This Document Contains Experiments 12 to 18 Name____________________________________________________Section________________Date___________ Experiment 12: Magnetic Fields Invitation to Inquiry Is it possible to produce electricity in an extension cord that is not plugged into a circuit? Hook a 50 ft extension cord to a galvanometer and move it as a jump rope, cutting the magnetic field lines around the earth. Figure out how you are going to attach the cord to the galvanometer and how you are going to move it across the earth’s magnetic field lines. Can you think of any practical uses for “jump-rope electricity?” Background magnet moved into the space nears a second magnet, experiences a force as it enters themagnetic field of the second magnet. The magnetic field model is a conceptual way of considering how two magnets interact with one another. The magnetic field model does not consider the force that one magnet exerts on another one through a distance. Instead, it considers the condition of space around a magnet. The condition of space around a magnet is considered to be changed by the presence of the magnet. Since this region of space, or field, is produced by a magnet, it is called a magnetic field. A magnetic field can be represented by magnetic field lines. By convention, magnetic field lines are drawn to indicate how the north pole of a tiny imaginary magnet would point when in various places in the magnetic field. Arrowheads indicate the direction that the north pole would point, thus defining the direction of the magnetic field. The strength of the magnetic field is greater where the lines are closer together and weaker where they are further apart. Magnetic field lines emerge from a magnet at the north pole and enter the magnet at the south pole. Magnetic field lines always form closed loops. Magnetic field strength is defined in terms of the magnetic force exerted on a test charge of a particular charge and velocity. The magnetic field is thus represented by vectors (symbol B) which define the field strength, also called the magnetic induction. The units are: = newton meters second Since a coulomb(s) is an amp, this can be written as (coulomb)( ) B = amp meternewton⋅ which is called a tesla (T). The tesla is a measure of the strength of a magnetic field. Near the surface, the earth’s horizontal magnetic field in some locations is about 2 × 10-5 tesla. A small bar magnet produces a magnetic field of about 10-2 tesla, but large, strong magnets can produce magnetic fields of 2 tesla. Superconducting magnets have magnetic fields as high as 30 tesla. Another measure of magnetic field strength is called the gauss (G) (1 tesla = 104 gauss). Thus, the process of demagnetizing something is sometimes referred to as “degaussing.” In this experiment you will investigate the magnetic field around a permanent magnet. Procedure Tape a large sheet of paper on a table, with the long edge parallel to the north-south magnetic direction as determined by a compass. Center a bar magnet on the paper with its south pole pointing north. Use a sharp pencil to outline lightly the bar magnet on the paper. Write N and S on the paper to record the north-seeking and south-seeking poles of the bar magnet. Place the bar magnet back on its outline if you moved it to write the N and the S. Slide a small magnetic compass across the paper, stopping close to the north-seeking pole of the bar magnet. Make two dots on the paper, one on either side of the compass and aligned with the compass needle. See Figure 12.1. Slide the compass so the south pole of the needle is now directly over the dot that was at the north pole of the needle. Make a new dot at the north pole end of the compass, exactly in front of the needle. See figure 12.2. Continuing the process of moving the compass so the south pole of the needle is over the most recently-drawn dot, then making another new dot at the north pole of the needle. Stop when you reach the bar magnet or the edge of the paper. Draw a smooth curve through the dots, using several arrowheads to show the direction of the magnetic flux line. Repeat procedure steps 3 through 6, by starting with the compass in a new location somewhere around the bar magnet. Repeat the procedures until enough flux lines are drawn to make a map of the magnetic field. Place a second large sheet of paper on a large rigid plastic sheet (or glass plate) on top of the bar magnet. Sprinkle a thin, even layer of iron filings over the plastic, tapping the sheet lightly as you sprinkle. Sketch the magnet flux lines on the paper as shown by the arrangement of the filings. Results How is using iron filings (a) similar, and (b) different from using a magnetic compass to map a magnetic field? They have the same shape. In terms of a force, or torque on a magnetic compass needle, what do the lines actually represent? Explain. The lines represent the direction of the magnetic field at any position. Do the lines ever cross each other at any point? Explain. No -- it is not possible for the magnetic field at any point to have more than one direction. Where do the lines appear to be concentrated the most? What does this mean? Near the ends of the magnet -- this is where the magnetic field is the strongest. Name____________________________________________________Section________________Date___________ Experiment 13: Reflection and Refraction Invitation to Inquiry Boil water, cool, then freeze it in various shapes of lenses. Boiling the water drives the airout of solution, making clear ice. Use the ice lenses to ignite paper by focusing sunlight on a piece of paper. Make a prism by fastening three glass microscope slides together to form a triangle. Sealthe edges and end with cellophane tape, fill it with water, then seal the other end. You can also seal the two ends by pushing them into a block of wax. Use the prism to experiment with sunlight and strong sources of light. Fill a small aquarium with water that has been made murky by mixing in chalk dust,talcum powder, or some other insoluble substance that will make a beam of light visible. Experiment with aiming a strong beam of light through the side of the aquarium and up to the inner surface, directing the beam at increasing angles. Background The travel of light is often represented by a light ray, a line that is drawn to represent the straight-line movement of light. The line represents an imaginary thin beam of light that can be used to illustrate the laws of reflection and refraction, the topics of this laboratory investigation. A light ray travels in a straight line from a source until it encounters some object. What happens next depends on several factors, including the nature of the material making up the object, the smoothness of the surface, and the angle at which the light ray strikes the surface. If the surface is perfectly smooth, rays of light undergo reflection. If the material is transparent, on the other hand, the light ray may be transmitted through the material. In these cases the light ray appears to become bent, undergoing a change in the direction of travel at the boundary between two transparent materials (such as air and water). The change of direction of a light ray at the boundary is called refraction. Light rays traveling from a source, before they are reflected or refracted, are called incident rays. If an incident ray undergoes reflection, it is called a reflected ray. Likewise, an incident ray that undergoes refraction is called a refracted ray. In either case, a line perpendicular to the surface, at the point where the incident ray strikes, is called the normal. The angle between an incident ray and the normal is called the angle of incidence. The angle between a reflected ray and the normal is called the angle of reflection. The angle between a refracted ray and the normal is called the angle of refraction. These terms are descriptive of their meaning, but in each case you will need to remember that the angle is measured from a line perpendicular to the surface, the normal. Procedure Part A: Reflection of Light Using a ruler, draw a straight line across a sheet of plain (unlined) white paper. Place the paper on a smooth piece of cardboard that has been cut from a box. Label the line with a B at one end and B′ at the other end (B is for boundary). Attach a small, flat mirror to a block of wood as shown in Figure 13.1. Place the mirror and block combination on the paper with the back of the mirror (the reflecting surface) on line BB′. Stick a pin straight up and down into the paper about 10 cm from the mirror and slightly to the right side as shown in Figure 13.1. On the left side, carefully align the edge of a ruler with the reflected image as shown in the illustration. Then firmly hold the ruler and draw a pencil line along this edge. Move the mirror and extend this line to the mirror boundary line BB′. Label the point of reflection with the letter P. Place a protractor on line BB′ and mark a point 90˚ from the line. From this point, use the ruler to draw a dashed normal (NP). Complete your ray diagram by using the ruler to draw a line from the point of reflection (P) to the source of the light ray at the pin (I). Place arrows on line IP and line PR to show which way the light ray moved. Figure 13.1 Use the protractor to measure the angle of incidence and the angle of reflection. Record these angles in Data Table 13.1 on page 112 under Trial 1. Place the mirror with its back edge again on line BB′ and conduct a second and third trial at different sighting angles. Record these measurements in Data Table 13.1. Part B: Refraction of Light Place a clean sheet of white (unlined) paper on the cardboard. Place a glass plate approximately 5 cm square flat on the center of the paper. Use a pencil to outline the glass plate, then move the plate aside. Figure 13.2 Use a ruler to draw a straight line from the upper left edge of the plate outline, making an angle of about 60˚ to the edge. Label this line IB as shown in Figure 13.2. Place one upright pin at point B immediately outside the plate outline and a second upright pin at point I. Return the glass plate to the outline. Bring the cardboard, paper, and glass plate near the edge of the table so you can sight through the glass plate toward the two pins. Position a ruler so that one edge aligns with the two pins as shown in Figure 13.2. Draw a line along the ruler and label the line B′R. Move the glass plate aside for a second time. Draw a line from B to B′, showing the path of the light ray through the glass. Overall, the path of the light ray is from IB to BB′ to B′R, showing that the light ray was bent twice. Draw normals to the surface of the glass at B and B′. Show the angle of incidence and the angle of refraction with curved arrows at both boundaries. Results Describe any pattern you found in the data between the angle of incidence and the angle of reflection. They are equal. Describe what happens to a light ray as it travels (a) from air into glass and (b) from glass into air. Bends toward the normal. Bends away from the normal. Assuming that light travels faster through air than it does through glass, make a generalized statement about what happens to a light ray with respect to the normal as it moves from a faster speed in one material to a slower speed in another. A light ray bends toward the normal when it travels from one medium to another medium with greater optical density. The light ray bends away from the normal when it traveled from a medium with greater optical density. What rules or generalizations do your findings suggest about reflection? How much more data would be required to make this a valid generalization? The angle of incidence equals the angle of reflection. What rules or generalizations do your findings suggest about refraction? How much more data would be required to make this a valid generalization? The angle of incidence equals the angle of reflection. Was the purpose of this lab accomplished? Why or why not? (Your answer to this question should be reasonable and make sense, showing thoughtful analysis and careful, thorough thinking.) (Students should know the difference between reflection and refraction, should be able to identify the normal, the angle of incidence, angle of reflection, and the angle of refraction.) Data Table 13.1 Reflection of Light Trial Angle of Incidence Angle of Reflection 1 2 3 Name____________________________________________________Section________________Date___________ Experiment 14: Physical and Chemical Change Invitation to Inquiry In many communities the recycling of aluminum, paper, and plastics is started by first segregating items made of these materials from the rest of the trash. A major problem in recycling plastics is the many different types of plastics that exist, all with different chemical and physical properties. Some of these materials are more desirable for recycling than others, so they must be sorted. One way of sorting plastics is to read the code that might be stamped on the bottom. Here are some letter and number codes from some common plastic items. The number usually appears inside the recycling arrow logo: 2 (HDPE) milk jugs, bleach, and detergent bottles; 1 (PETE) soft-drink bottles; 5 (PP) ketchup bottles, yogurt cups; 6 (PS) transparent plastic drinking cups, CD boxes, and; 4 (LDPE) plastic squeeze bottles. Can you find a way to separate a mixture of pieces of plastic from each of these 5 groups by taking advantage of the differences in chemical or physical properties? Consider cutting pieces of plastic from each of the group, and finding important properties that could be used in a separation scheme. Background Matter undergoes many changes and most of the common, everyday changes are either physical or chemical. A physical change is one that does not alter the identifying properties of a substance. It can be a change in form, state, or energy level, but no permanent change occurs in the properties of the substance. A piece of paper torn into two parts, for example, still has the properties of the original paper and no new substance has been formed. Thus, tearing a piece of paper into two parts is a physical change. Evaporation, condensation, melting, freezing, and dissolving often produce physical changes. A chemical change is one that produces new substances with new properties. A piece of paper that burns produces gases and ash that have different properties than the original paper, so this is an example of a chemical change. Heat, light, and electricity often produce chemical changes. Chemical changes occur as (1) atoms combine to form new compounds, (2) compounds break down into simpler substances, and (3) compounds react with other compounds or elements to form new substances. In this experiment you will determine if certain changes in matter are physical or chemical changes. Procedure Dissolve about 2 g of sodium chloride (ordinary table salt) in 50 mL of water in a clean graduatedcylinder. Observe (a) if the water level changes when the salt is added, and (b) the taste of the solution. CAUTION: Taste chemicals only when directed to do so. Record your observations in Data Table 14.1 on page 118. Continue dissolving 2 g samples of sodium chloride until 10 g are dissolved. Record your observations about the water level and the taste after each of the five additions. Place about 5 mL of the solution in an evaporating dish and heat slowly until a dry solid remains. When the solid has cooled, carefully taste the solid and record your observation in the data table. Carefully examine a 5-cm length of nichrome wire, noting the color, luster, flexibility, and other properties you can observe. Hold the wire with tongs and heat it strongly in the flame of a burner until it glows brightly. After the wire cools, again examine the wire, comparing the properties before and after heating. Record your observations. Examine a 5-cm length of magnesium ribbon, noting the color, luster, flexibility, and other properties you can observe. Hold the ribbon with tongs and ignite the end in the flame of a burner. CAUTION: Do not look directly at the magnesium while it is burning. Compare the properties of the ash with the properties of the original magnesium ribbon, recording your observations in the data table. Pour about 5 mL of silver nitrate solution in a small test tube. Add several drops of dilute hydrochloric acid, gently swirling the mixture with the addition of each drop. Record your observations in the data table. Filter the solid that forms (the precipitate), using small amounts of water as necessary to flush all the precipitate from the test tube. (Figure 14.1 shows how to Step 1: Fold along diameter. Step 2: Fold over a second time. Step 3: Open one fold to make cone. Step 4: Place cone in funnel. Figure 14.1 fold a filter paper and Figure 14.2 shows how to set up apparatus to filter a liquid.) Discard the filtered liquid (the filtrate) and place the filtered precipitate in direct sunlight. Record your observations after the precipitate has been in the sunlight for two or three minutes. Figure 14.2 Pour about 125 mL of copper(II) chloride solution in a beaker. Observe if the beaker feels cool or warm to touch. Obtain a piece of aluminum foil approximately 2 cm by 15 cm and form it into a loose coil. Drop the aluminum foil coil into the copper(II) chloride solution. Observe the results, if any, and cautiously touch the beaker after about five minutes. Record your observations in the data table. Results What kind of change occurs when sodium chloride dissolves? Give an explanation for your answer. The salt crystals become smaller and smaller, then disappear. The sodium and chloride ions are being pulled into and incorporated into the aqueous solution. What changes occurred with the water level, if any, as more and more sodium chloride was added to the solution? Provide an explanation for your observation. There appears only to be a slight increase in volume, which may or may not be measurable. The ions fit between the water molecules. Compare the evidence of a new substance being formed before and after heating (a) a nichrome wire, and (b) a magnesium ribbon. Nichrome - No new substance. An oxide film might be formed , but generally will not change the properties of the wire; depends on the temperature of the fire. Magnesium - A new substance formed, a white powder. No longer a metallic ribbon. What evidence indicates if new substances with new properties were or were not formed when hydrochloric acid was added to a silver nitrate solution? A white precipitate was formed. This is a new substance, silver chloride. Did the precipitate exposed to sunlight undergo either a chemical or physical change? Give an explanation for your answer. The white precipitate burns dark brown or black. This is the photo conversion of silver chloride to metallic silver, a chemical change. What changes did you observe when the aluminum was added to the copper(II) chloride solution? (Hint: There were more than five changes.) The aluminum goes into solution. A reddish-brown copper precipitate. The solution changes color to a milky-green solution. The solution becomes hot. Which of the changes in this experiment were physical changes? Give reasons for your conclusion. The increase of temperature is a physical change. It results from increased molecular kinetic energy and the molecules are not otherwise changed. Which of the changes in this experiment were chemical changes? Give reasons for your conclusion. All changes except the temperature change were chemical changes. New substances were formed. Was the purpose of this lab accomplished? Why or why not? (Your answer to this question should be reasonable and make sense, showing thoughtful analysis and careful, thorough thinking.) (The student should know the difference between physical change and chemical change. The student should be able to give examples of each type of change.) Name____________________________________________________Section________________Date___________ Experiment 15: Conductivity of Solutions Invitation to Inquiry Use some small gauge wire, a flashlight battery, and a flashlight bulb and holder to set up a portable conduction tester. Test a variety of materials to find which are conductors and which are nonconductors. Generalize what categories of materials are conductors and what categories seem to be nonconductors. Background Ionic compounds that dissolve in water do so by ions being pulled from the crystal lattice to form a solution of free ions. When free ions (charged particles) are present in a solution, the solution is a good conductor of electricity. Many covalent compounds that dissolve in water form molecular solutions of noncharged particles. Such a solution will not conduct an electric current. Some covalent compounds are pulled apart into ions; that is, the compound is ionized in a water solution. Strong ionization results in a solution that is a good conductor of electricity. Partial ionization results in a solution that is a poor conductor of electricity. The amount of current that flows through such a solution is roughly related to the amount of ionization. In this experiment you will use a conductivity apparatus (see Figure 15.1) to compare the conductivity of different solutions. Figure 15.1 CAUTION: The solutions in this experiment are dilute and relatively safe provided that you exercise care in their use. Acids and bases, however, can irritate the skin and damage clothing and other materials. If you spill an acid or base on your skin or clothing, wash it off immediately with plenty of water and inform your instructor. Procedure Set up the apparatus as shown in Figure 15.1. With the lamp unplugged, attach a patch cord with alligator clips to the two electrodes. Plug in the lamp to make sure the apparatus works, then unplug the lamp. Remove the patch cord and set it aside. Test each substance listed in Data Table 15.1 on page 123, using the conductivity apparatus with the following procedure: For solids: With the lamp unplugged, lower the electrodes until both are touching the solid. Plug in the lamp and record if it glows brightly, dimly, faintly, or not at all. Unplug the lamp and remove the electrodes from the solid. Wipe the electrodes with a clean cloth before proceeding. For solutions: With the lamp unplugged, lower the electrodes to the same depth in each solution. Plug in the lamp and record if the bulb glows brightly, dimly, faintly, or not at all. Unplug the lamp and remove the electrodes from the solution. Wash the electrodes by immersing them in distilled water before proceeding to the next solution test. Your instructor might specify other solids or solutions to be tested, which should be added to the list in Data Table 15.1. Figure 15.2 Results Compare your observations of the conductivity of dry sodium chloride and a solution of sodium chloride. What conclusions can you make about these observations? Dry sodium chloride does not conduct an electric current. A water solution of sodiumchloride does conduct an electric current. Dry sodium chloride is an ionic compound that has its ions tied up in the crystal lattice.When dissolved in water the ions are free to move about, so the solution is a good conductor of electricity. Compare your observations of the conductivity of distilled water and tap water. Offer an explanation for this comparison. Distilled water does not conduct an electric current at all. Tap water (in mostlocations) does conduct an electric current. It must be that the tap water contains some free ions from dissolved salts that are ableto carry the current. Compare the conductivity of hydrochloric acid and vinegar, which is also an acid solution. If the amount of conductivity is an indication of the degree of ionization of an acid solution, what would you predict about the degree of ionization of these two solutions? Both conduct an electric current, but hydrochloric acid appears to be a betterconductor. The prediction should be that hydrochloric acid is more highly ionized. What solutions were good conductors of electricity? What solutions were nonconductors of electricity? Describe any patterns that you can find in this summary. Good conductors were water solutions of sodium chloride, hydrochloric acid, and sodiumhydroxide. Non-conductors were dry sodium chloride, ethyl alcohol, glycerin, and sugar solution. In general, what kind of compounds conduct electricity when they are in solution? Explain. In general, water solutions of ionic compounds and water solutions of certain covalent compounds conduct electricity. These solutions contain free ions, meaning charged particles, to conduct a current. In general, what kind of compounds do not conduct electricity when they are in solution? Is this always true? Explain. In general, covalent compounds do not conduct electricity in a water solution. No, this is not always true. Hydrogen chloride, for example, is a covalent compound thatbecomes ionized in a water solution. Was the purpose of this lab accomplished? Why or why not? (Your answer to this question should be reasonable and make sense, showing thoughtful analysis and careful, thorough thinking.) (Students should know the difference between ionic and covalent compounds and how they will conduct an electric current.) Name____________________________________________________Section________________Date___________ Experiment 16: Metal Replacement Reactions Invitation to Inquiry Make a plan for using the fewest number of steps possible in an experimental study to place six metals (Mg, Zn, Fe, Sn, Cu, and Al) in an activity series from highest to lowest. Describe the procedure you would use for determining the position of a Ni salt in this series. Background Based on what happens to the reactants and products, there are basically four types of chemical reactions: (1) combination, (2) decomposition, (3) replacement, and (4) ion exchange reactions. This experiment is concerned with metal replacement reactions. A metal replacement reaction occurs when a more active metal is added to a solution of the salt of a less active metal. In generalized form the reaction is XY + A → AY + X, where A represents the more active metal and XY represents the salt of a less active metal. A metal replacement reaction occurs because some metals have a stronger electron holding ability than other metals. Metals that have the least ability to hold on to their electrons are the most chemically active. An activity series ranks the most chemically active metals at the top and the least chemically active at the bottom. This means that the activity series is upside down with respect to electron holding ability. Metals at the top of the activity series have the least ability to hold on to their electrons and those at the bottom have the greatest ability to hold on to their electrons. A metal replacement reaction occurs when a more active metal (with less electron holding ability) is added to a solution containing the ions of a less active metal (with greater electron holding ability). The more active metal loses electrons to the less active metal, so they trade places; that is, the active metal loses electrons to form metallic ions in solution, and the less active metal gains electrons and comes out of solution as a solid metal. In this experiment you will study the chemical reaction of three metals and rank them according to their electron holding ability. CAUTION: Soluble lead compounds are poisonous. Silver nitrate solutions will stain the skin. Avoid contact with both of these metal salt solutions. Wash thoroughly with soap and water if contact is suspected. Wash your hands after completing the experiment. Procedure Obtain three clean, dry test tubes and a test tube rack. Add 10 mL of silver nitrate solution to each test tube. Obtain three metal strips each of copper, zinc, and lead. Sandpaper all nine strips until they are clean and bright. Make a right-angle bend near the end of each metal strip. A 10-cm length of thin copper wire with one end formed into a hook can be used to fish the bend of a metal strip from a test tube. Use the copper wire to remove the metal strips from their solution for inspection as needed. Place a strip of each metal into a test tube with the silver nitrate solution. Observe the metals for 10 minutes or so, looking for evidence of a replacement reaction. Carefully touch each test tube at the outside bottom to observe any temperature changes. If a black deposit appears on a metal strip, observe for an additional 5 minutes or longer. In Data Table 16.1 on page 128, record your observations of any changes in the solutions and any changes on the metal strips. Repeat procedure steps 1 through 3 with fresh pieces of metal and 10 mL of copper nitrate solution. Observe the metals for 10 minutes or more, again looking for evidence of a replacement reaction. Repeat procedure step 3 again, this time with fresh pieces of metal and 10 mL of lead nitrate solution. Again observe the metals for 10 minutes or longer and record your observations in the data table. Results According to Data Table 16.1, which metals had the (a) greatest number of reactions; (b) middle number of reactions; (c) least number of reactions? Zinc Lead Copper Are your findings in question 1 consistent with the activity series of metals? Explain. Yes, zinc is the most active of the three and copper is the least active. Were any changes observed in the color of the solutions? Offer an explanation for this observation. Yes. Different ions going into or coming out of solution would change the color of the solution. Complete the following equations, writing “no reaction” if none occurred. Cu + AgNO3 → 2Ag + Cu(NO3)2 Zn + AgNO3 → 2Ag + Zn(NO3)2 Pb + AgNO3 → 2Ag + Pb(NO3)2 Cu + Cu(NO3)2 → No Reaction Zn + Cu(NO3)2 → Cu + Zn(NO3)2 Pb + Cu(NO3)2 → Cu + Zn(NO3)2 Cu + Pb(NO3)2 → No Reaction Zn + Pb(NO3)2 → Pb + Zn(NO3)2 Pb + Pb(NO3)2 → No Reaction Write a general rule describing when metal replacement reactions occur and when they do not occur. The reactions occur when a more active metal is placed in a solution of less active metal ions. The less active metal has a greater ability to "hold onto" electrons, taking some from the more active metal. This causes the more active metal to go into solution as ions and the less active metal ions to come out of solution as a metal. Was the purpose of this lab accomplished? Why or why not? (Your answer to this question should be reasonable and make sense, showing thoughtful analysis and careful, thorough thinking.) (The student should be able to describe a metal replacement reaction. The student should also be able to predict whether a simple reaction will be a metal replacement reaction or not.) Name____________________________________________________Section________________Date___________ Experiment 17: Identifying Salts Invitation to Inquiry Foods naturally contain enzymes, biochemical compounds that originate in plants and animals. Cooking changes enzymes, and the purpose of blanching, or steaming food shortly before freezing is intended to destroy enzymes. One of the enzymes in foods is a catalyst that will speed the decomposition of hydrogen. You can design a home experiment to use hydrogen peroxide (3% solution) to test fresh, blanched, and cooked crushed food (or juices from the foods) for the catalyst enzymes. If you can find a way to quantify the measurements, perhaps you can come up with specific recommendations about how hot the food should be heated, and for how long. Temperature is one of the more important factors that influence the rate of a chemical reaction. You can use a “light stick” or “light tube” to study how temperature can influence a chemical reaction. Light sticks and tubes are devices that glow in the dark and have become very popular on July 4th, Halloween, and other times why people might be outside after sunset. They work from a chemical reaction that is similar to the chemical reaction that produces light in a firefly. Design a home experiment that uses light sticks to find out the effect of temperature on the brightness of light and how long the device will continue providing light. Perhaps you will be able to show by experimental evidence that use at a particular temperature produces the most light for the longest period of time. Background A salt is any ionic compound except those with hydroxide (OH- ) or oxide (O-2) ions. Table salt, which is sodium chloride (NaCl), is but one example from the large group of ionic compounds known as salts. Simple salts consist of a metallic ion (such as Na+) in an ionic crystalline structure with nonmetallic ions (such as Cl- ). When a salt is dissolved in water the ion crystal structure is separated into a solution of metallic and nonmetallic ions. The name of a salt provides a clue about the ions present in a solution. Lithium chloride (LiCl) becomes Li+ and Cl- ions when dissolved in water. Likewise, calcium iodide (CaI2) becomes a solution of Ca+2 and I- ions, and iron(II) carbonate [Fe2(CO3)3] becomes a solution of Fe+3 and CO3-2 ions. As the water of a salt solution is evaporated, the salt ions again form an ionic crystal structure. Knowing the ions present in a given salt solution will therefore identify the salt in the solution. In this experiment, you will use a flame test to identify metallic ions present in a salt solution. Each metal ion gives a characteristic color to a burner flame. The nonmetallic ions will be identified by chemical tests. CAUTION: Acids will damage human tissue and other materials. Silver nitrate will stain the skin. Some of the chemicals are poisonous when swallowed. Be sure to wash thoroughly if any chemicals are spilled and inform your instructor. Wash your hands thoroughly after the experiment. Procedure Part A: Flame Tests Obtain seven test tubes and a test tube rack. Wash the test tubes and rinse them thoroughly with distilled water. Cleanliness is very important throughout this experiment. Pour about 25 mL of dilute hydrochloric acid into one test tube and label it in a rack. Pour about 5 mL of each of the six nitrate solutions into test tubes and label each in the rack. Set up a burner and adjust it for an inner blue cone and no yellow color. Clean a platinum or nichrome wire by dipping it into dilute hydrochloric acid, then holding it in the pale blue part of the flame. Repeat the cleaning procedure until the wire gives no color to the flame. The flame test you are about to do is a very sensitive test and even a trace of contamination of sodium from your fingers will give the characteristic color of the sodium ion (a yellow coloration). Dip the clean wire into one of the solutions and place just the tip of the wire into the light-blue burner flame. Record your observations in Data Table 17.1 on page 132. Clean the wire, then repeat the flame test with each of the solutions. Record the results obtained for each in Data Table 17.1. Use two thicknesses of cobalt glass to view the potassium and sodium flames. The sodium ion gives a color that masks the color given by the potassium ion. The cobalt glass filters out the color produced by any sodium ions. Part B: Chemical Tests Obtain four clean test tubes and a test tube rack. Pour about 5 mL of each solution of sodium chloride, potassium sulfate, calcium carbonate, and one of the nitrate solutions from the flame test into separate test tubes. Label each in the test tube rack. Chloride ion test: To 5 mL of sodium chloride solution, add 5 drops of silver nitrate solution. Mix the solutions well, then describe the results in Data Table 17.2 on page 133. Sulfate ion test: To 5 mL of potassium sulfate solution, add 10 drops of barium chloride solution. Mix well, then describe the results in Data Table 17.2. Carbonate test: To 5 mL of calcium carbonate solution, add several drops of dilute hydrochloric acid. Describe the results in Data Table 17.2. Nitrate ion test: To 5 mL of a nitrate solution, add 10 mL of iron(II) sulfate solution. Use a long, thin pipette to slowly add 2 mL of concentrated sulfuric acid so it runs down the inside of the tube without mixing. Observe the interface between the acid and the solution, recording the results in Data Table 17.2. Part C: Unknown Solutions Wash test tubes of all salt solutions and rinse thoroughly with distilled water. Your instructor will supply unknown solutions, each containing a single metallic and a single nonmetallic ion. Using the procedures and findings from part A and part B, identify the ions in the unknown solutions. Record your findings in Data Table 17.3 on page 134, describing your test results as very positive, positive, trace, negative, or unsure. Data Table 17.2 Chemical Tests for Nonmetallic Ions Nonmetallic Ion Results Chloride AgNO3 + HCl → H+ + NO3- + AgCl↓ Sulfate K2SO4 + BaCl2 → K+ + Cl- + BaSO4↓ Carbonate H2CO3 + HCl → H+ + Cl- + H2O + CO2↑ Nitrate NO3- + FeSO4 → Fe(NO)2+ NO3- + 3Fe2++ 4H+ → NO + 3Fe3+ + 2 H2O Fe2++ NO → Fe(NO)2+ (Brown Ring) Data Table 17.3 Unknown Solutions Unknown Metallic Ion Tests Nonmetallic Ion Tests Compound 1 Sodium: Lithium: Strontium: Calcium: Barium: Potassium: Chloride: Sulfate: Carbonate: Nitrate: Name: _________________ Formula: _________________ 2 Sodium: Lithium: Strontium: Calcium: Barium: Potassium: Chloride: Sulfate: Carbonate: Nitrate: Name: _________________ Formula: _________________ 3 Sodium: Lithium: Strontium: Calcium: Barium: Potassium: Chloride: Sulfate: Carbonate: Nitrate: Name: _________________ Formula: _________________ Name____________________________________________________Section________________Date___________ Experiment 18: Natural Water Invitation to Inquiry Washing soda, borax, or trisodium phosphate are often added to laundry products to soften the wash water. Add a small amount of each to 5 mL of tap water, then measure the hardness with the soap-drop method described in this experiment. Is one of the softening chemicals more effective than the others? Is there a practical way to obtain pure water from a hard water source? Design, use, and evaluate an apparatus for purifying water—by boiling, freezing, filtering, precipitation of dissolved minerals—or by any means you can imagine. Evaluate your technique in terms of usefulness and effectiveness. Background Water is sometimes referred to as the “universal solvent” because so many gases, liquids, and solids readily dissolve in it. In addition, the solubility of carbon dioxide in water produces an acid (H2CO3) that contributes to the dissolution of normally insoluble carbonate, phosphate, and sulfite minerals. Thus, natural water can contain a significant amount of dissolved mineral matter as well as suspended solids. This occurs naturally from weathering and erosion of the earth’s surface by rainwater and by erosion of stream beds. Natural water that contains dissolved mineral salts in the form of calcium ions or magnesium ions is called hard water. It is called hard water because the metal ions react with soap, making it hard to make soap lather. Instead of producing suds, soap in hard water produces an insoluble, sticky precipitate. This requires a greater amount of soap to produce suds since all the metal ions must be precipitated out of solution before the soap will lather. The precipitate also produces a bathtub ring, collects on clothes, and results in other undesirable effects. When boiled, hard water can produce a solid deposit that restricts water flow in water heaters and pipes. Water hardness is usually measured in parts per million (ppm), with hard water identified as any concentration of calcium or magnesium ions greater than about 75 ppm. In this experiment you will learn a method of analyzing water hardness as well as analyzing the type of hardness present in your water supply. Procedure Part A: Suspended Solids Natural water usually contains suspended solids that are removed by (a) natural settling, (b) addition of a gelatinous precipitate, and (c) filtering through sand. Compare these methods by obtaining four beakers, each with 100 mL of muddy water. Allow beaker #1 of muddy water to stand undisturbed as you complete the rest of this section. This will provide a comparison of natural settling to the other methods of removing suspended solids. Make a sand filter by placing a loose plug of glass wool in a funnel, then covering it with several cm of clean sand. Pour beaker #2 of muddy water through the filter and collect the filtrate in a clean beaker. In beakers #3 and #4, produce the gelatinous precipitate of aluminum hydroxide by adding to each 10 mL of alum solution, then 4 drops of ammonium hydroxide while stirring. Allow both suspensions to settle. Prepare a second sand filter as in procedure step 3, then filter one of the settled beakers of aluminum hydroxide suspension from procedure step 4. Collect this filtrate in a clean beaker and save it for part C of this experiment. Compare the appearance of the muddy water from each procedure step in Data Table 18.1 on page 140. Part B: A Test for Hard Water The calcium and magnesium ions of hard water react with soap, forming an insoluble precipitate. Soap will form suds only after the ions have been removed, so the amount of soap needed to produce suds is an indication of the amount of water hardness. Measure 5 mL of the calcium chloride solution and pour it into a test tube. This solution has been prepared to have a water hardness of 100 ppm. Add one drop of soap solution to the calcium chloride solution and insert a clean stopper. Shake the test tube vigorously, then check the surface for soap suds. If none are visible, or if visible suds do not persist for at least one minute, add a second drop and shake the test tube vigorously again. Continue adding one drop at a time and shaking, keeping track of the number of drops added until a layer of suds persists for at least a minute. Record the number of drops required for permanent suds in Data Table 18.2 on page 141. Repeat procedure step 3 with a clean test tube and 5 mL of distilled water. A certain number of soap drops will be required to make a permanent suds layer because of the mechanism by which suds are produced. In Data Table 18.2, record the number of soap drops required for this mechanism to occur in distilled water. Complete the calculations required in Data Table 18.2 to find the hardness equivalent to one drop of the soap solution. Part C: Water Hardness Repeat procedure step B-3 with 5 mL of the flocculated and filtered water saved from part A of this experiment. Record the number of drops of soap required in Data Table 18.3 on page 142, and calculate the hardness of the water. Repeat the soap-drop test with 5 mL of ordinary cold tap or well water. Record the number of soap drops required in Data Table 18.3 and calculate the water hardness. Water hardness caused by calcium or magnesium bicarbonate can be removed by boiling, so it is called temporary hardness. The hardness remaining after boiling is permanent hardness. Temporary hardness is therefore the difference between total hardness and permanent hardness. Pour 25 mL of cold tap water or well water into a small beaker. Place the beaker on a wire screen on a ring stand, with a clean watch glass on top of the beaker to prevent water loss by splattering. Gently heat the water to boiling, continuing a slow boil for 5 or 6 minutes. Cool the water to room temperature, then pour the water into a graduated cylinder. Add distilled water as necessary to bring the volume back to 25 mL. Pour 5 mL into a test tube. Use the soapdrop method to determine the permanent hardness. Calculate the temporary hardness. Record all data in Data Table 18.3. Results Compare the appearance of muddy water after natural settling and the other means of removingsuspended solids. Which method produced the best apparent purity? Is there any impurity that the method does not remove? After natural settling, the water probably has a somewhat tint of murky or milkiness.Filtering might remove some of the tint, depending on local conditions, but most will be removed by the use of the gelatinous precipitate. The apparently best purity was produced by the settling of the aluminum hydroxidesuspension, the filtration. Yes, dissolved ions such as salts are not removed. Explain how a soap solution was standardized to measure water hardness. A standardized 100 ppm hard sample was prepared. The number of soap drops needed to form suds was determined. A 0 ppm standard was prepared. The number of soap drops needed to form suds was determined for this sample. A hardness equivalent was determined after subtracting off the sudsing mechanism. What is hard water? Hard water contains calcium and/or magnesium ions that react with soap, forming a precipitate. Explain how temporary and permanent hardness can be measured. Temporary hardness can be removed by boiling, but permanent hardness cannot. Therefore you could determine the total hardness of a sample of water using the soap drop method. The water could then be boiled to remove the temporary hardness, then the soap drop method could be used again. The temporary hardness is the difference between the two measurements. What is the white solid that forms on water outlets in hard water areas? How could you test the substance to confirm your answer? It is probably calcium or magnesium carbonate that formed from temporary hardness. A few drops of vinegar or any other acid solution will cause the carbonate to fizz and dissolve. Was the purpose of this lab accomplished? Why or why not? (Your answer to this question should show thoughtful analysis and careful, thorough thinking.) (The student should be able to define hard water, both permanent and temporary. Students should also be able to identify the ions present in the water supply.) Data Table 18.2 Hardness Equivalent of Soap Solution 1. Drops of soap solution required for calcium chloride solution (100 ppm) 10 ___________ drops 2. Drops of soap solution for distilled water 2 ___________ drops 3. Subtract sudsing mechanism (row 1 – row 2) 8 ___________ drops 4. Hardness equivalent (100 ppm ÷ row 3) 12.5 _________ ppm/drop NOTE: Answers will vary with the dropper size and soap solution. Solution Manual Experiment for Integrated Science Bill W. Tillery, Eldon D. Enger , Frederick C. Ross 9780073512259
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