Lubricating the glass with water or glycerine is helpful. When diluting acids, always add the acid to water, never water to the acid. Do not pour organic solvents down a sink in the open laboratory. Dispose of them as directed by your instructor, or down a drain in the fume hood. Flush with plenty of water. Do not wind the electric cord around a hot plate if it is still warm. The hot plate might melt the rubber insulation. Do not dispose of matches, paper, or solid chemicals in the sink.
Matches, after you are sure they are extinguished, and paper should be discarded into a wastebasket. Solid chemicals should be disposed of in whatever facility is provided in your laboratory. Do not put broken glassware into wastebaskets. Dispose of it in designated places. Have a nearby student call the instructor for aid. If you spill any chemical, solid or liquid, be sure to clean it up so another student does not come into contact with it and perhaps be injured by it.
Chemical characteristics, hazard levels, and safety instructions for the chemicals you use in the laboratory are described in Material Safety Data Sheets MSDS that are generally available in the laboratory. Follow directions given by your instructor in regard to these sheets.
Pay close attention to particular safety precautions your instructor talks about before you begin each experiment. Before leaving the laboratory, wipe the desk top and wash your hands with soap and water.
The following procedures will help minimize the possibility of contamination: 1. Avoid handling more than one reagent bottle at a time, so you do not interchange their stoppers by mistake. Introduction n Safety in the Laboratory 3 3. When selecting a reagent bottle, read the label twice to be sure you have the chemical you want.
Do not lay tops of reagent bottles or stoppers on the laboratory bench. Use separate spatulas to remove different solid chemicals from their bottles.
Never remove a liquid reagent from a stock bottle with an eye dropper. Pour a small portion into a clean, dry beaker, and use your eye dropper to remove the liquid from the beaker.
When a quantity of a chemical is removed from its original container, whether it is a solid or a liquid, do not return any excess to the stock bottle. Dispose of the unused portion as directed by your instructor. Never weigh a chemical directly on a balance pan. Use a preweighed container. Weighing paper is acceptable for most solid chemicals. Some chemicals react with some stoppers.
If you are going to store a chemical or solution in a bottle other than its original container, be sure the stopper you select glass, rubber, cork is suitable for that substance. Never leave a stock bottle uncovered. Be sure you cover the bottle with the proper cover. For beginning students this assumption is often wrong. This section discusses some of these methods.
The names and formulas of different chemicals may be almost identical. The names differ by one letter, and the formulas differ by 1 in a subscript. We strongly recommend that you read all chemical names and formulas twice in the laboratory manual and twice again on the supply bottle. When you need a chemical, take the container in which you will place it to the station from which the chemical is distributed.
Transfer the chemical to the container there. Do not take a supply bottle to your work area. Solid chemicals are generally distributed in wide-mouth, screw-cap bottles.
When you remove a cap from a bottle, place it on the desk with the top, or outside, of the cap down. This prevents contaminating the inside of the bottle from a dirty desk when the cap is returned to the bottle. Using a clean scoop, remove the amount of chemical you need. If you have removed too much chemical, do not return the excess to the bottle; instead throw it away.
The waste from this procedure is less of a problem than the contamination that will eventually occur if excess chemicals are returned to supply bottles.
It follows that you should judge your requirement carefully and take no more chemical than you need. After you have transferred the chemical you need, return the cap to the bottle and tighten it securely. If any solid has been spilled on the desk, clean it up before leaving the distribution station. Figure LP. This technique may also be used when pouring from a beaker, as shown in Figure LP. When using a dropper for removing the liquid, be sure to hold the dropper vertically with the rubber bulb at the top so that the liquid does not Figure LP.
The proper procedure is to pour some of the liquid into a small beaker and then use your dropper to transfer the liquid from the beaker to your container. Excess liquid should be thrown away, as noted above.
Estimate your needs carefully so the excess can be kept to a minimum. When working with such chemicals, it is best to work in a fume hood. In fact, trying to get that exact amount is a waste of time, both your time and the time of other students who will be delayed because you tie up a balance. While using an exact quantity of a chemical is not important, knowing as accurately as possible the quantity actually used is essential if that quantity becomes a part of your calculations.
Again the exact quantity is not important, and it is a waste of time to measure it with a graduated cylinder. One milliliter therefore can be estimated simply as two droppers-full. Volumetric laboratory equipment is calibrated to measure volume by sighting to the bottom of the meniscus, as shown in Figure LP. Notice that it is essential that the line of sight be perpendicular to the calibrated vessel if you are to read it accurately.
It is also important that you hold the vessel vertically. Four types of calibrated glassware are used in the experiments in this book. The most accurately calibrated are the volumetric pipets and burets used in Experiments 20 and Most of your volume measurements will be made in graduated cylinders. Their main purpose is to measure volumes and they are designed and calibrated accordingly. Volumes estimated by these calibrations should never be used in calculations.
No attempt will be made to duplicate those instructions here. Chemicals are never weighed directly on the pan of a laboratory balance. Instead, the mass is determined by a process known as weighing by difference. A suitable container—a small beaker, or perhaps a test tube that is to be used in the experiment—is weighed empty on the balance.
The desired chemical is added to the container, and the total mass of the combination is determined. Throughout this book the word container is used to include any and all objects that pass through the entire experiment unchanged in mass. Sometimes students use containers that are not actually part of the experiment in taking samples of solid chemicals.
Most common is the practice of placing a piece of paper on the pan of a balance, transferring the required quantity of chemical to the paper, and then transferring it to the vessel to be used in the experiment. Then transfer the chemical, and bring the paper back for a second weighing. This way your difference will be the mass of the chemical actually transferred, unaffected by any chemical that may have remained on the paper unnoticed.
In this method you should use a hard, smooth paper—waxed paper is best—rather than coarse paper, such as paper towel, which is certain to trap powders and tiny crystals. Laboratory balances are subject to corrosion. Both the balances and the balance area should be kept clean, and spilled chemicals should be cleaned up immediately. Close the side doors or hood of a milligram balance while weighing. Record all digits allowed by the accuracy of the balance used, even if the last digit happens to be a zero on the right side of the decimal point.
Weigh objects that are wet evaporation of water will change the mass. Weigh volatile liquids in uncovered vessels. Forget to check the zero on a milligram balance before and after weighing.
Forget to record the mass to as many digits as the accuracy of the balance allows—and no more. As a group the burners are called Bunsen burners, although most burners used today are improvements over the original Bunsen design. All have the same general features, and the Tirrill burner described in Figure LP. Gas enters the barrel of the burner from the center of the base, controlled by a valve in the base. Air enters through an opening at the bottom of the barrel where it screws onto the base.
When the barrel is screwed down, the air opening is small. You usually become aware of this condition by the noise produced in burning.
Be careful, however, because the barrel of a burner that is striking back becomes very hot. If you are not familiar with laboratory burners, it is recommended that you light one and experiment with the various adjustments to see how they work. It is also a rugged device that you will not damage without trying to. Second, heating would expand the glass and probably destroy the accuracy of the calibration. A rubber stopper with one or more holes is cut from the side to one of the holes.
The split stopper may then be held open, as shown in Figure LP. In this experiment you will examine some of the characteristics of matter and be introduced to some of the language of science in which these characteristics are described.
Physical properties are those characteristics of a substance that can be observed without changing the composition of the substance. Common physical properties are taste, color, odor, melting and boiling points, solubility, and density. Chemical properties describe the behavior of a substance when it changes its composition by reacting with other substances or decomposing into two or more other pure substances. The ability to burn and the ability to react with water are chemical properties.
Matter can undergo two types of changes, physical and chemical. Physical changes do not cause a change in composition, only in appearance. For example, when copper is melted, only a change of state occurs; no new substance is formed. In a chemical change, substances are converted into new products having properties and compositions that are entirely different from those of the starting materials. Wood, for example, undergoes a chemical change when it burns by reacting with oxygen in the air, forming carbon dioxide and water vapor as the new products.
When two liquids are mixed, the mixture may be completely uniform in appearance. In this case the liquids are said to be miscible. Some liquids are miscible in all proportions, while others have a limited range of 15 16 Introduction to Chemical Principles: A Laboratory Approach n Weiner and Harrison miscibility. If the two liquids are not at all miscible, i.
When a solid is added to and dissolves in a liquid, it is soluble in that liquid. The mixture formed is called a solution. A liquid solution is always clear; it may be colorless, or it may have a characteristic color. If the solid does not dissolve, it is said to be insoluble. When two solutions are combined, a chemical change, or reaction, may occur in which new products form. If so, it will be evidenced by one of several visible changes. Among them are: 1.
Formation of a precipitate, or a solid product. If allowed to stand, the precipitate will settle to the bottom of its container. Formation of a gaseous product. The gas produced bubbles out of the solution, a process called effervescence. Occurrence of a color change. Usually a color change indicates the formation of a product with a color not originally present among the reactants.
Sometimes the color will be the same as that of one of the reactants, but a darker or lighter shade. In many cases no reaction occurs when two solutions are brought together. Skin contact with these three liquids, or with hydrochloric acid, should be avoided. If it occurs, rinse the affected area thoroughly with water, and then wash with soap and water. Be sure to wear approved eye protection throughout the experiment. Trichloroethane and xylene mixtures should be collected in stoppered bottles.
Do not pour them down the drain. Solutions containing heavy metal precipitates should be collected in a separate container. Mixing Liquids A. Place about 20 drops of trichloroethane into a small-size test tube. Note this approximate quantity, because you will have several occasions in the experiment to estimate this volume in a test tube. Add about 10 drops of water and gently shake the test tube, or mix the contents with a stirring rod. Are the two liquids miscible?
Record your observation on the work page. If the two liquids are not miscible, identify the liquid, trichloroethane or water, that is on top. You may determine which liquid is on top by the relative quantities placed into the test tube: you added two Experiment 1 n Properties and Changes of Matter 17 times as much trichloroethane as you did water.
Record the name of the top liquid on your work page. If the liquids are not miscible, again record on the work page which liquid is on top. Using a clean test tube, or the original one thoroughly rinsed with water, repeat the procedure with about 20 drops each of methanol methyl alcohol and water.
It is not necessary this time or hereafter to reverse the order of liquids, as in Parts 1A and 1B. Again record your observations and conclusions. Using a clean and thoroughly rinsed test tube, repeat the procedure again, this time with about 20 drops each of water and xylene.
Record your observations and conclusions as before. Using a clean and dry test tube there must be no water present , perform the experiment once again, now using about 20 drops each of trichloroethane and xylene. Record your observations. Dissolving a Solid in a Liquid In this part of the experiment and the next, you will be preparing solutions.
Shake the test tube gently, or stir the contents with a clean, dry stirring rod. If none of the solid appears to dissolve, the substance is insoluble. If any of it dissolves, but a small amount does not, add more water to get all of the solid into the solution.
Place a small quantity of barium chloride, BaCl2, in water as described above. Does the solid dissolve in the water?
Record your observations and save the solution for further use. Add a small amount of sodium sulfate, Na2SO4, to about 10 mL of water in a second test tube. Does the solid dissolve? Record your observations and save the solution. Combine the contents of the test tubes from Steps 2A and 2B in a large test tube. Set the test tube aside for 5 to 10 minutes and examine it again.
Record what you see. Add a small amount of barium sulfate, BaSO4, to about 10 mL of water. Is this compound soluble? Mixing Solutions A. In another test tube, dissolve a small amount of potassium thiocyanate, KSCN, in about 2 mL of water. Mix the two solutions and record your observations. In a small test tube, dissolve a small amount of sodium chloride, NaCl, in about 2 mL of water. In a small test tube, dissolve a small amount of sodium carbonate, Na2CO3, in about 2 mL of water.
Add 2 to 3 drops of hydrochloric acid, HCl, watching carefully for any evidence of a chemical reaction. Then add some more HCl and watch for a reaction. In a small test tube, dissolve a small amount of calcium chloride, CaCl2, in about 2 mL of water. In another test tube, dissolve a small amount of sodium carbonate, Na2CO3, in about 2 mL of water.
Add concentrated ammonia solution, NH3 aq , to it, a drop at a time. Ammonia solutions are sometimes labeled ammonium hydroxide, NH4OH. Experiment 1 Advance Study Assignment 1. Distinguish between physical and chemical properties. Give an example of each. Classify each of the following as a physical or chemical change: a. Identify three forms of evidence that a chemical reaction has occurred: a. Is it possible from Parts 1A—1D to determine which of the liquids, xylene or trichloroethane, is more dense?
If so, identify the liquid with the greater density and explain how you reached your conclusion; if not, explain why. Is it possible from Part 1E alone to determine which of the liquids, xylene or trichloroethane, is more dense?
Mixture of contents of test tubes from 2A and 2B: Immediate appearance: Appearance 5—10 minutes later: 2D. In fact, everything that surrounds us is a chemical. The air we breathe is a mixture of gaseous elements, oxygen, and nitrogen. The food we eat contains carbohydrates, proteins, large molecular mass vitamins, etc. You may not think about it, but you handle all kinds of chemicals every day. Qualitative analysis is used to identify components of a solution or solid.
A reagent that causes an easily recognized reaction with a particular ion present is added to a sample of the unknown. If the reaction occurs, the ion is present; if the reaction does not occur, the ion is absent. In this experiment you will perform tests on known compounds that show the presence of certain ions. Finally, you will be given an unknown compound.
Any spilled acid or base should be washed off promptly. Be sure to wear goggles or safety glasses while performing this experiment. Dispose of excess solids as directed by your instructor. Solutions containing precipitates should be collected in a waste container. A white precipitate of silver chloride, AgCl, will form. Place a few crystals of table salt in a small test tube.
Dissolve the solid in about 10 drops of deionized water. Add 2 drops of 1 M nitric acid, HNO3. Then add 2 or 3 drops of 0. Garden Fertilizers Some of the active ingredients of ordinary garden fertilizers are ammonium salts. These compounds are the source of nitrogen, an element essential for the growth of plants. This can be detected by a piece of red litmus paper that has been moistened with deionized water. The paper will turn blue if NH3 is present. Pour about 10 drops of 1 M ammonium chloride, NH4Cl, into a small test tube.
Add about 10 drops of 3 M sodium hydroxide, NaOH. Hold a piece of moist red litmus paper in the mouth of the test tube. Do not allow the paper to come into contact with the side of the test tube, since it may have NaOH on it. If you notice no change, gently warm the test tube in a hot-water bath and check with litmus paper again. Place a small amount of garden fertilizer into a test tube.
Add about 10 drops of 3 M NaOH to the solid. Hold a moist strip of red litmus paper in the mouth of the test tube. It is commonly used to prepare soothing baths, and it is sometimes used as a purgative.
Pour about 10 drops of 1 M sodium sulfate, Na2SO4, into a small test tube. Place a small amount of Epsom salt into a test tube. When this compound reacts with an acid, gaseous carbon dioxide is produced.
The formation of bubbles indicates that at least one of these ions is present in a sample. Using a spot plate, dissolve a small amount of baking soda in about 10 drops of deionized water. Add about 6 to 8 drops of 1 M hydrochloric acid, HCl. Again record your observations, noting any difference between the two reactions. Take another small amount of solid baking soda and add 10 drops of commercial vinegar to it.
Record your observations, noting any difference between this reaction and the ones in Step 4A. Detergents One of the common ingredients of laundry detergents and wall-washing compounds is sodium phosphate, Na3PO4.
A yellow, powdery precipitate will form. Sometimes gentle heating in a water bath is necessary to hasten the reaction. Test by dipping a stirring rod into the solution and touching the wet rod to a strip of blue litmus paper. The solution is acidic if the color changes to red. Then add 6 to 8 drops of 0. Repeat the procedure with a small amount of laundry detergent. Dissolve the solid in about 10 drops of deionized water, acidify it, and add the molybdate reagent.
Heat the solution in a water bath and note any change, if there is one. Identification of an Unknown Obtain a solid unknown and record its number on your work page. Keep enough unknown to make three additional tests, in case you wish to repeat one or more of the procedures. Your unknown will contain only one ion. Identify the ion and record it on your work page. Record it as such. This page intentionally left blank Name Experiment 2 Advance Study Assignment 1.
What would happen if you used tap water instead of deionized water in this experiment? You have a piece of chalk. How would you determine if it contains CaCO3? If you have a white powder that could be either Epsom salt or table salt, how could you decide which one it is? Why is it dangerous to mix household chemicals indiscriminately?
A solvent mixture, carrying the pigments, was allowed to pass through a glass column packed with chalk. At the end of the experiment, the pigments were separated in colored bands at various distances from the starting level. This method is now known as column chromatography.
Chromatography may now be applied to colorless compounds and to ions. Paper chromatography is a more recent and much faster separation technique than column chromatography.
The mixture of solvents used to carry the substances along the paper is called the mobile phase, or solvent system. In practice, a sample of the solution containing the substances to be separated is dried on the paper. The end of the paper is dipped into the solvent system so that the sample to be analyzed is slightly above the liquid surface.
As the solvent begins to soak the paper, rising by capillary action, it transports the sample mixture upward. Each component of the mixture being separated is held back by the stationary phase to a different extent.
Also, each component has a different solubility in the mobile phase and therefore moves forward at a different speed. A combination of these effects causes each component of the mixture to progress at a different rate, resulting in separation.
The RF value is a characteristic property of a species, just as the melting point is a characteristic property of a compound. Each ion forms a different colored complex when sprayed with a solution containing potassium hexacyano ferrate II , K4[Fe CN 6]. Fumes of acetone and concentrated hydrochloric acid are objectionable and, to some degree, harmful. These chemicals should be used in the hood.
Be sure to wear safety glasses. Using a graduated cylinder and working in the hood, prepare the following solvent system: 19 mL acetone; 4 mL concentrated hydrochloric acid, HCl; 2 mL water or use 25 mL of a pre-prepared solvent mixture, if available.
This allows the atmosphere within the beaker to become saturated with solvent vapor and helps to give a better chromatographic separation. Obtain a piece of chromatography paper 24 to 25 cm long by 11 to 14 cm wide. Draw a pencil line about 1 cm from the long edge of the paper. You must use an ordinary pencil for this line. Ink or colored pencil often contains substances that may be soluble in the solvent, producing chromatograms of their own.
This line will indicate the origin see Figure 3. Also draw a line about 1 cm long, 6 cm above the center of the penciled line. Using a different capillary tube for each solution do not mix them! Apply the spots evenly over the line, leaving a margin of about 3 cm from each short edge of the paper. Use a separate, clean capillary tube for each solution; or, if the solutions are to be Experiment 3 n Separation of Cations by Paper Chromatography 37 obtained from beakers in which a capillary tube is provided, be sure to return the tube to its proper beaker.
With a pencil, identify each spot by writing on the paper directly beneath the spot. The solutions are: A. Any of the unknowns furnished be sure to record its number F. Another unknown again, record the number. Dry the paper under a heat lamp or air blower. Form the paper into a cylinder without overlapping the edges. Fasten the paper with staples, as shown in Figure 3.
Place the beaker in a position on your desk where it will remain undisturbed throughout this step. Taking care that the origin line remains above the solvent, carefully place the cylinder into the beaker, Chromatography paper Line about 1 cm long 6 cm Pencil line origin Figure 3.
Do not move the beaker or the solvent front will be uneven. NOTE: In this and all remaining steps, when the paper is wet, be sure not to lay it down on any surface that is not clean. When the solvent has risen above the short line drawn 6 cm above the origin in Step 2, remove the cylinder from the beaker and quickly mark the solvent front position with a pencil. Remove the staples and dry the paper under a heat lamp. Measure and record in millimeters the distance between the origin and the solvent front X in Figure 3.
Next, measure and record the distance between the origin and the center of each spot in the chromatograms for solutions A through D. Calculate the RF value for each ion, using Equation 3. Record that value as a decimal fraction to two decimals i. From the spots above your solutions E and F, indicate by a check mark in the table the ions present in the unknowns. If the center of a spot is right on the solvent front, the RF value equals 1. Solvent front X Y Figure 3. Experiment 3 Advance Study Assignment 1.
What would you observe if you used a ballpoint pen, instead of a pencil, to mark the chromatography paper? Why do you have to cover the beaker while the solvent is moving up the paper? What problem would be caused by moving the beaker during the development of the chromatogram? Solution Unknown No. To determine the density of a substance, you must measure both the mass and volume of the same sample of the substance.
Density is then calculated by dividing the mass by the volume, as indicated in Equation 4. Mass is measured by the usual weighing techniques.
The volume of a liquid may be measured in a graduated cylinder. The dimensions of a solid with a regular geometric shape rectangular block, cylinder, sphere may be measured with a ruler, and these measurements can then be used to calculate the volume. The volume of a solid with an irregular shape may be determined by measuring the volume of a liquid displaced when the solid is immersed in the liquid.
In Part 1 of this experiment you will be asked to determine experimentally the density of a known substance and then to calculate the percent error in your determination. Error is expressed as an absolute value, i. Absolute value is indicated by enclosing the quantity between vertical lines.
Thus Equation 4. This uncertainty dictates that all liquids be regarded as potentially dangerous and treated accordingly.
This includes the known liquid, trichloroethane. Liquid samples should be obtained from a dispensing station in the hood. If taken from the hood, liquids should be in containers that are stoppered or covered with a plastic sheet or metal foil. Avoid contact between all liquids and your skin; if it occurs, wash the exposed area thoroughly with soap and water. Safety glasses must be worn at all times. Depending on the nature of your liquid unknown, disposal directions will be given by your instructor.
Trichloroethane should be collected in a stoppered bottle. Length measurements are to be recorded in centimeters to the nearest 0. Record liquid volume measurements in milliliters to the nearest 0. Density of a Liquid A. Your mL graduated cylinder and a piece of plastic wrap e. Being sure the cylinder is clean and dry, weigh it and the Saran wrap—the container—to the nearest 0.
Take the cylinder and plastic wrap to the hood. Pour 12 to 15 mL of 1,1,1-trichloroethane into the cylinder; do not attempt to make the amount exactly 12, 13, 14, or 15 mL. Cover the cylinder with the plastic wrap. Estimate the volume to the nearest 0.
After making sure the outside of the cylinder is dry, measure and record the mass of the container plus liquid on the centigram balance. Dispose of your liquid as directed by your instructor. Experiment 4 n Densities of Liquids and Solids 47 E. Determine and record its mass to the nearest centigram. Make whatever measurements may be necessary to calculate the volume of the object, listing these measurements to the closest 0.
Because these objects are of various shapes, the data table contains blank spaces in which to describe the shapes and identify the measurements length, diameter, etc. Density of an Irregular Solid A.
Place 20 to 25 mL of water into the cylinder from Part 1. Record the volume to the nearest 0. Determine the mass of the cylinder plus water to the nearest centigram. This is the mass of the container for Part 3.
Place enough of the solid into the graduated cylinder to cause the liquid level to rise by more than 10 milliliter markings. Be sure all of the solid is below the surface of the liquid. Also measure the mass of the container and its contents to the nearest centigram.
Dispose of your solid material into the recovery facility that has been set up in your laboratory. Be careful not to mix unknown solids. Repeat Steps 3A and 3B for as many unknowns as are required by your instructor, or for a second run with the same unknown. Density of a Liquid Find the mass of the liquid by difference—by subtracting the mass of the container from the mass of the container plus liquid. The density of the liquid is found by dividing the mass of the liquid sample by its volume, as indicated in Equation 4.
Percent error may be calculated by substituting into Equation 4. The accepted value for the density of 1,1,1-trichloroethane is 1. The volume of a cylinder is the area of the base times the height. Density of an Irregular Solid Both the mass and the volume of the sample are found by difference.
Density is again calculated by substitution into Equation 4. Experiment 4 Advance Study Assignment 1. The volume of an unknown liquid is Calculate the density of the liquid. When Calculate the density of the metal. The accepted value for the density of a certain metal is 5.
Calculate the percent error in a laboratory experiment that yields a value of 5. Calculation Setups for Density Determinations: Name The atoms are held together by chemical bonds. It has been shown experimentally that the ratio of moles of the elements in a compound is nearly always a ratio of small, whole numbers. The few exceptions are known as nonstoichiometric compounds.
The formula containing the lowest possible ratio is known as its simplest formula. It is also called the empirical formula. At times it may be the same as the molecular formula; often, however, the molecular formula is an integral multiple of the simplest, empirical formula. For example, the simplest formula of the compound benzene C6H6 is simply CH, indicating that the ratio of carbon atoms to hydrogen atoms is one to one.
From these data the moles of atoms of each element may be calculated. By dividing these numbers by the smallest number of moles, you obtain quotients that are in a simple ratio of integers, or are readily converted to such a ratio. The ratio of moles of atoms of the elements in a compound is the same as the ratio of individual atoms that is expressed in the empirical formula.
In Option 1 you will react a measured mass of copper with excess sulfur. The excess sulfur is burned away as sulfur dioxide. In Option 2 the reaction is between a measured quantity of tin and excess nitric acid.
The excess acid is boiled off. Option 3 involves the reaction of a measured mass of magnesium with excess oxygen from the air. If so, plan your use of time. The procedure includes some periods in which you wait for a crucible to cool.
Calculate the simplest formula of the oxide from the following data: Mass of container Be careful not to reach through one in reaching for some object behind it. Be sure to use crucible tongs when handling hot crucibles, including the lid. Laboratory hardware gets hot, too. Harmful gases are released in Options 1 and 2. These reactions must be performed in a fume hood, as stated in their procedures.
Be careful of hot chemical spattering from crucibles when they are heated. Be sure to wear goggles throughout this experiment. Dispose of any solid residue as directed by your instructor. The purpose of this step is to remove moisture from the crucible. Support a clean, dry porcelain crucible and its lid on a clay triangle, as shown in Figure 5. Set the crucible and lid aside on a wire gauze to cool. When the crucible and lid are cool to the touch, weigh them on a centigram balance.
Record this value as the mass of the container. Place a loosely rolled ball of copper wire or medium shavings, about 1. Weigh them, with the lid, on a centigram or milligram balance, and record the weight as the mass of the container plus metal. Sprinkle about 1 to 1. Place the lid on the crucible and begin heating it in a fume hood.
Finally, heat the crucible strongly for about 5 minutes, making sure that no excess sulfur is present on the lid or on the sides of the crucible. Set the container and its contents aside to cool. Do not open the lid until the crucible is cool, because air oxidation is apt to occur. When the crucible is cool, lift the lid and examine the contents. There should be no evidence of sulfur in the crucible or on the lid.
If sulfur is present, heat the crucible again until the sulfur is completely burned away. Allow the crucible to cool. Weigh the container and its contents again. Record your measurement as the mass of the container plus compound. Set the container aside while you complete your calculations. Just before discarding the compound, press it to the bottom of the crucible. Notice the difference between the physical properties of the compound and those of the elements from which it was formed. The compound should be discarded as directed by your instructor.
Crucible Clay triangle Figure 5. Heat a porcelain crucible and its lid as described in Option 1, Step A. Allow it to cool. Weigh as described in Option 1, Step B. Place a loosely rolled ball of tin foil, weighing 1 to 1.
Weigh the crucible, the lid, and the metal on a centigram balance. Record the mass obtained. Under the fume hood, add concentrated nitric acid, HNO3, drop by drop, to the crucible until all the tin has reacted and a damp white paste remains.
Cool the crucible and compound to room temperature and weigh it. Record the mass of the container and compound. Keep your compound in the crucible until all calculations are completed. This may save you time if it becomes necessary to add more nitric acid. Weigh the crucible and lid on a milligram balance.
Record this value on the work page as mass of the container. Place a loosely folded magnesium ribbon, weighing 0.
Weigh the crucible, lid, and metal on a milligram balance and record the mass. Remove the lid and hold it near the crucible with a pair of tongs. Start heating the crucible, and as soon as the magnesium begins to burn, replace the lid. To convert the possible side product, magnesium nitride, to the oxide, let the crucible cool, add 10 drops of deionized water to it, and then gently heat to vaporize excess water.
Allow the crucible to cool, then weigh the cool crucible, lid, and product on a milligram balance. Experiment 5 Advance Study Assignment 1. Circle one of the following formulas that is correctly written as an empirical formula: NaSO1.
If the product weighs When the reaction is complete, the product has a mass of 3. What mass of sulfur should be used in the simplest formula calculation? It is referred to as water of hydration. The formula of a hydrate consists of the formula of the anhydrous without water compound followed by a dot, the number of water molecules that crystallize with one formula unit of the compound, and the formula of water.
Generally, water of hydration can be driven from hydrates by heating, leaving behind the anhydrous salt. The process may be accompanied by physical changes, such as a change in color or physical appearance. In this experiment you will be instructed to determine the mass of a sample of an unknown hydrate by difference, using a preweighed crucible as the container. From this you can calculate the moles of anhydrous compound in the original sample.
From the mass of water in the original sample you can calculate the moles of water. You therefore repeat the heating, cooling, and weighing procedure. Another heating is therefore required. The heating, cooling, and weighing sequence is repeated until two successive duplicate weighings are recorded. Your instructor may require you to perform the experiment twice to obtain duplicate results or to run more than one unknown. The procedure includes some periods in which you must wait for a crucible to cool.
In this way you perform both runs simultaneously. Be sure to use crucible tongs in handling hot crucibles, including the lids.
Goggles should always be worn when working with chemicals, and particularly while heating them, as in this experiment.
The anhydrous solids should be discarded in a container or you should follow directions given by your instructor. See Figure 5. Heat the container to constant mass in 5-minute heating cycles until duplicate masses within 0. Place 1 to 1. Record as the mass of the container plus hydrate. Cool the crucible and lid, and when they feel cool to the touch, weigh them again.
Heat to constant mass in 4- to 6-minute heating cycles until duplicate masses are reached. Set the container and its contents aside while you complete your calculations. Behavior of a Hydrate A. Holding the test tube tilted at an angle, and with its mouth pointing away from you and all others, heat the test tube gently.
After the test tube has cooled to room temperature, add a few drops of water. Hold the test tube against the back of your hand. From these, calculate the percentage water of hydration, using Equation 6.
Return to the original data masses of container, container plus hydrate, and container plus anhydrous salt , and identify and perform the subtraction that will yield the mass of the anhydrous salt. Also convert grams of water of hydration to moles. From the ratio, determine the formula of the hydrate. Experiment 6 Advance Study Assignment 1. How can you make sure that all of the water of hydration has been removed? Why do you have to cool the crucible before weighing it?
Calculate the percentage water of hydration and the formula of the hydrate if the residue after heating weighed 2. Experiment 6 Work Page Data and result tables for Experiment 6 appear on pages 78 and Observations when a hydrate is heated in a test tube. Observations when water is added to an anhydrous salt.
You will compare your experimental result with the theoretical percentage calculated from the formula KClO3. While potassium chlorate decomposes simply by heating, the reaction is intolerably slow.
A catalyst, manganese dioxide, MnO2, is therefore added to speed the reaction. Although it contains oxygen, the catalyst experiences no permanent change during the reaction and does not contribute measurably to the amount of oxygen generated. As with all catalysts, the quantity present at the end of the reaction is the same as the quantity at the beginning.
The experimental procedure is to weigh a quantity of potassium chlorate, heat it to drive off the oxygen, and then weigh the residue, which is assumed to be potassium chloride. The loss in mass represents the oxygen content of the original potassium chlorate. The above procedure will be carried out in a test tube. It will include the constant mass of catalyst that remains in the test tube throughout the experiment, plus whatever device is used to hold the test tube and its contents while they are being weighed.
If your milligram balance has a suspended pan and there is provision for hanging an object to be weighed, you can clamp the test tube in a test-tube holder and hang the entire assembly from the hook provided. In this case the test-tube holder is a part of the container.
If the pan on your milligram balance is supported only from beneath, you can stand your test tube in a small 85 86 Introduction to Chemical Principles: A Laboratory Approach n Weiner and Harrison beaker each time it is weighed, and include the beaker in the mass of the container. You must be sure, of course, to use the same beaker for each weighing. It is possible that some oxygen may still be present after the second heating, too.
The procedure is therefore repeated again, as many times as necessary, until there is negligible change in mass no more than 0. The formation of a gas at the bottom of a test tube may result in a sudden expansion, blowing hot chemicals out of the test tube.
This will not occur if the test tube is handled properly during heating. When heating a solid in a test tube, tip the tube until it is almost horizontal and tap it carefully until the contents are distributed over the lower half of the length of the tube, as shown in Figure 7.
Do not concentrate the heat in any one area, particularly near the bottom of the test tube. Be very sure your test tube is not pointing toward anybody, including yourself, while it is being heated. Be aware of what those around you are doing while this experiment is being performed in the laboratory. Obviously, wearing goggles is absolutely mandatory while you or anyone near you is performing this experiment.
Dispose of the residue in the designated container. Place 0. Heat the test tube over a Bunsen burner for about 3 to 4 minutes to drive off any moisture that may be present in the catalyst and test tube.
When the tube is cool to the touch, measure the mass of the entire container on a milligram balance. Record the mass in the space provided. Buy this product. K educators : This link is for individuals purchasing with credit cards or PayPal only. Newly updated based on extensive reviewer feedback, this introductory text remains focused on the essentials necessary for success in General Chemistry. Introduction to Chemistry Principles , Eleventh Edition focuses on the most important topics — omitting organic and biochemistry chapters — and teaches the problem-solving skills students need.
Each topic is introduced and developed step by step until reaching the level of sophistication required for further course work.
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Introduction to Chemical Principles, 11th Edition. Stephen Stoker. Description Newly updated based on extensive reviewer feedback, this introductory text remains focused on the essentials necessary for success in General Chemistry.
Emphasis on problem solving throughout uses dimensional analysis whenever possible. This equips students with a powerful and widely applicable tool that requires no mathematics beyond arithmetic and elementary algebra. Worked examples with detailed commentary show students the proper way to mentally dissect and solve a problem. Over 5, questions and problems give students more opportunities than any other text to become proficient problem-solvers.
Emphasis on significant figure concepts in all problem-solving situations provides two answers to the example: the calculator answer which does not take into account significant digits , and the correct answer which is the calculator answer adjusted to the correct number of significant figures. The Human Side of Chemistry vignettes are brief biographies of scientists who helped develop the foundations of modern chemistry.
Chemical Insights add perspective to worked-out examples that center on specific compounds. It focuses on the chemical compound itself, its relationship to the environment, its relationship to living systems biochemistry , etc. Answer Double Check , found in the majority of worked-out problems, encourages students to consider if their solution is a "reasonable answer" in terms of numerical magnitude, number of significant figures present, sign convention plus or minus and direction of change increase or decrease.
Multiple-Choice Practice Test questions in the end-of chapter reviews help students prepare for exams. New to This Edition.
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