in sample salt. RESULTS A) Determination of water content in hydrated salt 1st reading i ii iii iv v Weight of crucible, g Weight of crucible + hydrated salt, g Weight of crucible + dehydrated salt,g Weight of water (ii iii) Water content in hydrated salt, % [ iv/(ii-i)] x 100% vi Average of water content in hydrated salt, % vii Mole of dehydrated salt (Na2SO4) viii Mole of water ix Formula of the hydrated salt 2nd reading 3rd reading

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B) Determination of sulfate content in sulfate-hydrated salt. 1st reading i ii iii Weight of crucible, g Weight of sulfate-hydrated salt, g Weight of crucible + BaSO4 precipitate, g iv Weight of BaSO4 (iii i) vii Mole of BaSO4 viii Percentage of sulfate in the hydrated salt, % POST-LAB QUESTIONS 1) If the heating process is not really completed, what is the effect on the percentage of water containing in the hydrated salt. 2) Let say 1.5 g of Na2SO4.10H2O is heated for 10 minutes, what is the balance weight that supposed to be obtained after 10 minutes heating? 3) Why Na2SO4 hydrated need to be converted to BaSO4? 2nd reading 3rd reading

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EXPERIMENT 5: DETERMINATION OF ALUM FROM ALUMINIUM CANS INTRODUCTION Many soft drinks can are mainly made by aluminium (Al). The percentage of Al content is about 98 - 99%, which is more than Mg. The relative abundance of Al in nature is 8.1%, which is the highest abundance of metallic in the world. Because of its low density, stronger stretch and tolerable to erosion, it is always be used as the material in the manufacturing of airplanes, vehicles and aluminium cans. Recycle of tin is one of the positive contributions to save the environment and energy resources. In the recycle process, Al will be melted and reused to make another types of Al products. The importance of this process is that it can be used to prepare alum potassium, one kind of hydrated salts with the formula of KAl(SO4)2.12H2O. This alum is always being used in wastewater treatment and also in fire extinguisher. Al metallic can react with hot solution of KOH and form aluminate salt, KAl(OH)4 in aqueous form. 2 Al (s) + 2 KOH (aq) + 6 H2O (l) 2 KAl(OH)4 (aq) + 3 H2 (g) When KAl(OH)4 solution react with H2SO4 acid, the solid Al(OH)3 will be formed and it will dissolve in excess hot H2SO4 acid. 2 KAl(OH)4 (aq) + H2SO4 (aq) 2 Al(OH)3 (s) + K2SO4 (aq) + 2 H2O (l) 2 Al(OH)3 (s) + 3 H2SO4 (aq) Al2(SO4)3 (aq) + 6 H2O (l) Alum potassium will be produced after this concentrated solution is cooled down. K2SO4 (aq) + Al2(SO4)3 (aq) + 12 H2O 2 KAl(SO4)212 H2O (s)

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REAGENTS AND METHOD Reagents Aluminium cans, KOH solution 1.5 M, H2SO4 acid 6 M, ethanol Apparatus Analytical balance, hotplate, thermometer, retort stand, Buchner filter, filter paper, beaker (250 mL), glass rod, melting point capillary tube, steam bath, rubber band, petri dish. Method A. Preparation of alum from aluminium can Weigh about 2 g of Al can (already been cut into small pieces) and put it into a 250 mL beaker. Record the weight of the Al can to the nearest 0.01 g. Add 60 mL of KOH solution 1.5 M into the beaker to react with the Al. Heat the mixture slowly on the hotplate until the Al totally disappear. (Note: H2 gas will be released in this step, do it in fume chamber). After the entire Al fully reacted, filter the hot solution slowly into another 150 mL beaker using the cotton in the filter tunnel. Let the solution cool down. Add slowly H2SO4 acid 6 M into the beaker until Al(OH)3 formed (stir it at the same time). Heat the mixture slowly until all the Al(OH)3 dissolve. Leave the hot solution cool down in the cool water bath. After about 20 minutes, the crystal of alum will appear. If not, stir the solution slowly until crystal is formed. Weigh a piece of filter paper and filter out the crystal from its solution using Buchner filter. Rinse the crystal with ethanol twice. Place the crystal and filter paper on the petri dish and dry it. Weigh the crystal + filter paper and get the weight of the crystal. B. Melting point of alum Add a small piece of alum into a melting point capillary tube. Place the capillary tube on a thermometer using a rubber band and immerse it into water bath. Heat the water slowly, and observe the condition of solid (any phase change or not) inside the capillary tube. Once the crystal starts melting, record the temperature.

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RESULTS A. Preparation of alum from aluminium cans Weight of aluminium cans (in piece) Weight of filter paper + KAl(SO4)2.12 H2O Weight of filter paper Weight of KAl(SO4)2.12 H2O Moles of Al cans used Moles of alum formed (experimental) Moles of theoretical alum Yield of KAl(SO4)2.12 H2O (%) B. Melting point Melting point

POST-LAB QUESTIONS 1. Why we use ethanol to rinse the crystal instead of distilled water? 2. Better yield of crystal can be obtained if the alum solution is placed into refrigerator for a few days. Explain why. 3. Solubility of alum, KAl(SO4)212 H2O in an aqueous solution at 0 C is 114 g/L. In this experiment, calculate the weight of alum remain in the solution after crystallization completed. (Volume = 40 mL)

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EXPERIMENT 6: QUANTITATIVE ANALYSIS BY USING ETHYLENEDIAMINETETRAACETIC ACID (EDTA) INTRODUCTION Complexometry titration is a stoichiometry titration which involves the formation of complex when the titrant is added into the sample that we want to analyze. In order to make sure the reaction is completed, the reactions must hit all the requirements such as faster rate, follow to the reaction equation and the end point. The best reagent used in the complexometry titration is ethylenediaminetetraacetic acid (EDTA). There are some reasons why EDTA is said to be the best reagent in complexometry titration: a) EDTA can form the complex faster with stoichiometry 1:1. At the same time, it is stable and able to dissolve most of the metallic ions. b) Equivalent point can be detected easily. c) This titration is suitable in wide range of concentration. d) The disodium salt from EDTA (Na2H2Y. 2H2O) can be used as the primary standardize chemical after been dried under temperature 80 C. The solution from disodium salt obtains pH 4 to 5. If Na2H2Y is used, the complexation reaction with a metallic ion can be represented as below: Mn+ + H2Y- [(MY)n-4] + 2 H+ Where the formation constant K, can be written as below: K = [(MY)n-4] [H+]2 / [Mn+] [H2Y2-] Therefore, this reaction is very sensitive to pH (due to [H+] in the formula), and all the procedures where EDTA is used as the titrant, must involve a suitable buffer solution to control the concentration of H+ during the titration. The formula of EDTA,HOOCH2C N HOOCH2CSemester I, 2011 - 2012

CH2COOH CH2 CH2 N CH2COOHPage 23


This formula always be summarized into H4Y, where Y can represent as,-

OOCH2C N CH2 CH2 N

CH2COO-

-

OOCH2C

CH2COO-

With this formulation, the molecular weight for (Na2H2Y.2 H2O) is 372.25 g/mol e) The metallic ions can be classified into three groups according to their stability constant. In general, metallic ion in Group I can be titrated in base condition (pH 8 10), the metallic ion in Group II will be in acidic condition to slightly base condition (pH 4 7) and finally the metallic ion in Group III will be in acid condition (pH 1 4). Table 6.1 lists out the metallic ions, which are classified into Group I, II and III. Table 6.1 Metallic ions in Group I, II and III Group Metallic ions I Mg2+ , Ca2+ , Sr2+ , Ba2+ II Mn2+, Fe2+, Co2+, Ni2+, Cu2+, Zn2+, Cd2+, Al3+, Pb2+, TiO2+, V2+, VO2+ and lanthanoid ions III Hg2+, Bi3+, Co3+, Fe3+, La3+, In3+, Sc3+, Ti2+, Ti3+, V3+, Th4+ Determination of end point: The most effective way to determine the end point of complexometry titration is using the metallochromic indicator. Metallochromic indicator is a substance that acts as a complexation agent for the metallic ion. In an optimum condition, the metallic complexindicator, which is a metallic complex link with indicator, will obviously show different colour than the uncomplexed indicator. Stability constant of metallic complex indicator is necessary in this titration. If the stability constant is too large, it means that the sample is over titration, whereas if the stability constant is too small, it means that the titration is not complete. There are several types of indicator that can be used as the metallochromic indicator and eriochrome black T (EBT) is one of them. The chemical name of EBT is 1-(1-hydroxy-2naphthylazo)-6-nitro-2-naphtol-4-sulfonate (II), the short form of this chemical is H3In. This indicator has three acid sites and two of them involve in the changes of the colour. Figure 6.1 shows the colour changes of EBT.

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HO OH N N

SO3H

H3InN2O

- H+OH

+H+HO SO3-

(H2In)Red

N N N2O

pH 6.3 - H+ON N N2O

+H+HO SO3-

(HIn)2Blue

pH 11.5 - H+OH N N N2O

+H+-

O

SO3-

In3Orange

Figure 6.1 Structural and colour change of EBT according to pH.

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With the presence of Mg2+,

When pH 10, Mg2+ will form a complex with EBT indicator, therefore the optimum pH for the titration of Mg2+ with EDTA is about 9 11. During the equivalent point, the colour 9-11. changes from dark red to sky blue when the last Mg2+ release from (MgIn)- complex by EDTA. However, EBT is not suitable to determine Ca2+ because the complex formed between EBT and Ca2+ is not as stable as Mg2+. REAGENTS AND METHOD Reagents MgCl26H2O powder, EDTA solution 0.01 M, EBT indicator, NaOH solution 1.0 M, nitric acid 0.1 M, buffer solution pH 10, sodium sulfide 1%, pyrocatechol indicator, unknown Bi3+ solution, NH4OH 0.1 M. Apparatus Beaker 100 mL, measuring cylinder 25 mL, glass rod, volumetric flask 100 mL, conical flask 250 mL, pH paper Method A. Preparation of standard solution of Mg2+ Weigh accurately 0.203 g of MgCl26H2O in a 100 mL beaker. Add about 50 mL . of distilled water into that beaker and dissolve the MgCl26H2O contain in the beaker. Transfer the solution into a volumetric flask 100 mL. Top up the solution with distilled water until to the mark of the flask. Shake the volumetric flask thoroughly to mix the solution completely. Use 25 mL of this solution for each titration in Part B. Calculate the molarity of Mg2+ in this solution. n

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B. Standardization of EDTA solution 0.01 M The concentration of EDTA solution supplied is about 0.01 M. Use the measuring cylinder 25 mL to measure about 25 mL Mg2+ solution, which you have prepared it in Part A and transfer it into 250 mL conical flask. Drop in NaOH 1.0 M solution to adjust the pH until 10. Check the pH using pH paper. Add teaspoon of EBT indicator into the solution. Do the titration as soon as possible to avoid the formation of Mg(OH)2. During the end point, the colour will change from dark red to sky blue. Repeat the same procedures two more times. Calculate the concentration of the EDTA provided. C. Determine the concentration of Mg2+ in tap water Add 100 mL of tap water into a 250 mL conical flask. Add 2 mL of buffer solution pH 10 and 2 drops of sodium sulfide 1% solution. Leave the mixture for 30 minutes. Add 1/6 teaspoon of EBT indicator and do the titration using EDTA until all the red colour become sky blue. Repeat the same procedures two more times. Calculate the concentration of Mg2+ in the tap water. D. Determination of bismuth You are given a bismuth nitrate [Bi(NO3)3] solution with unknown concentration. Transfer about 10 mL of a given solution into a 100 mL volumetric flask. Top up the solution with distilled water until to the mark of the flask. Shake the solution thoroughly to mix the solution completely. Transfer 25 mL of the solution into a 250 mL conical flask. Follow either one of the titration methods as below: D.1 Using pyrocatechol as the indicator Add 0.5 grams of pyrocatechol into the solution. This solution must appear in light yellow colour. Check the pH using pH paper to make sure the pH is about 2. Adjust the pH of the solution using nitric acid 0.1 M until the pH is about 2. Do the titration until the solution change to colourless. Repeat the whole procedures two more times. Calculate the concentration of Bi3+ in unknown solution. D.2 Using xylenol orange/potassium nitrate mixture as indicator Adjust the pH of the solution approximately to 1 by adding dilute nitric acid 0.1 M. Check the pH using pH paper. Add 30 mg (0.03 g) of xylenol orange/ potassium nitrate mixture (see preparation procedure as below) and then titrate it with standard 0.01 M EDTA solution until red colour starts to fade. From this point, add the titrant slowly until the end point is reached and the indicator color changes to yellow. Repeat the whole titration procedures two more times. 1 mole EDTA = 1 mole Bi3+.

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*For preparing the xylenol/potassium nitrate mixture, mix 1 g of xylenol orange into 99 g of potassium nitrate and shake the mixture thoroughly. RESULTS A. Preparation of standard solution of Mg2+ Weight of MgCl2.6H2O Mole of MgCl2.6H2O Molarity of MgCl2.6H2O solution

B. Standardization of EDTA solution 0.01 M Concentration of EDTA solution given Volume of Mg2+ solution used Volume of EDTA solution needed Actual concentration of EDTA solution Average concentration of EDTA solution C. Determination of Mg2+ concentration in tap water Volume of tap water used Volume of EDTA solution needed Concentration of Mg2+ in tap water Average concentration of Mg2+ in tap water

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D. Determination of Bi3+ concentration in unknown solution Volume of Bi(NO3)3 solution used Volume of EDTA solution needed Concentration of Bi3+ in unknown solution Average concentration of Bi3+ in unknown solution POST-LAB QUESTIONS 1. What type of ligand is EDTA? 2. Draw the structure of complex compound formed from Mg2+ and EDTA. 3. Why the formation of metal-ion/EDTA complexes is dependent upon the pH?

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EXPERIMENT 7: SYNTHESIS AND ANALYSIS ON A POTASSIUM TRIOXALATOALUMINATE (III) TRIHYDRATE COMPLEX, K3[Al(OX)3].3H2O INTRODUCTION Complex can be defined as substance contains a large number of elements which are directly attached to metal especially transition metals. The elements that are attached to the metal are called ligand. There are many types of ligands, for example, H2O, NC5H5 and etc. In other cases, many metallic salts exist as a hydrated salt, for example CoCl2.6 H2O, which contains the ion complex such as [Co(H2O)6]2+. Six water molecules are attached to the central metal ion, Co, therefore the water molecules are the ligand in this complex. Those ligands attached to the central metal always arranged in proper way and their geometrical structure of the complex can be determined. For example, if the coordination number of a complex is 4, geometrical structure of this complex could be either tetrahedral, seesaw or square planar. If the coordination number is 6, then it will be octahedral. Many complexes present in solid form and their geometrical structure can be determined using Fourier Transform-Infra Red (FT-IR), Nuclear Magnetic Resonance (NMR), UVSpectroscopy and X-ray crystallography. In this experiment, the complex consists of K+ ion and [Al(OX)3]3- complex ion. In the normal condition, this complex will dissolve in water but not dissolve in organic solvents such as ethanol. Therefore, the organic solvent, which is miscible with water such as ethanol, acetone and methanol, can be used for precipitation from aqueous solution. REAGENTS AND METHOD Reagents Aluminium powder, potassium hydroxide, oxalic acid, ethanol 95%, acetone, KMnO4 solution 3.2 g/L, H2SO4 acid 0.1 M and NaOH 0.1 M Apparatus 100 mL Beaker, 25 mL measuring cylinder, Bunsen burner, Buchner filter, filter paper, hotplate, glass rod, petri dish, test tube, 100 mL volumetric flask, thermometer, 250 mL conical flask.

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Method A. Synthesis of potassium trioxalatoaluminate (III) trihydrated, K3[Al(OX)3].3H2O Weigh about 0.5 g aluminium powder into a 100 mL beaker. Add 3 g of potassium hydroxide (KOH) into that beaker and add 25 mL of distilled water. (This step has to be done in the fume chamber due to the emission of H2 gas). When the reaction has completed (bubbling stop), heat the solution until it boils for 1 minute. Filter out the residue using Buchner filter and rinse the residue with 5 mL of hot distilled water. Leave the residue and collect the filtered hot solution for the following steps. Add 7.0 g of oxalic acid slowly into the solution and stir it. Use the glass rod to scratch the beakers wall and at the same time put it in the ice bath until the precipitate form. Filter out the precipitate from the solution with Buchner filter. Rinse the precipitate with 5 mL ethanol twice and followed by 5 mL of acetone once. Put the precipitate on a piece of petri dish and leave it dry in the open air. It will only take few minutes in this process. Once the precipitate dry, weigh it and keep it for next procedures. B. Analysis on K3[Al(OX)3].3H2O i) To test whether the complex will decompose or not under analysis condition Get two tubes; fill each of them with 1 mL of KMnO4, 3.2 g/L solution and 3 mL of distilled water. For the first test tube, add 2 drops of H2SO4 acid 1.0 M, whereas for the second test tube, add 2 drops of NaOH 0.1 M solution. For each of the test tube, insert just a little complex that have been synthesized and heat them. Observe the changes inside the test tube. ii) To standardize KMnO4 solution Weigh accurately 0.64 g of oxalic acid dehydrated (H2C2O42 H2O) and dissolve it in 50 mL of distilled water. Transfer the solution into a 100 mL volumetric flask. Use the distilled water to top up the solution to the mark of the flask. Transfer 25 mL of oxalic acid solution that you have prepared into a conical flask and add about 50 mL of H2SO4 acid 1.0 M. Heat the mixtureSemester I, 2011 - 2012 Page 31


solution until 65 C and do the titration directly with KMnO4 solution. Before the end point, make sure the temperature of the solution still remain 55 C. If not, heat it to 60 C. This titration must be repeated for at least two times. The chemical equation of the filtration is shown as below. 2 MnO4- + 5 H2C2O4 + 6H+ 2 Mn2+ + 10 CO2 + 8 H2O iii) To analyze the percentage of oxalate in the complex Weigh accurately 0.2 g of the complex and dissolve it into 50 mL of H2SO4 acid 0.1 M in a conical flask. Heat this solution and do the titration following the procedures in the part of standardization of KMnO4 solution above. Repeat this part two times.

RESULTS A. Synthesis of K3[Al(OX)3].3H2O

Colour of the precipitate Weight of the precipitate, g

B.

Analysis on K3[Al(OX)3].3 H2O (i) To test whether the complex will decompose or not under the analysis condition Observation Equation

Test tube 1

2

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(ii)

To standardize KMnO4 solution

Weight of oxalic acid used Concentration of oxalic acid prepared Volume of KMnO4 solution used Reaction equation (oxalic acid with KMnO4) Mole ratio (oxalic acid: KMnO4) Concentration of KMnO4 solution determined Average concentration of KMnO4 solution (iii) To analyze the percentage of oxalate in the complex

Theoretical percentage of oxalate group in K3[Al(OX)3].3 H2O complex, % Weight of complex used Concentration of KMnO4 solution determined Volume of KMnO4 used Reaction equation of oxalate group with KMnO4 in acidic condition Experimental percentage of oxalate group in the complex, % Compare the theoretical and experimental percentage, what is the final conclusion that you can make?

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POST-LAB QUESTIONS 1. What is the main purpose of scratching beakers wall during the synthesis of the complex? 2. Why ethanol followed by acetone is used to rinse the precipitate? 3. What is the residue that may contain in the precipitate? 4. Why KMnO4 has to be standardized before being used for the titration? 5. Why oxalic acid solution must be heated before being used for the titration?

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EXPERIMENT 8: SYNTHESIS AND TESTING ON COPPER (II) COMPOUND INTRODUCTION The common oxidation states of copper are +1 and +2. In oxidation state of +1, copper solution always colourless. Whereas for copper(II), blue or blue-green are the common colours. However, there are some exception to this, for example CuI and Cu2O, which are yellow and red in colour respectively.

In this experiment you will synthesize a compound by adding NH3 to a concentrated aqueous solution of copper sulfate. The blue CuSO4 solution will turn a still deeper blue and a mass of small deep blue-to-violet crystals will form as ethyl alcohol is added. On the basis of the analysis of this solid for Cu2+, SO42- and NH3, you will be able to propose a formula for the compound.

Four principal species are present initially in the reaction mixture: copper (II) ions [actually Cu(H2O)62+ ions], ammonia molecules (NH3), sulfate ions (SO42-), and water. The product of the synthesis is therefore presumed to be formed by the reaction of two or more of these species. Ethanol is also present, but it is an indirect participant in the reaction. In aqueous solution ethanol, which is miscible with water but of lower dielectric constant, the solubility of ionic compounds will be decreased. The marked color change that occurs in the reaction is an important clue to the nature of the product. The product is analyzed for copper (II) ions, sulfate ions, and ammonia molecules. Water is determined as the mass of a sample of the compound that is not accounted for as one of these three species. REAGENTS AND METHOD Reagents CuSO45H2O, aqueous ammonia 8 M, aqueous ammonia 1 M, ethanol, distilled water, zinc granule, SnCl2 solution, NaS2 solution

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Apparatus Beaker 250 mL, spatula, glass rod, measuring cylinder 50 mL, filter tunnel, conical flask 250 mL, test tubes, dropper Method A. Preparation of copper ammine sulphate 1. Weigh out 2.0 g of copper sulfate pentahydrate, CuSO45H2O, on a triple beam balance, and place the crystals in a 100 mL beaker. 2. Weigh filter paper and watch glass. 3. Add 2 - 4 mL of distilled water to the solid and then in a hood add 4 mL of 8 M NH3 (concentrated ammonia). Stir to dissolve the crystals as ascertained by holding the beaker up to the light. 4. Over a period of 1 minute, slowly add 4 mL of 95% ethanol to the solution, stir, and cool to room temperature. 5. Set up an apparatus for vacuum filtration. Moisten the filter paper and turn on the aspirator. Carefully filter the slurry of crystals that has formed in the copper-containing solution and suck off all the solution. 6. Dry the precipitate filtered in desiccators overnight. 7. Weigh the dried precipitate and calculate the percentage yield of the complex. B. Copper (II) ions testing Weigh approximate 0.5 g of CuSO45H2O and dissolve it in 50 mL of distilled water. In each of four test tubes, insert 5 to 6 drops of CuSO4 solution and conduct the following tests: a) Zn metal test To the first test tube, add in a piece of Zn granule. Observe the colour changes in the test tube. b) Tin (II) chloride test To the second test tube, insert 10 to 15 drops of SnCl2 solution 1.0 M and leave the precipitate settle down. Record the colour of the precipitate and solution.

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c) Ammonia test To the third test tube, insert 10-15 drops of NH3 solution 1.0 M and observe the reaction occurs. d) Sodium sulphide test To the fourth test tube, insert 4 to 5 drops of NaS2 solution 1 M and observe the reaction occurs. C. Spectroscopic Analysis Obtain the IR spectrum of the cobalt complex and identify the important bands. If you have access to a spectrometer capable of measuring the reflectance spectra of solid samples in the UV-vis region, record the spectra of the complexes and relate the spectra to the colors of the compound.

RESULTS A. Preparation of copper ammine sulphate i) Quantitative data: Mass of CuSO4.5H2O Mass of filter paper and watch glass Mass of filter paper, watch glass and dried complex Mass of dried complex Theoretical yield, %

Experimental yield, %

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ii) Qualitative data: Describe the process of dissolving the copper sulfate pentahydrate and characteristics of the resulting solution. Describe the changes that occurred as the ammonia was slowly added. Describe the changes that occurred upon addition of the ethanol.

B. Copper (II) ions testing Test Zn metal test Observation Equation

Tin (II) chloride test

Ammonia test

Sodium sulphide test

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POST-LAB QUESTIONS 1. Why is it necessary to cool the solution containing the blue crystalline compound in order to get crystals to form? 2. Why is it necessary to know the masses of reactants used in the reaction used in the reaction if you are going to calculate the theoretical yield later?

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STK1084 Lab Manual

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