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Chemistry 16 Lab Manual

Table of Contents Laboratory Safety and Laboratory Guidelines Common and Special Laboratory Equipment Materials and Other Requirements Common Laboratory Operations and Techniques Experiment 1 ………………………………………………………………………….. 10 Properties of Matter Experiment 2 ………………………………………………………………………….

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12 Chemical Changes Experiment 3 …………………………………………………………………………. 15 Classification of Matter Experiment 4 …………………………………………………………………………. 17 Chemical Nomenclature: The Language of Chemistry Experiment 5 …………………………………………………………………………. 22 Water of Hydration

Experiment 6 …………………………………………………………………………. 25 Gases Experiment 7 …………………………………………………………………………. 27 Oxygen Experiment 8 …………………………………………………………………………. 29 Heat and Volume Effects Experiment 9 …………………………………………………………………………. 31 Flame Test Experiment 10 ……………………………………………………………………….. 32 Electromotive Series Experiment 11 ………………………………………………………………………… 33 Oxidation – Reduction Reactions/ Some Aspects of Corrosion Experiment 12…………………………………………………………………………. 35 Colligative Properties GENERAL INSTRUCTIONS TO THE STUDENTS Apparatus Check each piece of apparatus, which you find in your locker from the duplicate list furnished to you by your instructor. Sign your name and submit to your instructor. The instructor signs the checklists and gives one copy to you for your safekeeping. ? Provide your locker with reliable padlock. You are responsible for all the apparatus issued to you. Towards the end of the semester you have to replace or give a deposit for any piece which you have lost or broken. If you have partners, each of you will share equally any loss or breakage of apparatus kept in your lockers and those orrowed from the stockroom. A clearance duly signed by the laboratory attendant is a requirement for taking the final examination. NO CLEARANCE, NO FINAL EXAMINATION. ? General apparatus, e. g. , Bunsen burner, thermometer, iron stand, clamps, etc. and special apparatus may be borrowed from the laboratory attendant. ? Borrowing of apparatus from the stockroom should be done during the first 30 minutes of the laboratory period. Materials and Other Requirements You have to provide yourself with the following materials and supplies besides the apparatus in the laboratory locker and the stockroom: Group |Individual | |Masking/paper tape |Tissue paper |Vials with cover (5 pcs) |Lab notebook | |Pair of scissors |Rags |Medicine dropper (3-5 pcs) |Lab manual | |Aspirator |Marking pens |Rubber tubing (2 ft) |Lab gown | |Wire gauze |Filter paper |Newspaper/scratch paper |Hand towel | |Wash bottle |Tray |Stirring rod |Mask | |Liquid detergent |Match |Corks/rubber stoppers |Goggles | |Test tube brush |Test tube holder |Padlock with keys | | Laboratory Work Laboratory work is an integral and essential part of any chemistry course. Chemistry is an experimental science – the compounds and reactions that are met in the lecture and classroom work has been discovered by experimental observation. The purpose of laboratory work is to provide an opportunity to observe the reality of compounds and reactions and to learn something of the operations and techniques. Safety is Top Priority ? All students are required to wear a lab gown during each experiment. This will be strictly enforced to avoid accidents caused by chemical spills and the like. Safety glasses, goggles or eye shields must be worn during the experiment. Contact lenses should not be worn. ? Shorts, skirts, sandals, slippers are not allowed in the laboratory. Secure long hair. ? Never taste, smell, or touch a chemical solution unless specifically directed to do so. Individual allergic or sensitivity responses to chemicals cannot be anticipated. If any chemical comes in contact with any other parts of your body or clothes, wash thoroughly with plenty of water. ? Procedures involving the liberation of volatile or toxic flammable materials shall be performed in a fume hood (e. g. , H2S, HCN). ? Never heat a flask or apparatus that is not opened to the atmosphere. Always pour waste acid, used KMnO4, organic solvents and solutions of heavy metals into their respective disposal jars, never into the sink. ? Replace the cover of every container immediately after removal of reagent. Deposit insoluble refuse such as pieces of paper, wood, glass cork in the waste basket, never into the sink or on the floor ? All accidents, injuries, breakages and spillages, no matter how minor, must be reported immediately to the instructor. ? Eating, drinking, smoking and playing inside the laboratory are strictly prohibited. Your hands may be contaminated with “unsafe” chemicals. ? Unauthorized experiments, including variations of those in the laboratory manual, are strictly prohibited.

If your chemical intuition suggests further experimentation, consult with your instructor first. ? Unauthorized person(s) shall not be allowed in the laboratory. ? Maintain a wholesome, businesslike attitude. Horseplay and other careless acts are prohibited. ? The tabletop must be cleared of unnecessary materials. Put all bags and books in designated areas. ? Solids, water and other liquids spilled on your tabletop must be cleaned up as soon as possible ? No electronic equipment (laptops, ipod, mp3s, cellphone, etc. ) will be switched on while working in the lab. For Economic Reasons ? Always turn off the burner as soon as you are finished using it. Get only the amount of the reagent, which you need in the experiment. Use spatula for solid reagents and pipet for liquid ones. ? Never return any excess reagent to a bottle, unless specifically directed, to avoid contamination Before leaving the room, see to it that: ? Your locker is locked ? Your assigned water and/or gas outlet(s) are turned off ? The tabletop and the floor near your working area are clean and dry Collecting Data ? Record all data as they are being collected on the laboratory notebook. Data on scraps of paper (such as mass measurements in the balance room) will be confiscated. ? Record the data in ink as you perform the experiment. If a mistake is made in recording data, cross out the incorrect data entry with a single line (do not erase, white out or obliterate) and clearly enter the corrected data nearby. If a large section of data is deemed incorrect, write a short notation as to why the data are in error, place a single diagonal line across the data, and note where the correct data are recorded. Assessment: Evaluation of the students’ progress will be based on performance laboratory experiments; written reports of laboratory work and exams. The distribution is as follows: Exams35% Performance/ Attendance15% Written Laboratory report35% Pre-laboratory write-up/ Data notebook15%

Laboratory Course Policies: 1. Arrive on time. The overview and description of the lab exercise, and the questions you need to answer in your written reports are usually given at the start of each session. These could be valuable to the success of you laboratory course.

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2. Note all laboratory safety policies at all times. You are required to wear lab coats and safety glasses while in the lab. You must wear your protective gear at all times that any lab work is underway. Failure to observe safety precautions may result in your being dismissed from the laboratory class. 3. Request all chemicals and materials that you may need from the stock room at least 30 mins. head of the scheduled experiment. At this stage in your studies, you are expected to be able to work independently and responsibly. 4. Written reports of laboratory work are due at the start of the following lab session. Reports that are late will be penalized for each day of late submission(10% deduction per day). 5. Laboratory techniques, including your preparedness and participation in each laboratory activity, good note-keeping and ability to work well with your partner will be graded accordingly. 6. Read and plan you work before every laboratory class. Prepare a pre-laboratory write-up at the start of the lab period and prior to starting your laboratory work.

You will not be allowed to perform the experiment without a pre-lab write up. The pre lab should include the following sections,: Experiment #, Title of Experiment, Date, Objectives of the Experiment’ Theoretical Framework, Materials and Methods, Expected Outcome. Sign and Date each pre-lab write-up. During the conduct of the experiment, record all your raw data in the same notebook. 7. Written Reports should be written on a short-sized bond paper and will have the following components: Name, Laboratory partner/s, Discussion of Results, Calculation/s(if any), Question/s, and Answer/s, Conclusion/s, Comment on you and your partner’s contribution towards the successful completion of the laboratory activity.

Submit your lab report as a group, write your group number and experiment number as the subject of the email. COMMON LABORATORY OPERATIONS AND TECHNIQUES 1. BUNSEN BURNER A. Lighting the Burner a. Examine the parts of the Bunsen burner. Make a sketch of the burner, label and state the function of its parts. b. Attach the rubber tubing from the burner to the gas outlet on the lab bench. Bring the lighted match or striker up 4-5 cm above the barrel while opening the gas valve. c. Adjust the gas supply so as to have a flame of not more than 8 cm high. Close the air holes of the burner and observe the appearance of the flame. Hold the porcelain dish on this flame for a moment.

What is deposited on the porcelain dish? d. Open the air holes until the flame is pale blue and has two or more distinct cones. A slight buzzing or roaring sound is characteristic of the hottest flame from the burner. Too much air may blow the flame out. Adjust the air intake until the roaring stops. What is the effect on the flame upon opening of the air holes? Does this type of flame have the same effect on the porcelain dish? Why? Spray powdered charcoal on the flame and note its effect. What makes the flame luminous? e. When the best adjustment is reached, three distinct cones are visible. Always use this kind of flame unless directed otherwise. f.

Extinguish the flame when it is not being used, by closing the gas valve. B. Determining the Flame Temperatures a. Wet a piece of cardboard and hold it vertically through the center of the flame, with the lower end of the cardboard resting against the top of the burner. b. Remove the cardboard as soon as it shows a tendency to char. From the scorched portions note the relative temperature of the different parts of the flame. c. Draw a sketch of the flame to illustrate the different regions. 2. GLASS MANIPULATION A. Cutting a. Place the glass tubing flat on the table. Make a single scratch with a sharp triangular file 30cm from one end of the glass tubing. b.

Grasp the glass tubing with both hands and place the thumbs one cm beside the scratch. Position the thumbs such that they are opposite the scratch. c. Break the glass tubing by applying a gentle pressure. If it does not yield to gentle pressure, make a deeper scratch. d. The edges of the cut glass tubing are sharp and should be polished by rotating it at the non-luminous portion of the burner’s flame. This is to prevent the sharp edges of the glass from ruining corks and rubber tubing as well as cutting your fingers. B. Bending a. Take a piece of glass tubing about 30 cm long and hold it lengthwise over the flame. b. To bend the glass tubing properly, it must be heated uniformly over a length of 5 to 8 cm.

This can be done using a flame spreader. c. Roll the tube back and forth until it has become quite soft. d. When it has become sufficiently soft, (i. e. , the glass tubing begins to take a pink color and sag gently) take it out of the flame. e. Bend quickly to the desired angle (30° or 90°) and hold until it hardens. Try to get a good idea of the angle before you begin to work so that you may work rapidly and secure the desired bend at once. f. Make one right angle and one 30O bent glass tubing. NOTE: Reheating and re-bending produces unsightly and often frail apparatus. C. Drawing Out a. Roll the center of a 10cm glass tube over the flame until it softens.

The tube must be constantly rotated, to prevent the softened portion from sagging. b. Quickly remove it from the flame, and while holding it in a vertical position, gently pull the ends apart until the bore at the stretched portion is of the desired diameter. c. Cut to the desired nozzle length and fire polish the tip. D. Boring corks and rubber stoppers a. Select a cork that will fit into the mouth of the flask or test tube. b. Soften by rolling it between the tabletop and the palm of your hand. Select a sharp cork borer one size smaller than the glass tube that will be inserted. c. Place the cork on the desk and gently twist the borer in until it is halfway through the cork.

Then withdraw the borer and finish the hole from the other end of the cork. d. Smoothen the hole in the cork with a round file. e. If the hole is too small, enlarge it by carefully filing with a round file. Only small adjustment should be made in this way. f. Rubber stoppers are bored in the same manner as mentioned. Select a very sharp borer one size larger than the hole to be made, and wet it with glycerin. Proceed as in boring the cork, but do not apply too much pressure. E. Inserting a glass tubing through a cork/rubber stopper NOTE: This operation is the most common cause of accidents in the laboratory. a. Wet the cork and the glass tubing with water. b.

Place your hand on the tubing 2-3 cm away from the stopper. Protect your hand with a towel. c. Simultaneously twist and push the tubing slowly and carefully through the hole. 3. CLEANING OF GLASSWARE a. Clean all glassware with a soap or detergent solution. Use a brush if appropriate. b. Once the glassware is thoroughly cleaned, rinse several times with tap water and then once or twice with distilled water. c. Roll each rinse around the entire inner surface of the glass wall for a complete rinse. Discard each rinse through the delivery point of the vessel (e. g. , beaker spout). d. Invert the clean glassware on a clean paper towel or rubber mat to dry.

Do not dry any glassware over direct flame. e. The glassware is clean if, following the final rinse, no water droplets adhere to the clean part of the glassware. f. If you must use a piece of glassware while it is still wet, rinse it with the solution to be used in the manner described in step 5c below. 4. TRANSFERRING OF LIQUIDS/SOLUTIONS a. When the liquid or solution is to be transferred from a reagent bottle, remove the glass stopper and hold it between the fingers of the hand used to grasp the reagent bottle. Never lay the glass stopper on the laboratory bench; impurities may be picked up and thus contaminate the liquid when the stopper is returned. b.

To transfer a liquid from one vessel to another, hold a stirring rod against the lip of the vessel containing the liquid and pour the liquid down the stirring rod, which, in turn, should touch the inner wall of the receiving vessel. Return the glass stopper to the reagent bottle. c. Do not transfer more liquid than is needed for the experiment; do not return any excess liquid or unused liquid to the original reagent bottle. 5. MEASURING VOLUME OF LIQUID/SOLUTIONS a. The eye should always be level with the meniscus when you are making a reading. b. For measurements of clear or transparent liquids/solutions, the volume is read using the lower meniscus. For colored liquids/solutions, the upper meniscus is used. 6. HEATING A LIQUID/SOLUTION IN A TEST TUBE NOTE: Never fix the position of the flame at the base of the test tube and never point the test tube to anyone.

The contents may be ejected violently if the test tube is not heated properly. a. The test tube should be less than one third full. Hold the test tube with a test tube holder at an angle of about 45? with the cool flame. A cool flame is a nonluminous flame supplied with a reduced amount of fuel. b. Move the test tube circularly in and out of the flame, heating from top to bottom. 7. PRECIPITATION a. Place 2 mL of sodium chloride solution in a test tube and slowly add 2 mL of silver nitrate solution. Write the balanced chemical equation for this reaction. NOTE: Be careful in handling silver nitrate solution. This solution may leave dark stains on skin, clothes or bench top. b.

The solid formed is the precipitate and in this case, the slightly soluble silver chloride. Allow the precipitate to settle. c. Add a few drops of silver nitrate solution. Continue addition until no precipitation is observed. Divide the mixture into two portions and keep these for procedure 8. 8. SEPARATING A LIQUID FROM A SOLID A. Filtration a. Preparation of the filter paper to be used for gravity filtration: i. Cut out a 5” x 5” piece of filter paper. Fold the filter paper in exact halves and fold it again crosswise into two. ii. Make a small tear in one corner. This tear seals the paper against the inflow of air to the underside of the filter paper. iii.

Open the folded paper so as to form a cone. iv. Place it in a funnel. Moisten it with a little water and press it against the top wall of the funnel to form a seal. The filter paper must always be smaller than the funnel. v. Support the funnel with a clamp or a funnel rack. b. Transfer the precipitate formed from the previous activity by carefully pouring the mixture, with the aid of a glass rod, into the filter paper. The liquid that passes through the liquid is called the filtrate. c. The tip of the funnel should touch the wall of the receiving beaker to reduce any splashing of the filtrate. d. Fill the bowl of the funnel until it is less than two-thirds full. e.

Always keep the funnel stem full with the filtrate; the weight of the filtrate creates a slight suction on the filter in the funnel, thus this hastens the filtration process. f. Set aside both precipitate and filtrate for the next two activities. B. Decantation a. Transfer the precipitate retained in the filter paper into a beaker by rinsing the filter paper with jets of water from a wash bottle. b. Allow the solid to completely settle at the bottom of the vessel for several minutes. c. Transfer the liquid (called supernatant) into another container with the aid of a clean stirring rod. d. Do this slowly so as not to disturb the solid. Is this method applicable for the separation of all solid-liquid mixtures? Why? e. Rinse the precipitate with water and decant again. f. Which of the two separation methods (i. e. decantation or filtration) is better in isolating the precipitate? Why? E. Evaporation a. Pour the filtrate obtained from 8A into the evaporating dish. Place the evaporating dish on a wire-gauze supported on an iron ring clamped to an iron stand. Heat the dish over a cool flame. b. Continue heating until crystals begin to appear. Cover the dish with a watch glass and allow the contents to cool. The solid remaining after evaporation is called the residue. What is the composition of the residue? 9. WEIGHING a. Weigh 0. 5 g of sand. Weighing may be done on platform balance or on an analytical balance. Rough weighing (to the nearest half gram), can be done on the platform balance.

The analytical balance is used to get more accurate mass measurements. b. The properties of the substance will often determine the nature of the container where it is to be weighed. Use a weighing paper, a watch glass, a beaker, or some container to measure the mass of the chemicals. c. Do not place the chemicals directly on the balance pan. When in doubt as to what container to use, ask your instructor. TECHNIQUE IN HANDLING CHEMICALS d. A reagent is a substance which has a definite composition and a set of specific properties. The strong solutions are marked “concentrated” and the weak solutions, “dilute”. Some examples of the reagents are: Sulfuric acidH2SO4Ammonia NH3

Hydrochloric AcidHClSodium hydroxide NaOH Acetic acidCH3COOHCalcium hydroxide Ca(OH)2 e. Before getting the desired amount, read the label twice to be sure it is the correct chemical at the right concentration. Transfer the needed amount into the receiving container. Once removed, these should never be returned. f. Do not take out more than what is needed to minimize waste. Do not return excess chemicals to the reagent bottle. In pouring reagents from bottles, don’t place the stopper on the table but hold it between your fingers. g. Never touch, taste or smell chemicals unless specifically directed to do so. ExPERIMENT Properties of Matter

This experiment presents several of the properties used to identify a sample of matter. The data gathered are interpreted by the use of some quantitative method. For safety and accuracy of results, the experimenter should make sure that all set-ups used should be properly checked for possible connection leaks and other errors. Stirring rod should be used to ensure uniform distribution of heat when heating liquids in an open container. The heat should also be regulated especially when heating closed set-ups. Laboratory techniques included are: measurement and transferring of liquids, weighing and heating of liquids and solids. MATERIALS AND APPARATUS 25 or 50-mL graduated cylinder |Thermometer |Cork and/or rubber stoppers | |50-mL distilling flask |Bunsen burner |Top loading balance | |250-mL beaker |Rubber tubings |Condenser | |25-mL Florence flask |Iron stand |Oil | |Test tube |Iron ring |Sulfur powder |2-3 iron clamps |Wire gauze |Isopropyl alcohol | | | |Lead pellets | PROCEDURE 1. BOILING POINT a) Measure 25 mL of isopropyl alcohol and record the initial temperature. 32 degrees a) Transfer it into a 50-mL distilling flask. Assemble the distillation set-up (consult the instructor). b) Warm the set-up gently with a Bunsen burner. Take temperature readings at one-minute time intervals until the liquid begins to boil, and two more minutes thereafter. c) Continue distilling until the flask is almost dry. Pour off the liquid still present in the flask. ) Transfer the distillate into the distilling flask and repeat the distillation process. e) Make a graph of your data with time on the x-axis and temperature on the y-axis. Compare the two graphs. f) Determine the boiling point of the liquid from the graphs. Look for the standard boiling point of isopropyl alcohol and get the % error of the boiling point obtained experimentally. 2. MELTING POINT a) Place about 1-gram of sulfur powder into a dry test tube. Clamp the test tube vertically into the oil bath. See to it that the solid is below the oil level. a) Hang the thermometer into the test tube such that it is covered by the solid and does not touch the sides and bottom of the test tube. ) Heat the oil bath gradually and take temperature readings at one-minute intervals until the solid has completely liquefied, and two more minutes thereafter. c) Make a graph of your data with time on the x-axis and temperature on the y-axis. Determine the melting point of sulfur from the graph. Look for the standard melting point of sulfur and get the % error of the melting point obtained in the experiment. NOTE: Stir the oil bath so that the heat is uniformly distributed. 3. DENSITY OF A LIQUID a) Clean and dry the Florence flask. Weigh the dry flask and the rubber stopper on a top loading balance and record the mass. b) Fill the flask with distilled water until the liquid level is nearly to the brim.

Put the stopper on the flask in order to drive all the air and excess water. Work the stopper gently into the flask so that it is firmly seated into position. c) Wipe any water on the outside of the flask and soak up all excess water from around the top of the stopper. d) Again, weigh the flask, which should be completely dry on the outside and full of water, and record the mass. e) Calculate for the precise volume of the flask given the standard density of water, the temperature of the laboratory and the mass of water in the flask. f) Empty the flask, dry it and fill it with isopropyl alcohol. Stopper and dry the flask as you did when working with water.

Record the weight of the flask filled isopropyl alcohol. g) Calculate the density of isopropyl alcohol and determine the % error using its standard density. 4. DENSITY OF A SOLID a) Use the same flask from the previous procedure for this part. Dry the flask completely and add small chunks of lead metal into the flask until it is about half full. b) Weigh the flask, with its stopper and the metal, and record the mass. Determine the mass of the metal in the flask. c) Fill the flask with water, leaving the metal in the flask, and then replace the stopper. Roll the metal around the flask to make sure that no air is trapped between the metal pieces. ) Refill the flask if necessary, and then weigh the dry stoppered flask full of water plus the metal sample. e) Compute for the density of the lead using the data obtained in this section and in part 3. Determine the density of the metal and compute for the % error. QUESTIONS 1. Interpret the graphs obtained in parts 1 and 2. What changes occur at the different regions of the graph? 2. What kind of properties are boiling point, melting point and density? 3. Which of these properties may be used to identify a sample of matter? Why? 4. Is one property sufficient to establish the density of the substance? 5. What is the identity of the distillate in Part 1? What is your basis?

ExPERIMENT CHEMICAL CHANGES This experiment presents different types of chemical change. Some quantitative methods are included to emphasize proper data handling and interpretation of results. Formula writing and setting up of simple chemical equations are introduced. It is to be emphasized that the experimenter should always take note of any physical evidence that a chemical reaction is taking place. Such physical evidences include the formation of a precipitate, change in color of the solution or precipitate, evolution of a gas, and absorption or evolution of heat. ? Evolution of gas. This evolution may be quite rapid or it may be a “fizzing” sound. Appearance or Disappearance of precipitate. The nature of the precipitate is important; it may be crystalline, it may have color, it may merely cloud a solution. ? Evolution or Absorption of Heat. The reaction vessel becomes warm if the reaction is exothermic or cools if the reaction is endothermic. ? Change in color. A substance added to the system may cause a color change. Also included are the common laboratory operations such as measurement and transferring of liquids, precipitation, decantation, filtration, washing and transferring of precipitates, drying of solids, weighing, testing for acidity and basicity, and testing for completeness of a reaction.

This experiment also emphasizes the need for gradual mixing of reactants to make certain the maximum recovery of the product, and the importance of washing, to ensure the purity of the product. MATERIALS AND APPARATUS |50-mL graduated cylinder |Watch glass |Zinc dust | |250-mL beaker |Evaporating dish |0. 100 M Cu(NO3)2 | |250-mL Erlenmeyer flask |Pair of scissors |6. 00 M NH3 | |Funnel |Filter paper |6. 0 M NaOH | |Bunsen burner |Litmus paper |6. 00M HCl | |Stirring rod |Medicine Dropper |6. 00 M H2SO4 | PROCEDURE 1. Precipitation of Copper (II) hydroxide a) Measure 10-mL of 0. 100 M Cu(NO3)2 solution in a 250-mL beaker. a) Add dropwise with constant stirring about 0. 5 mL 6. 00 M NaOH solution. b) Place a piece of litmus paper on a dry watch glass and moisten it with the solution using a stirring rod. c) If it is not yet alkaline, add more NaOH. Record any change in color of the solution and describe the precipitate. 2.

FORMATION OF COPPER (II) OXIDE a) Boil the contents of the beaker in part 1 for about 2 minutes with constant stirring to prevent “bumping” which may result in loss of material. The precipitate should change in color. b) Allow the copper (II) oxide precipitate to settle. Take note of the change in color of the precipitate. c) Test the supernate with a few drops of 6. 00M NaOH. If cloudiness is observed, continue the addition of the base until precipitation is complete. d) Heat the solution again with constant stirring, until all the precipitate has changed in color. Record the color changes that occur. What is the evidence of complete precipitation?

What is the composition of the supernate? 3. CONVERSION OF COPPER (II) HYDROXIDE TO COPPER (II) SULFATE a. Let the precipitate settle until the supernate is clear. Decant the supernate through a filter paper into the Erlenmeyer flask. b. Wash the precipitate in the beaker using 10 mL of water. Let the precipitate settle and decant the wash water through the filter paper into the Erlenmeyer flask containing the filtrate. c. Repeat the process, so that the precipitate is washed a total of four times. d. Wash the same filter paper with about 1 mL 6. 00 M H2SO4 dropwise, catching the filtrate in the beaker containing copper (II) oxide precipitate. e.

Rotate or stir the contents of the beaker to dissolve the solid. Add some more H2SO4 to dissolve the precipitate completely. f. Wash the filter paper again, this time with 10 mL water, catching the wash water in the same beaker. Record your observations. 4. REDUCTION OF Cu (II) IONS TO METALLIC COPPER a. To the solution (from 3), gradually add with constant stirring, about 1. 5 g zinc dust in minute amounts. CAUTION: Stir until no further reaction is observed before adding more zinc to make the solution colorless. b. Test for the completeness of the reaction by adding a few drops (1-2 drops) of the solution into a test tube containing about 1 mL of 6. 0 M NH3. If a colored solution is obtained, compare this with the control solution (prepare by adding a drop of 0. 100 M Cu(NO3)2 solution and 2 drops of NH3 to 1 mL water) and add more zinc into the solution with constant stirring. Repeat the process until the test with ammonia solution gives a colorless solution. c. Decant and discard the supernate in 4-b. Wash the precipitate in the beaker twice, each time using 10-mL portions of water. Decant and discard the wash water after each washing, taking care not to lose any solid. d. To the precipitate, add 10 mL water and 2 mL 6. 00 M HCl slowly and stir the contents until no more change is observed.

Let the precipitate settle, decant and discard the supernate into a waste acid jar. e. Wash the precipitate twice, each time using 10-mL portions of water. Decant and discard the wash water. f. Transfer the entire solid in the beaker to a previous weighed filter paper. Use as little water as possible to wash out the solid from the beaker. Discard the filtrate and wash water. g. Fold the filter paper containing the solid and press this between pieces of dry filter paper to remove most of the water. Place the partially dried filter paper containing the solid on a watch glass, and air dry in your locker until the next period. Weigh the solid and the filter paper.

Record all masses obtained. 5. OXIDATION OF COPPER a. Place a pinch of the weighed solid in an evaporating dish and heat the dish directly over a Bunsen burner. Observe and record your results. b. Submit the remaining solid, properly packaged and labeled, to your instructor. QUESTIONS 1. What type of process and/or chemical changes is observed in procedures 1-5? 2. Why must zinc be added very gradually to the solutions in procedure 4. a? 3. What is the purpose of the test using ammonia solution? 4. Why must HCl be added to the solid after the reaction with zinc dust is completed? 5. Why is it not advisable to dry the copper directly over a Bunsen flame? 6.

Calculate the percent recovery in the experiment. Does your result refute the law of conservation of matter? Explain. ExPERIMENT CLASSIFICATION OF MATTER Matter is classified according to its various properties and the type of changes it undergoes. There are two general types of matter, substances and mixtures. Substances are further subdivided into two types, elements and compounds. Mixtures are also of two kinds, homogeneous and heterogeneous. This experiment aims to differentiate several samples of matter. The samples are subjected to different conditions like temperature and solubility in some solvents. Chemical changes are illustrated by chemical equations. MATERIALS AND APPARATUS Beakers |Evaporating dish |Sugar crystals | |250-mL Erlenmeyer flask |Test tubes |Sodium chloride | |Funnel |Thermometer |Iodine Crystals | |Bunsen burner |Mortar and Pestle |Sulfur powder | |Glass tubing |Filter paper |Lead (II) nitrate | |Watch glass |Litmus paper |Magnesium ribbon | |Medicine dropper |Starch solution | | PROCEDURE 1. ubstances, homogeneous and heterogeneuos mixtures a. Measure out one gram of refined sugar in the balance. Dissolve the sample in 50 mL tap water. Compare the appearance of the solution with that of distilled water. Set up a simple distillation apparatus using the Erlenmeyer flask, thermometer and glass tubing. b. Distill the sugar solution and make a boiling point curve on the graphing paper. Collect the sugar solution and make a boiling point curve of the isopropyl alcohol (from experiment 1). Compare the boiling point curve of the sugar solution with that of the isopropyl alcohol. Which of the two is a substance and which is a mixture? c.

Test for the solubility of the powdered sulfur in water. Do the same with sodium chloride. Weigh out 0. 5 g of each chemical on the analytical balance. d. Grind the two together in a mortar. Note the appearance of the mixture. With a hand lens, observe the mixture more closely. Can you distinguish the sulfur from the sodium chloride crystals? e. Transfer half of the mixture into a beaker containing about 15 mL of water. Stir thoroughly then filter the resulting mixture. Identify the filtrate and the residue on the filter paper. f. Transfer the filtrate into an evaporating dish. Heat this to boiling. When the crystals begin to form, cover the dish with watch glass to prevent sputtering.

When the crystals are almost dry, stop heating the dish. g. Heat the other half of the original mixture in an evaporating dish until melting is observed. Examine the resulting product closely using a hand lens. Can you now differentiate the two components? Test its solubility in water. Record all observations. 2. ELEMENTS AND COMPOUNDS a. Take two small crystals of iodine. Place one crystal inside the test tube and heat it gently. Compare the heated and the unheated crystals with respect to state, color, solubility in water and their behavior in starch solution. b. Take a pinch of lead nitrate crystals. Observe carefully and list down its observable physical properties.

Heat it over a burner, gently at first, and then strongly afterwards until no further change is observed. List down your observations. 3. METALS AND NON-METALS a. Clamp a medium-sized test tube horizontally. Take a piece of magnesium ribbon and insert one end into a 10-cm piece of glass tubing. b. Heat the magnesium ribbon. When it begins to burn, insert the burning magnesium ribbon into the test tube until the metal has burned completely. c. Dissolve the residue in 3-mL water. Test the acidity and basicity of the solution with litmus paper. Repeat using a pinch of sulfur. QUESTIONS 1. Write all chemical equations involved. 2. Does the appearance of the sugar solution differ from that of distilled water? 3.

In part 1, which is an example of a homogeneous and heterogeneous mixture? How are the two types of mixtures differentiated? 4. What is the composition of the crystals formed after evaporation of the filtrate in 1. b? 5. Based on the results of part 1, how are substances different from mixtures? 6. Is there any evidence that would indicate a change in the identities of each of the substances heated? What are these evidences? 7. Differentiate the oxides of metals and non-metals. 8. From the results in part 2, define elements, compounds, metals and non-metal. 9. Iodine is liberated from seaweeds by the action of sulfuric acid on the ash of the seaweeds. How is it collected from the ashes? ExPERIMENT

The Language of Chemistry: Chemical Nomenclature Chemical Nomenclature is the system of naming substances. A systematic nomenclature was established by an organization of chemists called the International Union of Pure and Applied Chemistry (IUPAC). The standardized rules developed by the IUPAC are summarized below. 1. Binary Compounds 1. 1 Binary Compounds Containing Two Nonmetals If two nonmetals form a compound, the less electronegative is written first, followed by the more electronegative element. The same pattern is used in naming; the less electronegative is mentioned first, followed by the stem of the name of the more electronegative ending in –ide.

When more than one compound can be formed from the combination of two elements, Greek prefixes are used to indicate the number of atoms of each element. |CO2 |carbon dioxide | |PCl3 |phosphorous trichloride | |Cl2O |Dichlorine mon(o)oxide* | |HCl |Hydrogen chloride | *this is omitted when the more electronegative element begins with a vowel Greek Prefix |Number |Greek Prefix |Number | |Mono- |1 | Hexa- |6 | | Di- |2 | Hepta- |7 | | Tri- |3 | Octa- |8 | | Tetra- |4 | Nona- |9 | | Penta- |5 | Deca- |10 | 1. 2 Binary Compounds Containing a Metal and a Nonmetal The metal is always written first, in both the name and the formula. As with all binary compounds, the nonmetal takes an –ide ending.

There are two types that we must consider: metals with fixed (only one) oxidation number and those with variable (more than one) oxidation numbers. 1. 2. 1 Cations Monatomic ions cations retain their names as elements. The NH4+ ion, ammonium ion is named as if it were a metal ion because of its saltlike properties. |Li+ |lithium ion | |Na+ |sodium ion | |Mg2+ |magnesium ion | |Al3+ |aluminum ion | 1. 2. 2 Monatomic Anions

Monatomic anions are named using their names as elements and the suffix –ide. |C4- |carbide | |N3- |nitride | |O2- |oxide | |H- |hydride | 1. 2. 3 Metals with Fixed Oxidation Numbers The metals with fixed oxidation numbers are the IA and IIA, Aluminum and Zinc. All other metals have variable oxidation numbers. Note that no prefixes are used. NaCl |Sodium chloride | |Na2S |Sodium sulfide | |AgBr |silver bromide | |Al2O3 |aluminum oxide | 1. 2. 4 Metals with Variable Oxidation Numbers In a binary compound of a metal of this type with a nonmetal, the oxidation number of the metal must be indicated in the name. There are two methods of doing this, the classical system and the Stock or Roman numeral system. 1. 2. 4. Classical System This system can only be used for metals having two oxidation states. An –ic ending is used for the metal with the highest oxidation state and an –ous ending is used for the lowest. Also, the Latin name is used for iron (ferric and ferrous), copper (cupric and cuprous), tin (stannic and stannous) and lead (plumbic or plumbous). The classical system does not indicate the actual oxidation state. 1. 2. 4. 2 Stock System or Roman Numeral System The actual oxidation state is designated by a Roman Numeral placed in parenthesis immediately following the name of the metal. This is useful especially if the metal has more than two oxidation states. Formula |Classical System |Stock System | |CuCl |Cuprous chloride |copper(I) chloride | |CuCl2 |Cupric chloride |copper(II) chloride | |FeCl2 |ferrous chloride |iron(II) chloride | |FeCl3 |ferric chloride |iron(III) chloride | 1. 3. Compounds Named Like Binary Compounds Few other compounds take an –ide ending, like binary compounds. These include the following: |OH- |hydroxide |O22- |peroxide | |CN- |cyanide | |NH2- |amide | |I3- |triiodide | |N3- |azide | 1. 4. Trivial Names Some common binary compounds are assigned trivial names that have been assigned arbitrarily. These are universally used that they are allowed by the IUPAC rules of nomenclature. H2O |water | |NH3 |ammonia | |PH3 |phosphine | |AsH3 |arsine | 1. 5. Binary Acids A binary compound composed of hydrogen with a more electronegative element can act as a binary acid in water solution. For acids of this types, the prefix hydro- is added, and then the –ide ending is replaced by –ic acid. HF |hydroflouric acid | |HCl |hydrochloric acid | |HBr |hydrobromic acid | |HI |hydroiodic acid | 2. Ternary and Higher Compounds 2. 1 Oxyacids and Salts Oxyacids are composed of a nonmetal with more than one oxidation state, along with hydrogen and oxygen. A salt of oxyacid is formed when one or more of the hydrogen ions of an acid is replaced by a cation. The prefix hypo-, is used to denote the lowest oxidation state of the nonmetal with the characteristic ending –ous and the prefix per- is used to denote the highest oxidation state with the ending –ic. For acids whose names end in –ous, the corresponding salt ends with the suffix –ite, and those whose names ends in –ic, the name of the salt ends in –ate. Acid |Oxyanion |Salt | |H2SO3 |sulfurous acid |SO32- |sulfite |Na2SO3 |sodium sulfite | |H2SO4 |sulfuric acid |SO42- |sulfate |Fe2SO4 |iron(II) sulfate | |HClO |hypochlorous acid |ClO- |hypochlorite |NaClO |sodium hypochlorite | |HClO2 |chlorous acid |ClO2- |chlorite |KClO2 |potassium chlorite | |HClO3 |chloric acid |ClO3- |chlorate |NaClO3 |sodium chlorate | |HClO4 |perchloric acid |ClO4- |perchlorate |NaClO4 |sodium perchlorate | 2. 2 Salts of Polyprotic Acids These types of salts are formed when one or more hydrogen ions in polyprotic acids or acids with more than one replaceable H+ ion (e. g. , H2S, H3PO4, H2SO4) is/are replaced by metal ions. In naming, the word hydrogen is added to the name of the oxyanion. |NaH2PO4 |sodium dihydrogenphosphate |Na2HPO4 |disodium hydrogenphosphate | |Na3PO4 |trisodium phosphate | |NaHS |sodium hydrogensulfide | EXERCISES 1. Name the following. a. FeI2___________________________________ b. I2___________________________________ c. FeCl3___________________________________ d. Fe2(SO4)3___________________________________ e. FeS___________________________________ f. NCl3___________________________________ g. H2CO3___________________________________ h. CaCO3___________________________________ i.

Be2C___________________________________ j. SnSO4___________________________________ k. (NH4)2S___________________________________ l. N2O4___________________________________ 2. Write the correct chemical formula a. Barium chloride___________________ b. Stannous nitrate___________________ c. Stannic nitrate___________________ d. Aluminum carbide___________________ e. Magnesium phosphate___________________ f. Nitrogen dioxide___________________ g. Ammonium sulfate___________________ h. Barium carbonate___________________ i. Sodium carbonate___________________ j. Calcium hydrogen phosphate___________________ k. Disulfur dichloride___________________ 3. Complete the following table Formula |Name as acid |Formula of sodium |Name of salt | | | |salt | | |HNO3 | | | | |HNO2 | | | | |HBrO | | | | |HBrO2 | | | | |HBrO3 | |NaBrO3 | | |HBrO4 | | | | 4. Name the following as binary compounds or as salts from the anions of polyprotic or oxo acids. a. NaIO___________________________________ b. K2HPO4___________________________________ c. Na2SO3___________________________________ d. KMnO4___________________________________ e.

BaSO3___________________________________ f. FeSO4___________________________________ g. HClO3___________________________________ h. Na2SO4___________________________________ i. Fe(NO3)3___________________________________ j. Ca(ClO2)2___________________________________ 5. The spaces below represent portions of some of the main groups and periods of the periodic table. In the proper squares, write the correct formulas for the chlorides, oxides and sulfates of the elements of Groups 1, 2 and 3, respectively. Likewise, write the formulas of the compounds of sodium, calcium and aluminum with the elements of Groups 6 and 7. Two of the squares have been completed as examples. Period |Group 1 |Group 2 |Group 3 |Group 6 |Group 7 | |2 | LiCl | |(Omit sulfate) | | | | |Li2O | | | | | | |Li2SO4 | | | | | |3 | | | |Na2S | | | | | | |CaS | | | | | | |Al2S3 | | |4 | | | | | | | | | | | | | | | | | | | | |5 | | | | | | | | | | | | | | | | | | | | ExPERIMENT WATER OF HYDRATION Most solid chemical compounds will contain some water if they have been exposed to the atmosphere for any length of time.

In most cases the water is present in very small amounts, and is mere adsorbed on the surface of the crystals. Other solid compounds contain larger amounts of water that is chemically bound in the crystal. These compounds are usually ionic salts. The water that is present in these salts is called the water of hydration and is usually bound to the cations in the salt. In this experiment you will study some of the properties of hydrates. You will identify the hydrates in a group of compounds, observe the reversibility of the hydration reaction, and test some substances for efflorescence or deliquescence. Finally you will be asked to determine the amount of water lost by a sample of unknown hydrate on heating.

From this amount, if given the formula or the molar mass of the anhydrous sample, you will be able to calculate the formula of the hydrate itself. MATERIALS AND APPARATUS |watch glass |iron ring |crucible tongs | |test tubes |iron stand |triangular clay | |Bunsen burner |crucible |desiccators | PROCEDURE 1. Identification of Hydrates. Place about 0. g of the compounds listed below in small, dry test tubes, one compound to a tube. Observe carefully the behavior of each compound when you heat it gently with a burner flame. If droplets of water condense on the cool upper walls of the test tube, this is evidence that the compound may be a hydrate. Note the nature and the color of the residue. Let the tube cool and try to dissolve the residue in a few cm3 of water, warming very gently if necessary. A true hydrate will tend to dissolve in water, producing a solution with a color very similar to that of the original hydrate. If the compound is a carbohydrate, it will give off water on heating and will tend to char.

The solution of the residue in water will often be caramel colored. Nickel chloride Potassium chloride Sodium tetraborate (borax) Sucrose Potassium dichromate Barium chloride 2. Reversibility of Hydration. Gently heat a few crystals ~0. 3 g, of hydrated cobalt (II) chloride, CoCl2(6H2O, in an evaporating dish until the color change appears to be complete. Dissolve the residue in the evaporating dish in a few cm3 of water from your wash bottle. Heat the resulting solution to boiling (CAUTION! ), and carefully boil it to dryness. Note any color changes. Put the evaporating dish on the lab bench and let it cool. 3. Deliquescence and Efflorescence.

Place a few crystals of each of the compounds listed below on separate watch glasses and put them next to the dish of CoCl2 prepared in Part B. Depending upon their composition and the relative humidity (amount of moistures in air), the samples may gradually either lose water of hydration to, or pick up water from, the air. They may also remain unaffected. Any changes in crystal structure, color, or appearance of wetness should be noted. Observe the samples occasionally during the rest of the laboratory period. Since the changes tend to occur slowly, your instructor may have you compare your samples with some that were set out in the laboratory a day or two earlier. Na2CO3(10H2O (washing soda) CaCl2

KAl(SO4)2(12H2O (alum) CuSO4(5H2O 4. Percent Water in a Hydrate. Clean a porcelain crucible and its cover with 6 M HNO3. Any stains that are not removed by this treatment will not interfere with this experiment. Rinse the crucible and cover with distilled water. Put the crucible with its cover slightly ajar on a clay triangle and heat with a burner flame, gently at first and then to redness for about 2 minutes. Allow the crucible and cover to cool, and then weigh them to 0. 001 g on an analytical balance. Handle the crucible with clean crucible tongs. Obtain a sample of unknown hydrate from the stockroom and place about a gram of sample in the crucible.

Weigh the crucible, cover, and sample on the balance. Put the crucible on the clay triangle, with the cover in an off-center position to allow the escape of water vapor. Heat again gently at first and then strongly, keeping the bottom of the crucible at red heat for about 10 minutes. Center the cover on the crucible and let it cool to room temperature. Weigh the cooled crucible along with its cover and contents. Examine the solid residue. Add water until the crucible is two thirds full and stir. Warm gently if the residue does not dissolve readily. Does the residue appear to be soluble in water? DATA AND OBSERVATIONS A. Identification of Hydrates |Water appears |Color of residue |Water soluble |Hydrate | |Nickel chloride | | | | | |Potassium chloride | | | | | |Sodium tetraborate | | | | | |Sucrose | | | | | |Potassium dichromate | | | | | |Barium chloride | | | | | B. Reversibility of Hydration Summarize your observations on CoCl2(6H2O. Is the dehydration and hydration of CoCl2 reversible? C. Deliquescence and Efflorescence |Observation |Conclusion | |Na2CO3(10H2O | | | |CaCl2 | | | |KAl(SO4)2(12H2O (alum) | | | |CuSO4(5H2O | | | D. Percent water in a Hydrate |Mass of crucible and cover | | |Mass of crucible, cover, and solid hydrate | | |Mass of crucible, cover, and residue | | Mass of solid hydrate | | |Mass of residue | | |Mass of water lost | | |Percentage of water in the unknown hydrate | | |Formula mass of anhydrous salt (if furnished) | | |Number of moles of water per mole of unknown hydrate | | ExPERIMENT GASES

This experiment illustrates three of the common gas laws: Boyle’s law, Charles and Gay-Lussac’s law and Graham’s law. Boyle’s law states that the volume, V, of a certain quantity of an ideal gas is inversely proportional to its pressure, P, at a given temperature and amount of gas. Charles’ and Gay-Lussac’s law states that the volume of a gas is directly proportional to the absolute temperature, at a certain pressure and amount of gas. Graham’s law describes that the velocity of an ideal gas is inversely proportional to the square root of its molar mass. The first two gas laws will be validated using plots of the properties involved. Graham’s law will be determined by comparing the velocities of two sample gases. MATERIALS AND APPARATUS Glass syringe |250 or 400-mL beaker |Black cardboard | |Syringe holder |Iron stand |Ruler | |Thermometer |Iron ring |Graphing paper | |Glycerol |Wire gauze |Concentrated HCl | |Modeling clay |Glass tubing

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