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CHE 1401 School of Science & Engineering LABORATORY MANUAL FOR CHEMISTRY AND THE ENVIRONMENT Last Update: 7 July 2015 Last update: June 2011 1 CHE 1400 Name: ________________________ Section: ________________________ LABORATORY MANUAL FOR CHEMISTRY AND THE ENVIRONMENT Last Update: 7 July 2015 Last update: June 2011 1 CHE 1400 TABLE OF CONTENTS Laboratory introduction Common laboratory equipment iii viii Experiment 1: Physical Properties of a Compound 1 Experiment 2: Atoms and Light 7 Experiment 3: Solids in Cigarette Smoke 11 Experiment 4: Chemical Reactions 17 Experiment 5: Solubility 23 Experiment 6: Acids, Bases, Buffers and pH 26 Experiment 7: Acid Neutralization by Antacid 38 Experiment 8: Oxidation and Reduction 42 Experiment 9: Clarification of Water 46 Experiment 10: Saponification 51 Experiment 11: Water Analysis (field trip) 56 Appendix I Appendix II 62 64 Lab Manual ii CHE 1400 Laboratory introduction Laboratory safety, chemicals, equipment and techniques Laboratory safety Common sense precaution and a proper understanding of the techniques and chemicals being used make a chemical laboratory no more dangerous than parts of the home, such as the kitchen or garage. 1. Wear safety glasses. Safety glasses must be worn at all times when working in the laboratory. It’s particularly important to avoid the use of contact lenses since these lenses can result in more serious injuries if a toxic chemical is splashed into the eye. 2. Wear suitable clothing. Clothing acts as a protection against spilled chemicals or a burning liquid. Clothing which exposes large areas of bare skin can be a major hazard; open toed shoes or sandals increase the possibility of injury to the foot. A lab coat serves to clothes and skin, and should be worn in the laboratory at all times. Hair extending below the shoulder blades should also be secured to avoid damage to the hair itself, as well as to reduce the possibility of spilling a chemical. 3. Watch out for broken and hot glassware. Cuts and burns are the two most common types of laboratory injury. Do not use any glassware that is broken, chipped or badly cracked. If you have any questions on using a piece of glassware, check with the instructor. Hot glass looks no different than a piece of cold glass. Most burns result from trying to handle a piece of hot glass or other hot material. Be aware of any glassware or other items that have been heated on a hot plate or in a Bunsen burner flame. 4. Locate and be familiar with safety equipment. Know where to find and how to use safety and first aid equipment, such as fire extinguishers, eye wash stations, safety showers, etc. Remember that even if you know the location of safety equipment, it is of very little use unless you know how and when to use it. 5. Be careful with the chemicals. Consider all chemicals as potentially dangerous until you know otherwise. Avoid rubbing your eyes if there is any possibility that your hands may be contaminated with a chemical. Never eat or drink anything while working in the laboratory. 6. Read all labels. Read the label on any reagent bottle before you use it. Be sure to read the complete name of the chemical since many chemical names sound and look very similar. Also be sure to read the concentration of the reagent since a 5.0 M solution of sulfuric acid is very different than of a 0.005 M solution of the same acid. 7. Dispose of all chemicals properly. Discard any chemical in the proper container. Mixing of certain chemicals can start a fire or produce other hazards. Always dispose of unused chemicals; never return a chemical to the original container. 8. Wash any chemical off of your skin. If corrosive chemicals get on your skin or in your eyes, flush the affected area immediately with a large volume of clean water. Be sure you notify the laboratory instructor as soon as possible. 9. Smell any chemical cautiously. Never taste a chemical. If you need to smell a chemical do not smell the source directly. Instead use a cupped hand to bring a Lab Manual i CHE 1400 small amount of the vapor to your nose. 10. Use the fume hood when appropriate. Any chemical, which is irritant, dangerous or which has an unpleasant odor, should always be used in a properly operating fume hood. 11. Perform only authorized experiments. Perform only the experiments noted in the laboratory manual or those given to you by your laboratory instructor. 12. Add concentrated acids and bases to water. Always use the dilute acids and bases provided to you. If you are instructed to dilute a strong acid or base, pour the acid or base into the water. Remember that diluting many acids or bases can generate large amount of heat. 13. Clean up your workstation. Be sure to empty all reagents into the appropriate waste container, wash out any used glassware and wipe off the bench top. Be sure the gas, water and all electrical equipment are turned off. 14. In case of accident. Notify your instructor immediately if either you or another student had any type of accident. 15. Proper use of equipment. Misuse can lead to injury. Please do not hesitate to ask questions if you are unsure about the use of equipment. Chemicals Laboratory reagents shared by several students are stored in the fume hoods. Special reagents and unknowns will be issued by the laboratory instructor. 1. Aqueous solutions (water based) can, in most cases, be poured down the sink followed by a large volume of water. 2. Concentrated acids, concentrated bases, and organic solvents (halogenated and non- halogenated) should not be poured down the sink. Acids and bases should be first neutralized, then disposed of in the sink with a large volume of water. Water soluble organic acids should be diluted with water, then poured down the sink. Water-insoluble organic acids should be disposed in the containers provided. 3. Dispose of solids in the labeled waste containers: never dispose solids in the sink. 4. Do not place stoppers from the bottles on the bench top. Stoppers should be held in the hand until they can be returned to the original container. 5. Never return unused reagents to a reagent container. Dispose of the reagent in an appropriate manner. 6. Return reagent bottles to the work area shelf (reagents should never be taken out of the fume hoods). Do not take the reagent bottles to your work bench area where they will be unavailable for the other students. 7. Use only the amounts (approximately) of the reagent called for in the laboratory manual, or indicated by the laboratory instructor. Equipment Each student will have the use of a set equipment at an assigned work place. It should be remembered that each set of equipment is used by other students during other laboratory sessions. 1. Keep both the work area and the laboratory equipment clean and in good working condition. Do not leave dirty equipment laying around after the laboratory session 2. Do not borrow equipment from other desks. If you need extra equipment or are having problems with broken or malfunctioning equipment, obtain the additional equipment or replacement equipment from the stock room. Lab Manual ii CHE 1400 Laboratory techniques The laboratory principles that you learn in lecture and from the textbook are the result of numerous laboratory experiments. Years of observations, experimentation, interpretation and predication are necessary before a principle becomes well enough established to be included in an introductory chemistry course. Most laboratory work is made up of the same basic techniques. In this laboratory session you will study the basic rules, equipment and techniques used in a chemical laboratory. These include: - handling and transferring solid and liquid chemicals - using graduated cylinders, pipettes, burettes and volumetric flasks - using a laboratory balance - evaporating liquids - titrating a solution - separating a liquid and a solid Laboratory equipment From the list below and the pictures given on the following page, identify the common laboratory equipment at each work station and give a brief indication of what each piece of equipment is used for. Some of the equipment shown in the picture may not be at the work station. Number Item Use Common laboratory equipment 1-2 graduated cylinder ________________________________________________ 3 beakers ________________________________________________ 4 stirring rods ________________________________________________ 5 wash bottle ________________________________________________ 6 funnel ________________________________________________ 7-8 Erlenmeyer flasks ________________________________________________ 9-11 test tubes ________________________________________________ 12-13 test tubes racks ________________________________________________ 14 glass plate ________________________________________________ 15 wire gauze ________________________________________________ 16 crucible tongs ________________________________________________ Lab Manual iii CHE 1400 Number Item Use 17 spatulas ________________________________________________ 18 litmus paper ________________________________________________ 19 watch glasses ________________________________________________ 20 evaporating dish ________________________________________________ 21 dropping pipettes ________________________________________________ 22 test tube holder ________________________________________________ 23-24 test tube brushes ________________________________________________ Special laboratory equipment 1 reagent bottle ________________________________________________ 2 condenser ________________________________________________ 3 500-mL Erlenmeyer ________________________________________________ 4 1000-mL beaker ________________________________________________ 5 petri dish ________________________________________________ 6 Buchner funnel ________________________________________________ 7 Buchner flask ________________________________________________ 8 volumetric flask ________________________________________________ 9 500-mL Florence flask ________________________________________________ o 10 110 C thermometer ________________________________________________ 11 100-mL graduated cylinder ________________________________________________ 12 50-mL buret ________________________________________________ 13 glass tubing ________________________________________________ 14 U-tube ________________________________________________ 15 porous cup ________________________________________________ 16 crucible & cover ________________________________________________ 17 mortar & pestle ________________________________________________ 18 glass bottle ________________________________________________ 19 pipettes ________________________________________________ 20 ring stand ________________________________________________ 21 buret clamp ________________________________________________ 22 double buret clamp ________________________________________________ Number 23 Item Bunsen burner Lab Manual Use ________________________________________________ iv CHE 1400 24 burette brush ________________________________________________ 25 clay triangle ________________________________________________ 26 rubber stoppers ________________________________________________ 27 wire loop ________________________________________________ 28 pneumatic trough ________________________________________________ 29 rubber pipettes bulb ________________________________________________ 30 iron support ring ________________________________________________ Lab Manual v CHE 1400 COMMON LABORATORY EQUIPMENT Lab Manual vi CHE 1400 Lab Manual vii CHE 1400 Lab Manual viii CHE 1400 EXPERIMENT 1 Physical properties of a compound OBJECTIVES 1. To identify a compound based on its physical properties 2. To learn how to properly assemble and use a simple boiling point apparatus Relates to chapter 1 of “Chemistry for changing times, 13th Ed.”. INTRODUCTION Gold is a yellow solid, salt is a white crystal, lead is heavy, chlorine is a greenish-yellow gas and water is a clear colorless liquid. These are all physical properties of chemical substances. Physical properties can be used to identify a chemical substance. The common physical properties include color, odor, density, solubility, crystal structure, melting point and boiling point. Additional physical properties such as conductivity, malleability, etc. can also be determined. 1. Solubility The solubility of a substance is most accurately defined as the maximum mass (expressed in grams) of the test substance that dissolves in a known mass (usually 100 grams) of another substance at a given temperature. The test substance is referred to as the solute and the substance in which the test substance is being dissolved is the solvent. A given chemical has different solubilities in different solvents, depending on the similarities in molecular composition of the two substances. Usually, “like dissolves like“ and therefore highly polar compounds like sodium chloride (table salts) are soluble in high polar solvents like water, but not in low polar solvents like oils. Similarly, low polar substances like butter dissolve in low polar solvents (oils) but not in high polar ones (water). The degree of solubility can also be expressed in a rough way by the terms « soluble », « slightly soluble » or « insoluble ». A soluble solid dissolves quickly and easily in the solvent soluble liquid mixes quickly and easily with the solvent. An insoluble substance will not dissolve or mix with the solvent at all. A slightly soluble substance will dissolve or mix to a limited degree. 2. Density Density is the mass per unit volume. A substance with a high density has a large mass in a small volume. While we often say that lead is « heavy », what we really mean is that lead has a high density. Density is commonly given in terms of grams per milliliter (g/mL), although other units of mass and volume can be used. 3. Boiling point When a liquid is gradually heated, there is a point at which the temperature of the liquid no longer increases, but bubbles of vapor (gas) form spontaneously, and continue to do so until the entire volume of liquid has been converted to a gas. This constant temperature is called the boiling point of the liquid. The boiling is dependent on both the liquid being heated and the atmospheric pressure on the liquid. Boiling points are normally given at Lab Manual 1 CHE 1400 normal atmospheric pressure. EXPERIMENTAL PROCEDURE Note down the number of the experimental unknowns on your laboratory report sheet. 1. Solubility a. Solubility in water Into a 25-mm test tube add about 1 mL of water. Add about 1 mL of the liquid unknown. Shake the test tube several times and note if there is any change in the appearance of the mixture. If the two liquids mix (even if shaking is required) to give a single mixture without a meniscus, then that unknown is soluble in water. If the unknown liquid does not mix with water and you observe two distinct layers with a meniscus between them (even after shaking), then that unknown is insoluble in water. b. Solubility in ethanol Into a 25 mm test tube add about 1 mL of ethanol. Add about 1 mL of the liquid unknown. Shake the test tube several times and note if there is any change in the appearance of the mixture. If the two liquids mix (even if shaking is required) to give a single mixture without a meniscus, then that unknown is soluble in ethanol. The meniscus is the slightly curved surface at the top of a column of liquid. If the unknown liquid does not mix with ethanol and you observe two distinct layers with a meniscus between them (even after shaking), then that unknown is insoluble in ethanol. c. Solubility in cyclohexane Into a 25-mm test tube add about 1 mL of ethanol. Add about 1 mL of the liquid unknown. Shake the test tube several times and note if there is any change in the appearance of the mixture. If the two liquids mix (even if shaking is required) to give a single mixture without meniscus, then that unknown is soluble in cyclohexane. If the unknown liquid does not mix with ethanol and you observe two distinct layers with a meniscus between them (even after shaking), then that unknown is insoluble in cyclohexane Record your results for the solubility of the unknown in water, ethanol and cyclohexane on the laboratory report worksheet. 2. Density Place a 25-mm test tube into a 100 mL beaker and weigh the assembly to the nearest 0.1 g on a top loading balance (be sure that the balance is zeroed). Record the weight of the beaker/test tube assembly on the laboratory report worksheet. Pipette 1 mL of the unknown into the 25-mL test tube and reweigh. Use a rubber pipetting bulb to draw the liquid into the pipet. DO NOT pipet any liquid by mouth! Make sure that you pipet the correct volume of liquid. Lab Manual 2 CHE 1400 If you don’t know how to use a pipette, please check with the laboratory instructor. Record the weight of the beaker/test tube assembly containing the unknown on the laboratory report worksheet. Be sure that you use the same balance each time, otherwise you may get slightly different values due to differences in the balances. Calculate the mass of the unknown by subtracting the mass of the [beaker/test tube assembly] from that of the [beaker/test tube assembly plus unknown]. If you wish, you may tare the balance with the test tube/beaker assembly and then get the mass of the unknown directly. If you use this method, be sure to get the mass of the unknown immediately after you have tared the balance. Otherwise, another student may tare the balance to their glassware. Calculate the densities of the unknown by dividing the mass of the unknown by the volume of that unknown. Complete a second determination of the density and use the average of your two experimental results. 3. Boiling point 3.1. Assemble the boiling point apparatus Assemble the boiling point apparatus as shown in Figure 3.1. Place 1 to 2 mL of your unknown liquid into a 75-mm test tube. Next to the thermometer bulb secure a 10 cm capillary tube in an inverted position (open end down) using a rubber band. Place the thermometer bulb and capillary tube into the liquid. Place the apparatus into a water bath. Figure 3.1 Apparatus for determining the boiling point of a liquid. Lab Manual 3 CHE 1400 3.2. Measure the boiling point Make sure that the water bath contains 5-8 « boiling chips ». Slowly heat the water in the water bath. You should notice a bubble of gas escaping from the end of the tube every once in a while. As the temperature of the water increases, the rate at which the bubbles are formed should increase. When a rapid and continuous stream of bubbles escapes from the capillary tube, discontinue heating by either shutting off the hot plate, removing the test tube from the water bath, or removing the water bath and test tube from the hot plate. Caution! The water and the glassware may be quite hot. As soon as the bubbles from the capillary tube STOP, take a note of the temperature. Record the temperature of the water bath as the boiling point of the liquid unknown. To get as accurate a reading as possible, watch the capillary tube closely and record the temperature immediately. 3.3. Repeat the measurement Determine the boiling point of the unknown a second time. The capillary tube and all liquid must be removed. Reinsert the capillary tube before heating is resumed. Also, the water bath should be cooled by removing some of the hot water and adding some cool water along with some fresh boiling stones. 4. Identification of the unknown compound Based on your observation, experimental results, calculations and the information given in Table 1, determine the identity of the liquid unknown. If you cannot make a reasonable determination based on all of your data, make a determination on two of the three factors (solubility, density and boiling point) and repeat the experiment for the third factor. For example, if you observe that your unknown is soluble in all three solvents and has a boiling point of 54 °C (indicating it is acetone) but the density was calculated as 0.95 g/mL (instead of 0.79 g/mL), then carefully repeat the density experiment to see if you get the same value. Explain any problems you had in either the procedures or in the identification of the unknown. Lab Manual 4 CHE 1400 Experiment 2 Physical properties of a compound Name(s) Date Laboratory Instructor Unknown n°_ REPORT SHEET 1-Solubility Solvent Soluble Insoluble Water Ethanol Cyclohexane 2-Density Trial 1 Trial 2 Mass of test tube and beaker (g) ___________ ___________ Mass of test tube, beaker & unknown (g) ___________ ___________ Mass of unknown liquid (g) ___________ ___________ Volume of unknown liquid (mL) ___________ ___________ Density of unknown liquid (g/mL) ___________ ___________ Average density of unknown liquid ___________ g/mL 3-Boiling point Observed b.p. 1: ______ °C Observed b.p. 2: _____ °C Average b.p. : ________ °C 4.Identification of the unknown compound Unknown compound: ________________________________________ Explanations and comments: Lab Manual 5 CHE 1400 TABLE 1- Physical Properties of Some Common Laboratory Chemicals Compound Density Melting point Boiling point (g/cm3) (oC) (oC) Water Ethanol Acetone 0.79 -95 56 s s Acetamide 1.00 82.3 221 s s - - Acetanilide 1.22 114 304 - s s s Anthracene 1.28 216 - - - s s Benzamide 1.08 132 290 s s - s Benzoic acid 1.07 122 249 - s s s Benzoin 1.31 137 344 - s s - 2-Butanone 0.81 -86 80 s s s s Cyclohexane 0.79 6.5 81 - s s Cyclohexene 0.81 -104 83 - s s s Ethanol 0.79 -117 79 s s s Ethyl acetate 0.90 -84 77 s s s s Heptane 0.68 -91 98 - s s s n-Hexane 0.66 -95 69 - s - - Methanol 0.79 -94 65 s s s s Naphthalene 0.96 80.5 218 - s s s 1-Propanol 0.80 -127 97 s s s s 2-Propanol 0.79 -90 82 s s s s Lab Manual Solubility Acetone Cyclohexane s 6 CHE 1400 EXPERIMENT 2 ATOMS AND LIGHT Light is not just light OBJECTIVES 1. To observe the colors of light emitted from excited atoms and relate these wavelengths to the energy changes in the electrons within the atom. 2. To learn to use a diffraction grating to separate the emission spectrum into separate wavelengths of light. Relates to chapter 3 of “Chemistry for changing times, 13th Ed.”. BACKGROUND Atoms are made up of positive nuclei surrounded by negative electrons. These electrons have different amounts of energy as a rule, but the important fact is that electrons cannot possess just any amount of energy but only certain amounts of energy. How do we know that? This knowledge comes from a study of the electromagnetic spectrum in the x-ray, ultraviolet, and visible portions of the spectrum. In today’s experiment, a detailed observation of the light coming from hot atoms, will lead you to the same conclusion. When an atom is heated, electrons absorb energy in definite amounts and as they cool, they emit that extra energy which we see as a particular color of light. Electrons in different kind of atoms absorb and then emit different amounts of energy, which create different spectra. The spectra or combination of colors, observed from a discharge tube, can be used to identify what element is glowing. This was the way helium was discovered in the sun before it was known on earth. Without a diffraction grating the one most prominent color is seen. Atoms with many electrons, like tungsten, give off so many colors of light we observe a complete spectrum, i.e., white light. Through a diffraction grating, we would observe a rainbow of the complete spectrum from tungsten. A diffraction grating is like a prism in that it will separate the different wavelengths of light. It is interesting that, when electrons behave as waves as well as particles (just as light does), properties of waves lead to the same those electrons around nuclei can have only certain energies. Some instruments, such as the atomic absorption spectrometer make use of the fact that every element has a particular set of energies, which are absorbed when its electrons are excited. These particular energies, or wavelengths, of light can be used to detect the presence of a particular element. Other instruments identify elements by the wavelengths of light, which are re-emitted when the atoms are excited. When a solution containing the atoms of an element is heated in a flame, we can use these energy signatures to detect which elements are present in the solution. Lab Manual 7 CHE 1400 WASTE AND THE ENVIRONMENT Only concentrated hydrochloric acid poses a hazard. Concentrated acid can damage plumbing if not neutralized or diluted. Once used in the flame test, it should be flushed down the drain with lots of water. PROCEDURE NOTE: the room needs to be darkened for best results in this experiment 1. Hold a diffraction grating up to your eye and view the candle light through it. Record your observations. 2. Using the diffraction grating, view each of the following light sources when demonstrated by the instructor: a- candle b- helium light bulb c- neon light bulb d- argon light bulb e- krypton light bulb Record your observation(s) for each light source. 3. Without the grating, observe the colors as you place salts of the following metals into the flame of a Bunsen burner. Make sure to wash the spatula with concentrated hydrochloric acid between each salt. Record these observations. a- Lithium b- Sodium c- Barium d- Strontium e- Potassium f- Copper Lab Manual 8 CHE 1400 Experiment 2 Atoms and light Name(s) Date Laboratory Instructor Unknown n°_ REPORT SHEET I – Describe the view of each of the following through a diffraction grating: a- Candle ________________ b- Argon ________________ c- Krypton ________________ d- Neon ________________ e- Helium ________________ II – Describe the order of salts of these metals in a flame: a- Lithium ________________ b- Sodium ________________ c- Barium ________________ d- Strontium ________________ e- Potassium ________________ f- Copper Lab Manual ________________ 9 CHE 1400 III – Questions 1- Could the colors of the flame be used to identify different metal ions in a mixture? 2- Why is a yellow shirt yellow? 3- What can we learn from a study of light coming from different stars? Lab Manual 10 CHE 1400 EXPERIMENT 3 SOLIDS IN CIGARETTE SMOKE Smoke gets in your eyes……….. and that’s not all! OBJECTIVES 1. To use the process of mass difference to make quantitative measurements of the solids produced by a cigarette. 2. To determine whether the smoker or a person exposed to second-hand smoke captures the largest amount of smoke solid. 3. To make a quantitative determination of the segment of a cigarette which delivers the smallest mass of solid particles per mass of cigarette consumed. Relates to chapter 13 and 19 of “Chemistry for changing times, 13th Ed.”. BACKGROUND Cigarette smoke is generated by burning tobacco leaves. The smoke produced in this combustion reaction has at least two major components, which are harmful to the human body. Carbon monoxide (CO), one of the major components of cigarette smoke, is an odorless, invisible, but toxic gas. Carbon monoxide attaches to hemoglobin (making carboxyhemoglobin) and prevents hemoglobin (as oxy-hemoglobin) from carrying oxygen to the rest of the body. The affinity of hemoglobin for carbon monoxide is 200 times greater than for oxygen. A cigarette smoker often has a carboxyhemoglobin (COHb) level of 5% instead of the normal 0.5%. An increased level of COHb can cause increased stress on the heart, impaired time discrimination, and high blood pressure. Another harmful component in cigarette smoke is small solid particles. These particles may be fly ash, organic tar, or mineral dust. Some of these solids have been found to be carcinogenic or cancer causing. The size of these particles is a few microns (10-6 m) or less. When solid particles enter the lung, they can irritate and destroy alveoli (tiny sacs where gases are exchanged) and cause emphysema. The Surgeon General of the United States has published research indicating smoking is linked to lung cancer. A smoker may have 10 times the chance of lung cancer as that of a non-smoker. There is also research indicating that cigarette smoke not only causes damage in the lungs but also inhibits the reactions of the body to repair itself. An increased chance of heart trouble has also been linked to smoking. Not all of the smoke from a cigarette goes through the cigarette to the smoker, and a large part of the components of the smoke is exhaled by the smoker into the air. This causes an atmosphere of smoke for other people to breathe. Does this cause illness in non-smokers exposed to smoke? This is a hotly debated question and the research that has been done is not conclusive. However, a study in Japan indicated that lung cancer in women increases as the amount of smoke in the home increases. Another study concluded that chronic exposure to tobacco smoke at work is deleterious to the non-smoker and significantly reduces functioning of small airways. The most frightening research is on the effect of second-hand smoke on children. Babies of women who smoke during pregnancy tend to weigh less and develop more slowly than those of non-smoking mothers. The children of Lab Manual 11 CHE 1400 parents who smoke have been found to be more prone to bronchial ailments. WASTE AND THE ENVIRONMENT The damage to the environment that occurs during this experiment is to the air. PROCEDURE PART I: DETERMINING SOLIDS IN CIGARETTE SMOKES 1. Each pair of students will use the pre-assembled setup, shown in Figure 4.1, aimed at simulating the consumption of a cigarette and determining the quantities of solids released within the smokes. This setup comprises: - a piston-powered vacuum pump (P = 0.85 atm; air flow = 38 L/min) - vacuum rubber tubing - one T hose connector - 3 filter flasks named “Cigarette flask”, “Funnel flask” and “Pump guard” - 2 pieces of 6 μm filter paper (Whatman grade 3) - one piece of 2.5 μm filter paper (Whatman grade 5) - 3 one-hole cork stoppers - 2 funnels - one stand - 3 clamps - one cigarette. Note: The purpose of the filter flask named “pump guard” is to protect the (expensive) vacuum pump by preventing any residual fine solid from getting into it. Lab Manual 12 CHE 1400 Figure 3.1: Setup used to simulate the consumption of a cigarette. Lab Manual 13 CHE 1400 2. Remove the two pieces of filter paper from the apparatus (one from the cigarette flask, the other from the funnel flask) and weigh them. 3. Replace the two weighed pieces of filter paper in the apparatus. 4. Weigh a non-filter tip cigarette. 5. One pair of students inserts the cigarette into the funnel connected to the cigarette flask, as shown in Figure 3.1. 6. The other pair of students simply connects the second funnel into the funnel flask. This will serve to collect room smoke. 7. Clamp the cigarette funnel pointing up and the second-hand smoke funnel about a half-inch above the cigarette funnel. The funnels are to be far enough apart to allow air in but close enough together to catch all the smoke. 8. Make sure the pressure release valve of the pump is fully opened. Then, turn on the vacuum pump and light the cigarette. 9. If needed, adjust the pressure release valve so that the cigarette burns slowly, over a period of about one minute. 10. When the cigarette has burned down to the last centimeter, turn off the vacuum pump and put out the cigarette. 11. Remove the filter paper from both suction flasks and weigh each paper. 12. Weigh the unconsumed part of the cigarette. Do not weigh the ash. Subtract this weight from the initial weight of the cigarette. 13. Record all weights on the report page. 14. Calculate the number of milligrams of solids collected per gram of cigarette consumed for both smoker and second-hand smoker. 15. Repeat steps 2-14 for a filter-tip cigarette using new pieces of filter paper. 16. Subtract the initial weight of the filter paper from the final weight to obtain the weight of solids collected. For each cigarette complete steps 16-18. 17. Multiply the mass of the solids collected by 1000 and divide by the weight of cigarette consumed to obtain milligrams of solids per gram consumed. masssolid collected masscigarette ( g ) 1000 consumed (g) ________________ mg solid collected ________________ g cigarette 18. Divide the solids per cigarette of the cigarette flask (smoker smoke) by the solids per cigarette of the funnel flask (room smoke) to get the amount of solids in smoker smoke compared to solids in room smoke. Lab Manual 14 CHE 1400 PART II: SOLIDS FROM DIFFERENT PARTS OF A CIGARETTE 1. Using a pencil or marking pen, mark off the length of the tobacco part of a filter cigarette in thirds. 2. Weigh two pieces of filter paper. Place one piece in each flask. 3. Place the marked filter-tip cigarette into the funnel of the cigarette flask. Turn on the vacuum pump. Light the cigarette and as quickly as possible position the funnel over the cigarette. 4. Allow the cigarette to burn until one-third of the tobacco portion has been consumed. Extinguish the cigarette with a drop of water from an eyedropper as you turn off the pump. 5. Weigh both pieces of filter paper and record those weights on the report page. 6. Replace the 2 pieces of filter paper, turn on the pump, and relight the cigarette. Allow the second third of the cigarette to burn. When the second third is consumed, turn off the pump and extinguish the cigarette. Weigh the filter paper and record the weight on the report page. 7. Repeat step 6 for the last third of the cigarette. 8. Subtract the initial filter paper weight from the weight after one-third was consumed to find the number of grams of solids in the first third of the cigarette. 9. Subtract the weight after one-third consumed from the weight after two-third was consumed to find the grams of solids in the second third of the cigarette. 10. Subtract the weight after two-thirds consumed from the weight after last third was consumed to find the grams of solids from the last third. 11. Divide the grams of solids from the smoker smoke by the grams of solids from the room smoke for each third of the cigarette. Throw all solid residues in the trash. Be sure the cigarettes are totally extinguished. Lab Manual 15 CHE 1400 Experiment 3 Solids in smoke Name(s) Date Laboratory Instructor REPORT SHEET PART I - SOLIDS IN CIGARETTE SMOKE filter cigarette non-filter cigarette Cigarette flask Funnel flask Cigarette flask Funnel flask Initial filter paper weight (g) __________ __________ __________ __________ Final filter paper weight (g) __________ __________ __________ __________ Weight of solids collected (g) __________ __________ __________ __________ [=smoker solids] [=room solids] [=smoker solids] [=room solids] Initial cigarette weight (g) _________ __________ Final cigarette weight (g) _________ _________ Weight of cigarette consumed (g) _________ _________ masssolid collected masscigarette ( g ) 1000 consumed (g) ______ mg/g smoker solids room solids ______ mg/g _________ ______ mg/g ______ mg/g _________ 1. Which of the cigarettes in part I produced the most solids in smoker smoke and in room smoke? Can you explain the reason for your results? 2. Which cigarette produced the lowest ratio smoker solids ? room solids 3. What conclusions can be drawn from the data about being in a room with a smoker? Lab Manual 16 CHE 1400 PART II: EFFECTIVENESS OF A CIGARETTE FILTER filter cigarette Cigarette flask Funnel flask [= smoker smoke] [= room smoke] Initial filter paper weight (g) ___________ ___________ Final filter paper weight after ⅓ of cigarette consumed (g) ___________ ___________ Final filter paper weight after ⅔ of cigarette consumed (g) ___________ ___________ Final filter paper weight after a last third of cigarette consumed (g) ___________ ___________ msolids from the first third (g) ___________ ___________ msolids from the second third (g) ___________ ___________ msolids from the last third (g) ___________ ___________ smoker smoke for the first third = room smoke ___________ smoker smoke for the second third = room smoke ___________ smoker smoke for the last third = room smoke ___________ 4. Which portion of the cigarette in part II yielded the most solids in smoker smoke? Explain your result. 5. Does the portion of the cigarette affect the amount of solids in the room smoke? If so, why? 6. Does the filter protect a smoker from solids? Does a filter protect someone nearby? Lab Manual 17 CHE 1400 EXPERIMENT 4 CHEMICAL REACTIONS OBJECTIVES 1. To learn to observe and record events of a laboratory experiment and to draw conclusions from these observations. 2. To gain practice in writing and balancing chemical equations. Relates to chapter 5 of “Chemistry for changing times, 13th Ed.”. INTRODUCTION We live in a world where chemical reactions are continually taking place around and within us. Life itself is dependent on a very large number of highly complex chemical reactions. These include photosynthesis, which is the production of sugars by plants, and respiration, which is the burning of sugars by animals to produce energy. Chemical reactions are used throughout many industries to produce such products as synthetic fibers, drug, cosmetics, plastics, certain type of food detergents, fertilizers, metal alloys and many others. Numerous chemical reactions occur as part of the natural world such as the rusting of iron, tarnishing of silver, weathering of rocks, etc. A chemical reaction is a process during which a chemical change occurs. In the course of a chemical reaction, the starting material(s) undergo a chemical change, to form new chemical entities (products) that have different chemical and physical properties. Chemical reactions can be described by chemical equations. By convention, the reactants are written to the left and the products to the right of an arrow, therefore indicating the direction of the reaction. As an example, the reaction between the element sodium, Na, and chlorine, Cl2, (reactants) to form sodium chloride (product) can be written as: Na + Cl2 NaCl (not balanced) This equation is not balanced, because there are 2 chlorine atoms on the left side of the equation, but only 1 on the right side. This indicates that an atom of chlorine has been destroyed, which is a direct violation of the law of conservation of matter. According to this law, atoms can neither be created nor destroyed in a chemical reaction. If there are 2 chlorine atoms at the start of the reaction, then there must be 2 at the completion of the reaction. We cannot balance a chemical equation by changing the formulas of either the reactants or the products. This would be a violation of the law of definite proportions, which states that compounds are always found with the same proportion or ratio of each element. In the above example sodium chloride is always made up of one atom of sodium and one atom of chlorine. If we wrote the formula for sodium chloride as NaCl2, this would indicate that sodium chloride was composed of 1 atom of sodium and 2 atoms of chlorine. We can, however, balance a chemical equation by altering the number of formula units in the equation. Each formula unit of sodium chloride contains 1 atom of chlorine. Since there are 2 atoms of chlorine in the reactants, this would mean that 2 formula units of sodium chloride are produced: Lab Manual 18 CHE 1400 Na + Cl2 2 NaCl (not balanced) Now, however, there are 2 atoms of sodium on the right, and only one on the left. This would mean that 2 atoms of sodium are needed to carry out this reaction. 2 Na + Cl2 2 NaCl (balanced) As written, note that there are 2 atoms of sodium and 2 atoms of chlorine on both sides of the equation. Chemical equations sometimes indicate the physical state of the reactants and products. The most common notation is (s) for a solid, (l) for liquid and (g) for gas. The above equation could be written as: 2 Na (s) + Cl2 (g) 2 NaCl (s) In this experiment, where the emphasis is on learning how to write and balance chemical equations, these symbols will not be used. Chemical changes are usually accompanied by an observable physical change. This includes changes in color, the formation of an insoluble precipitate, the evolution of a gas, or a change in temperature. Other, more subtle changes are possible. The above reaction is an example of a synthetic reaction in which several reactants are put together to form a new compound. The opposite situation is a decomposition reaction in which a compound breaks down into elements or simpler compounds; a common example of this type of reaction is the decomposition of potassium chlorate (KClO3) into potassium chloride (KCl) and oxygen (O2) 2 KClO3 + heat 2 KCl + 3 O2 Another common type of reaction is the exchange reaction, where the cations and anions of two compounds are « exchanged ». For example, sodium bromide (NaBr) and lead nitrate (Pb(NO3)2) react to form sodium nitrate (NaNO3) and lead bromide (PbBr2) 2 NaBr + Pb(NO3)2 2 NaNO3 + PbBr2 In this reaction the PbBr2 product is a precipitate which can be noted either by underlining, with the use of an arrow pointing downwards, or by indicating the product is a solid. 2 NaBr + Pb(NO3)2 2 NaNO3 + PbBr2 (s) How do we know that the observed precipitate is PbBr2 and not NaNO3? A simple test would show that NaNO3 is very water-soluble and could not be a precipitate; which by definition is NOT water-soluble. The only other possible precipitate is the PbBr2. Although such tests are simple, it would require a considerable amount of time for you to make all the tests necessary to identify the precipitates in this experiment. Therefore, the two general solubility rules given below should be helpful. All compounds containing the cations Na+, K+ or NH4+, are water-soluble and therefore cannot be precipitates. All compounds containing the anion NO3- are water-soluble and therefore cannot be precipitates. For sections 1 and 2 of this experiment, you will mix chemicals in aqueous solution together, in small test tubes. In some cases a reaction will occur. This will be indicated by a change in the appearance (including temperature) of the mixture. In other cases there will be no evidence of a reaction. You need to make careful observations in each case, Lab Manual 19 CHE 1400 and record your observations on the data sheet. In some cases you will find that more than one reactant will react with a given reactant, in which cases you should identify the common component among the reactants. In sections 3 and 4 you will test certain chemicals with the indicators methyl orange and phenolphthalein. Carefully note any color changes during this part of the experiment and try to determine the chemical characteristics common to acids, and chemical characteristics common to bases. Table 3.1 indicates the colors given by each indicator. Table 3.1 Colors given by phenolphthalein and methyl orange Indicator Phenolphthalein Methyl orange Color when acidic colorless red Color when basic pink yellow EXPERIMENTAL PROCEDURE 1. Reaction with silver nitrate Into eight (8) 25-mm test tubes add approximately 0.5 mL of each of the 8 solutions listed below: 1) dilute hydrochloric acid HCl 2) barium nitrate Ba(NO3)2 3) sodium hydroxide NaOH 4) sodium chloride NaCl 5) hydrogen nitrate (nitric acid) HNO3 6) barium hydroxide Ba(OH)2 7) barium chloride BaCl2 8) potassium hydroxide KOH Then add about 0.5 mL of silver nitrate, AgNO3, to each of the tubes. Carefully observe to see if there is any precipitate. Record your observations in the space provided on the report sheet. If no observable reaction occurs, then report “NR” (No Reaction). 2. Reactions with sodium sulfate Into eight (8) 25-mm test tubes add approximately 0.5 mL of each of the 8 solutions listed below: 1) dilute hydrochloric acid HCl 2) barium nitrate Ba(NO3)2 3) sodium hydroxide NaOH 4) sodium chloride NaCl 5) hydrogen nitrate (nitric acid) HNO3 6) barium hydroxide Ba(OH)2 7) barium chloride BaCl2 8) potassium hydroxide KOH Lab Manual 20 CHE 1400 Then add about 0.5 mL of sodium sulfate, Na2SO4, to each of the tubes. Carefully observe to see if there is any precipitate. Record your observations in the space provided on the report sheet. If no observable reaction occurs, then report “NR” (No Reaction). 3. Reactions with Phenolphthalein Indicator Into eight (8) 25-mm test tubes add approximately 0.5 mL of each of the 8 solutions listed below: 1) dilute hydrochloric acid HCl 2) barium nitrate Ba(NO3)2 3) sodium hydroxide NaOH 4) sodium chloride NaCl 5) hydrogen nitrate (nitric acid) HNO3 6) barium hydroxide Ba(OH)2 7) barium chloride BaCl2 8) potassium hydroxide KOH Add one drop of phenolphthalein solution to each test tube and mix the solution. Note any change in the color of the solution. Record your observations in the space provided. No equation is needed. 4.Reactions with Methyl Orange Indicator Into eight (8) 25-mm test tubes add approximately 0.5 mL of each of the 8 solutions listed below: 1) dilute hydrochloric acid HCl 2) barium nitrate Ba(NO3)2 3) sodium hydroxide NaOH 4) sodium chloride NaCl 5) hydrogen nitrate (nitric acid) HNO3 6) barium hydroxide Ba(OH)2 7) barium chloride BaCl2 8) potassium hydroxide KOH Add one drop of methyl orange solution to each test tube and mix the solution. Note any change in the color of the solution. Record your observations in the space provided. No equation is needed. Note: The presence of dissolved CO2 in the water can give confusing results. If needed, add ONE DROP of the NaOH solution to any questionable test tube(s) and see if this alters your conclusions. Lab Manual 21 CHE 1400 Experiment 4 Chemical Reactions Name(s) Date Laboratory Instructor REPORT SHEET -A Reagents Equations Hydrogen chloride + silver nitrate Barium nitrate + silver nitrate Sodium hydroxide + silver nitrate Sodium chloride + silver nitrate Hydrogen nitrate + silver nitrate Barium hydroxide + silver nitrate Barium chloride + silver nitrate Potassium hydroxide + silver nitrate What ions are common to all the solutions that react with AgNO3? -B Reagents Equations Hydrogen chloride + sodium sulfate Barium nitrate + sodium sulfate Sodium hydroxide + sodium sulfate Sodium chloride + sodium sulfate Hydrogen nitrate + sodium sulfate Barium hydroxide + sodium sulfate Barium chloride + sodium sulfate Potassium hydroxide + sodium sulfate Which ion is common to all the solutions that react with Na2SO4? Lab Manual 22 CHE 1400 -C Reagents Color observed Hydrogen chloride + phenolphthalein Barium nitrate + phenolphthalein Sodium hydroxide + phenolphthalein Sodium chloride + phenolphthalein Hydrogen nitrate + phenolphthalein Barium hydroxide + phenolphthalein Barium chloride + phenolphthalein Potassium hydroxide + phenolphthalein Which ion is common to all the solutions that interact with phenolphthalein? -D Reagents Color observed Hydrogen chloride + methyl orange Barium nitrate + methyl orange Sodium hydroxide + methyl orange Sodium chloride + methyl orange Hydrogen nitrate + methyl orange Barium hydroxide + methyl orange Barium chloride + methyl orange Potassium hydroxide + methyl orange What ions are common to all the solutions that interact with methyl orange? Lab Manual 23 CHE 1400 EXPERIMENT 5 SOLUBILITY Vinegar and oil separate OBJECTIVES 1. To understand and predict which liquids will be miscible with each other. 2. To measure mass and volume and then calculate density for several liquids. 3. To observe soap as both an emulsifying agent and as a cleaning agent. Relates to chapter 6 of “Chemistry for Changing Times, 13th Ed”. BACKGROUND Molecules which are similar to each other in shape do usually dissolve in each other, that is, they are very miscible. Water is a polar molecule, which is usually immiscible with non-polar oils. Vinegar and oil dressing is an example of the denser water solution of vinegar forming a separate layer under the less dense oil layer. To blend these into one liquid, called an emulsion, requires an emulsifying agent. The emulsifier works by having one end soluble in water and the other end soluble in oil. The water and oil molecules are thus held together. Eggs are used as an emulsifying agent in mayonnaise. In an old jar of mayonnaise often a top layer of oil can be observed because some of the oil has separated. Immiscibility can be mimicked by two liquids if one is very viscous or thick. Viscosity is the ability to flow. A very viscous liquid will flow very slowly and thus will inhibit mixing and miscibility. Syrup and water will be used as an example in this investigation. The syrup and water can be mixed by putting energy into the system. The energy is added by stirring vigorously. Liquids that are miscible may dissolve in each other so well that the new volume is less than the sum of the volumes. The molecules fit inside each other and require less space together than by themselves. Another aspect of solubility is that one set of intermolecular attractions may be stronger than another. In this investigation, the attraction of oil for cloth will be stronger than the attraction of gum cloth. Thus the oil will loosen the gum/cloth attraction and allow the gum to be removed. Soap actually works much the same way in removing soil from cloth or other materials. The soil/clothe attraction is broken by stronger attractions of soap/soil. Soap acts as an emulsifying agent. Soap has a long organic end which is soluble in oil or grease and an ionic end which is soluble in water. Small balls of oil are surrounded by soap molecules with the organic end of the soap in the oil. The ionic end is then left sticking out to attract water. The entire ball is called a micelle. Watch the soap commercials on TV, which show dirt leaving a dish or cloth. The dirt is rolled into a ball. WASTE AND THE ENVIRONEMENT The solutions used in this experiment are not toxic. Lab Manual 24 CHE 1400 PROCEDURE Part I 1. Weigh a clean and dry 10-mL graduated cylinder. Fill the cylinder with exactly 10 mL of syrup and reweigh. Pour the syrup in a 50-mL beaker and set it aside. 2. Wash the cylinder, dry the outside, fill with exactly 10 mL of water and reweigh. Pour the water very carefully down the side of the beaker so that the water does not mix with the syrup. Set the beaker aside. 3. Dry the cylinder inside and out. Fill the cylinder with exactly 10 mL of cooking oil and reweigh. Pour the oil carefully down the side of the beaker. When you finish you should have three separate layers. 4. With a stirring rod mix the three layers well and set the beaker on the lab bench. Two layers will reform. 5. With an eyedropper remove enough of only the top layer to put exactly 5 mL into the 10-mL graduated cylinder. Weigh the cylinder and liquid. Record the masses in the report sheet. Clean and dry the cylinder. 6. With an eyedropper reach through the solution and remove enough of only the bottom layer to put exactly 5 mL into the 10-mL graduated cylinder. Weigh the cylinder and liquid. Record the masses on the report sheet. 7. Add 1 mL of liquid soap to the beaker. Stir well and set the beaker on the lab bench. Part II 1. Place exactly 5 mL of water in a 10-mL graduated cylinder. In a second 10-mL cylinder place exactly 5 mL of alcohol. Pour the two liquids together in the water cylinder. Read the new volume. 2. Chew a piece of gum to remove most of the sweet taste. Place the gum in a cloth square and squeeze to stick the gum to the cloth. Pour a small amount of oil and rub it into the cloth, loosening the gum from the cloth. When the gum is removed, wash the oil out with soap. Pour the solutions down the drain with plenty of water. Lab Manual 25 CHE 1400 Experiment 5 Solubility Name(s) Date Laboratory Instructor REPORT SHEET Part I Syrup Water Oil Weight of cylinder and liquid (g) __________ __________ __________ Minus weight of cylinder (g) __________ __________ __________ Weight of liquid (g) __________ __________ __________ __________ __________ __________ Density of liquid weight of liquid 10 mL Which liquid is the densest? Which liquid is on the bottom of the beaker? Which two layers mix in procedure 4? Top layer Bottom layer Weight of cylinder and liquid (g) __________ __________ Minus weight of cylinder (g) __________ __________ Weight of liquid (g) __________ __________ __________ __________ weight of liquid 5 mL Is the density of the mixture an average of the original two layers? Density of liquid Does the soap act as an emulsifying agent? Why would forming an emulsion be important when washing greasy dishes? Part II Volume of water: 5 mL. Volume of alcohol: 5mL. Measured volume of water (5 mL) + alcohol (5 mL): __________ mL Is the volume different, why? What is in oil that allows it to remove the gum? Why would using oil be better to remove gum from hair than scissors? Would peanut butter work similar to the vegetable oil to remove gum? Just like egg, soap would allow the vinegar and oil in a dressing to remain mixed, but would it be good to eat? Lab Manual 26 CHE 1400 EXPERIMENT 6 ACIDS, BASES, BUFFERS and pH OBJECTIVES 1. To gain an understanding of the relationship between pH and the concentration of acids and bases. 2. To observe the difference in properties between strong acids and weak acids. 3. To learn how a buffer is prepared and to observe how it functions. Relates to chapter 7 of “Chemistry for changing times, 13th Ed.”. INTRODUCTION Acids and bases We often encounter acidic solutions in our daily life. Many foods owe their characteristic tastes to the presence of specific acids. For example, citrus fruits contain citric acid, soft drinks often contain phosphoric acid and vinegar is little more than a solution of dilute acetic acid. Other acidic foods include tomatoes, coffee, apples and cabbage. Generally, the acids in food give them a noticeably sour taste. Our stomachs then secrete hydrochloric acid; which is needed to help to digest food. Commonly encountered bases include many cleaning agents such as soaps, ammonia, and lye (sodium hydroxide). A few foods such as egg whites are basic in nature, and many plants contain a special group of bases called alkaloids. Morphine, caffeine, nicotine and cocaine are all examples of plant alkaloids. Many commonly occurring minerals, such as limestone, magnesite, dolomite and soda ash are basic in nature. There are several definitions of both acids and bases, but the Arrhenius definition is one of the most useful. According to this definition, acids are compounds that can act as proton donors (forming hydronium ions, H3O+), and bases are compounds that act as proton acceptors (forming hydroxide ions, OH-). Strength of acids and bases When any acid reacts with water it releases the hydronium ion, H3O+, along with the corresponding anion. A strong acid is a compound that dissociates (breaks up) completely into ions when added to water. Examples of strong acids are HCl, HF, HNO3, H2SO4 and H3PO4. Since most of these acids do not contain any carbon they are often referred to as mineral acids. HCl + H2O H3O+ + ClH2SO4 + H2O H3O+ + HSO4- These reactions go totally to the right, meaning that no associated broken up HCl or H2SO4 exists in the solution. Therefore, one mole of HCl in water releases one mole of Hydronium ions along with one mole of chloride anions. One mole of H2SO4 in water releases one mole of hydronium ions along with one mole of HSO4- anions. Similarly strong base completely dissociates to form the hydroxide ion, OH- and a Lab Manual 27 CHE 1400 corresponding anion: Na+ + OH- NaOH CaOH+ + OH- CaO + H2O In contrast a weak acid, such as acetic acid CH3COOH or formic acid HCOOH, dissociate only to a small extent H3O+ + CH3COO- CH3COOH + H2O H3O+ + HCOO- HCOOH + H2O The double arrows in the equations for a weak acid denote that the reactions proceed in both directions (at different rates) at the same time. The longer arrow pointing to the left indicates that the majority of the molecules are in the associated form. Because nearly all of these acids contain at least one atom of carbon, they are often referred to as organic acids. The double arrow indicates that a given molecule, of acetic acid for instance, may give up an atom of hydrogen (H) to form a proton (H+), then later regain hydrogen from a passing hydronium. Although a given molecule of acetic acid may release and gain a hydronium ion repeatedly over time, the percentage of acetic acid molecules in the associated (CH3COOH) and dissociated (H3O+ + CH3COO-) states remain constant. One mole of a weak acid dissolved in water will generate less than one mole of hydronium ions. The amount of a weak acid that dissociates is defined by the equilibrium constant expression. This value is a property of a weak acid and relates the ratio of the moles of dissociated and associated acid present in the solution. For acetic acid (CH3COOH) this ratio is: H O CH 3COO [1] Ka 3 1.8 10 5 CH 3COOH As a first approximation it can be assured that the only source of hydronium ions is the acid and therefore (H3O+) = (CH3CCO-). It can also be assumed that only a small percentage of those acids are dissociated and therefore the value of (CH3COOH) does not change. The above equation can then be rearranged to: Ka H O 2 3 CH 3COOH H O 3 Ka CH 3COOH [2] Similarly for a weak base, such as ammonia, only a small number of the ammonia molecules dissociate into the ammonium and hydroxide ions. NH3 + H2O NH4+ + OHThe degree of dissociation can also be expressed through an equilibrium constant expression: NH 4 OH Kb 1.8 10 5 [3] NH 3 Note that, even though ammonium hydroxide is a base and acetic acid is an acid, they Lab Manual 28 CHE 1400 have the same degree of dissociation, as shown by the values of Ka and Kb. The strength of an acidic solution can be expressed by the concentration of the hydronium ion (H3O+) which is defined as the number of mole of H3O+ per liter of a solution. In pure water, a small number of water molecules dissociate into both the hydronium and hydroxide ions: to the extent that H 3O OH 1.0 107 Since the concentration of the two ions are numerically equal, the solution is considered to be neutral, neither acidic nor basic. Due to the nature of water, small concentrations of both H3O+ and OH- are always present. Numerically, this is expressed as the ion product of water and is defined as follows: [4] H 3O OH 1.0 1014 H2O H3O+ + OH- The (H3O+) from a strong acid, e.g HCl, is the same as the concentration of the acid (HCl). For example, in a 0.01 M solution of HCl, H 3O 0.01 M . For a 1.0 x10-5 M solution of HNO3, H 3O 1.0 10 5 M . Similarly, for a strong base, the (OH-) is the same as the base concentration. For example, in a 0.01 M solution of KOH, the OH 0.01 M . Since the product of (H3O+) and (OH-) ions is a constant, the concentration of the hydronium ion can be easily calculated. H O OH 1.0 10 10 1.0 10 H O 1.0OH 0.01 14 3 14 3 14 1.0 1012 Due to the large numerical range possible for the hydronium ion it is more convenient to express this value in terms of the power of the hydronium ion concentration, more commonly referred to as the pH. The pH is defined as: pH log H 3O [5] Since in pure water, H 3O 1.0 10 7 , then for a neutral solution of pure water the pH is 7.0. As the hydronium ion concentration increases, the value of the pH numerically decreases. As the hydronium ion concentration decreases, the value of the pH numerically increases. It should be remembered that the pH scale is a logarithmic one. A pH of 6.0 is ten times more acidic than water, a pH of 5.0 is 100 times more acidic and a pH of 4.0 is 1000 times more acidic. Buffer solutions Buffer solutions are solutions that tend to resist changes in pH. A typical buffer is a solution of a weak acid (such as CH3COOH) and the salt of a weak acid (such as CH3COO- Na+). If a small amount of base is added to the buffer solution, the base reacts with the weak acid to form more of the salt and water: CH3COOH + NaOH Lab Manual CH3COO-Na+ + H2O 29 CHE 1400 If some acid is added to the buffer solution it reacts with the anion of the salt (which behaves as a base, i.e. a proton acceptor) to reform the weak acid and an additional salt: CH3COO-Na+ + HCl CH3COOH + NaCl Since in either case the added base or acid is neutralized, there is very little change in the pH of the solution. The addition of acid or base to a buffered solution does change the pH of that solution, but only to a small extent. If a sufficiently large amount of base (or acid) is added so that all of the CH3COOH (or CH3COO-Na+) is consumed, then that capacity of the buffering solution has been exceeded, and it will lose its buffering ability. EXPERIMENTAL PROCEDURE 1 - pH values of common household materials Solutions of the items listed below, which are commonly found in the household, have different pH values. Ammonia Dishwashing detergent Table salt Vinegar Lemon juice Household cleaner Baking soda Oven cleaner Using a disposable pipette wet a small strip of indicator paper with one of the solutions. Determine the pH of the solution by comparing the color of the wet indicator paper with the color chart provided. Record the pH of the solution on the report sheet. Indicate if the solution is acidic (A), neutral (N), or basic (B). Continue with the remaining 7 solutions. 2- pH values of acids, bases and buffers 2.1- HCl, a strong acid 2.1.1 Put about 4 mL of 0.1 M HCl into a small test tube. Place a glass rod into the 0.1 M HCl solution, and then touch the tip of the rod on a small strip of the pH indicator paper. Determine the pH of the solution by comparing it to the color chart provided. Record the calculated (H3O+), assuming that the HCl completely dissociates. From the above (H3O+), calculate the expected pH from equation [5]. Record the observed pH. 2.1.2 Rinse a 10-mL graduated cylinder thoroughly with distilled water and then dilute 1.0 mL of sample of the 0.1 M HCl up to 10 mL with distilled water. Pour about 6 mL into a small test tube and save the remaining 4 mL. Wet a small strip of the indicator with the diluted solution from section 2.1.2. Determine the pH of the solution by comparing it to the color chart provided. Record the concentration of the HCl. Lab Manual 30 CHE 1400 Record the calculated (H3O+), assuming that the HCl completely disassociates. From the above (H3O+), calculate the expected pH from equation [5]. Record the observed pH. 2.1.3 Rinse the 10-mL graduated cylinder thoroughly with distilled water and then dilute a 1.0 mL sample of the HCl solution (saved from section 2.1.2) to 10 mL of distilled water and transfer to a small test tubes. Wet a small strip of the indicator with the diluted solution from section 2.1.3. Determine the pH of the solution by comparing it to the color chart provided. Record the concentration of the HCl Record the expected (H3O+), assuming that the HCl completely disassociates. From the above (H3O+), calculate the expected pH from equation [5]. Record the observed pH. 2.2- NaOH , a strong base 2.2.1 Put about 4 mL of the 0.001 M NaOH solution into a small test tube Wet a small strip of the indicator with the 0.001 M NaOH solution. Determine the pH of the solution by comparing it to the color chart provided. Record the calculated (OH-), assuming that the NaOH completely disassociates. Calculate the (H3O+) from equation [4]. From the above (H3O+), calculate the expected pH from equation [5]. Record the observed pH. 2.2.2 Rinse a 10-mL graduated cylinder thoroughly with distilled water and then dilute a 1.0mL sample of the 0.001 M NaOH up to 10 mL with distilled water. Pour about 6 mL into a small test tube and save the remaining 4 mL. Wet a small strip of the indicator with the diluted from section 2.2.2 solution. Determine the pH of the solution by comparing it to the color chart provided. Record the (NaOH) Record the calculated (OH-), assuming that the (NaOH) completely disassociates. Calculate the (H3O+) from equation [4]. From the above (H3O+), calculate the expected pH from equation [5]. Record the observed pH. 2.2.3 Rinse the 10-mL graduated cylinder thoroughly with distilled water and then dilute a mL sample of the NaOH solution (saved from section 2.2.2) up to 10 mL with distilled water and transfer to a small test tube. Wet a small strip of the indicator with the diluted solution from section 2.2.3. Determine the pH of the solution by comparing it to the color chart provided. Record the (NaOH). Record the calculated (OH-), assuming that the NaOH completely disassociates. Calculate the (H3O+) from equation [4]. Lab Manual 31 CHE 1400 From the above (H3O+), calculate the expected pH from equation [5]. Record the observed pH. 2.3. CH3COOH , a weak acid 2.3.1 Put about 4 mL of the 0.1 M CH3COOH into a small test tube. Wet a small strip of the indicator with 0.1 M CH3COOH solution. Determine the pH of the solution by comparing it to the color chart provided. Calculate the expected (H3O+) from equation [2]. From the above (H3O+), calculate the expected pH from equation [5]. Record the observed pH. 2.3.2 Rinse a 10-mL graduated cylinder thoroughly with distilled water, and then dilute a 1.0 mL sample of the 0.1 M CH3COOH with distilled water up to 10 mL. Pour about 6 mL into a small test tube and save the remaining 4 mL. Wet a small strip of the indicator with the diluted solution from section 2.3.2. Determine the pH of the solution by comparing it to the color chart provided. Record the concentration of the CH3COOH Calculate the expected (H3O+) from equation [2]. From the above (H3O+), calculate the expected pH from equation [5]. Record the observed pH. 2.3.3 Rinse a 10-mL graduated cylinder thoroughly with distilled water and then dilute a 1.0 mL sample of the CH3COOH solution (saved from section 2.3.2) to 10 mL with distilled water and transfer to a small test tube. Wet a small strip of the indicator with the diluted solution from section 2.3.3. Determine the pH of the solution by comparing it to the color chart provided. Record the concentration of the CH3COOH. Calculate the expected (H3O+) from equation [2]. From the above (H3O+), calculate the expected pH from equation [5]. Record the observed pH. 2.4. CH3COOH and CH3COO-Na+, a buffer solution 2.4.1 Mix 10 mL of 0.2 M CH3COOH with 10 mL of 0.2 M CH3COO-Na+. This will give 20 mL of a solution containing 0.1 M CH3COOH and 0.1 M CH3COO-Na+. Wet a small strip of the indicator with the CH3COOH/CH3COO-Na+ solution. Determine the pH of the solution by comparing it to the color chart provided. Record the observed pH of the CH3COOH/CH3COO-Na+ solution. 2.4.2 Add one drop of 0.1 M HCl to the CH3COOH/CH3COO-Na+ solution. Wet a small strip of the indicator with the CH3COOH/CH3COO-Na+ solution. Lab Manual 32 CHE 1400 Determine the pH of the solution by comparing it to the color chart provided. Record the observed pH of the CH3COOH/CH3CCO-Na+ solution. 2.4.3 Add nine more drops of 0.1 M HCl to the CH3COOH/CH3COO-Na+ solution in section 2.4.2. Wet a small strip of the indicator with the CH3COOH/CH3COO-Na+ solution. Determine the pH of the solution by comparing it to the color chart provided. Record the observed pH of the CH3COOH/CH3COO-Na+ solution. 2.4.4 Add one drop of 0.1 M NaOH to a fresh sample of the CH3COOH/CH3COO-Na+ solution. Wet a small strip of the indicator with the CH3COOH/CH3COO-Na+ solution. Determine the pH of the solution by comparing it to the color chart provided. Record the observed pH of the CH3COOH/CH3COO-Na+ solution. 2.4.5 Add nine more drops of 0.1 M NaOH to the CH3COOH/CH3COO-Na+ solution. Determine the pH of the solution by comparing it to the color chart provided. Record the observed pH of the CH3COOH/CH3COO-Na+ solution. 2.5 Water, an unbuffered solution. 2.5.1 Wash a test tube thoroughly and add about 4 mL of distilled water. Wet a small strip of the indicator with the water. Determine the pH of the water by comparing it to the color chart provided. Record the observed pH of the water. 2.5.2 Add one drop of 0.1M HCl to the water. Wet a small strip of the indicator with the water. Determine the pH of the water/acid solution by comparing it to the color chart provided. Record the observed pH of the water/acid solution. 2.5.3 Add nine more drops of 0.1 M HCl to the water/acid solution in section 2.5.2. Wet a small strip of the indicator with the water/acid solution. Determine the pH of the solution by comparing it to the color chart provided. Record the observed pH of the water/acid solution. 2.5.4 Add one drop of 0.1 M NaOH to fresh sample of distilled water. Wet a small strip of the indicator with the water/base solution. Determine the pH of the solution by comparing it to the color chart provided. Record the observed pH of the water: base solution. 2.5.5 Add nine more drops of 0.1 M NaOH to the water/base solution in section 2.5.4. Wet a small strip of the indicator with the water/base solution. Determine the pH of the solution by comparing it to the color chart provided. Record the observed pH of the water/base solution. Lab Manual 33 CHE 1400 Experiment 6 Acids, bases, buffers and pH Name(s) Date Laboratory Instructor REPORT SHEET 1. pH of common household materials. Indicate “A” (Acidic), “N” (Neutral) or “B” (Basic). 1. Ammonia solution pH: ______________ 2. Lemon juice pH: ______________ 3. Baking soda solution pH: ______________ 4. Dishwashing detergent pH: ______________ 5. Table salt solution pH: ______________ 6. Oven cleaner pH: ______________ 7. Vinegar pH: ______________ 8. Household cleaner solution pH: ______________ 2. pH values of acids, bases and buffers 2.1 HCl , a strong acid 2.1.1 2.1.2 2.1.3 (HCl) = 0.1 M Calculated (H3O+) = ______________ Expected pH = ______________ Observed pH = ______________ (HCl) = 0.01 M Calculated (H3O+) = ______________ Expected pH = ______________ Observed pH = ______________ (HCl) = + 0.001 M Calculated (H3O ) = ______________ Expected pH = ______________ Observed pH = ______________ Are the observed pH values consistent with the expected ones? Lab Manual 34 CHE 1400 2.2. NaOH, a strong base 2.2.1 2.2.2 2.2.3 (NaOH) = 0.1 M Calculated (OH-) = ______________ Calculated (H3O+) = ______________ Expected pH = ______________ Observed pH = ______________ (NaOH) = 0.01 M Calculated (OH-) = ______________ Calculated (H3O+) = ______________ Expected pH = ______________ Observed pH = ______________ (NaOH) = 0.001 M - ______________ + Calculated (H3O ) = ______________ Expected pH = ______________ Observed pH = ______________ Calculated (OH ) = How much do the pH values vary with each tenfold dilution of the base? Is this change expected? 2.3. CH3COOH , a weak acid 2.3.1 Lab Manual (CH3COOH) = 0.1 M Calculated (H3O+) = ______________ Expected pH = ______________ Observed pH = ______________ 35 CHE 1400 2.3.1.1 Are the observed pH values consistent with the expected ones? 2.3.1.2 How does the pH of this weak acid compare to the pH of a strong acid of the same molarity? 2.3.2 (CH3COOH) = + 2.3.2 0.01 M Calculated (H3O ) = ______________ Expected pH = ______________ Observed pH = ______________ (CH3COOH) = 0.001 M Calculated (H3O+) = ______________ Expected pH = ______________ Observed pH = ______________ 2.3.3.1 How do the changes in pH of a weak acid diluted 100 fold compare to the change in the pH of a strong acid diluted 100 fold? 2.4 CH3COOH and CH3COO-Na+, a buffer solution 2.4.1 Initial condition observed pH = ______________ 2.4.2 After adding 1 drop of HCl observed pH = ______________ 2.4.3 After adding 10 drops of HCl observed pH = ______________ 2.4.4 After adding 1 drop of NaOH observed pH = ______________ 2.4.5 After adding 10 drops of NaOH observed pH = ______________ What is the observed change in the pH of the buffer solution after the addition of: 1 drop of acid? ____________________________________ 10 drops of acid? ____________________________________ 1 drop of base? ____________________________________ 10 drops of base? ____________________________________ Does a solution of acetic acid and sodium acetate act as a buffer? Explain. Lab Manual 36 CHE 1400 2.5 Water, an unbuffered solution 2.5.1 Initial condition observed pH = ______________ 2.5.2 After adding 1 drop of HCl observed pH = ______________ 2.5.3 After adding 10 drops of HCl observed pH = ______________ 2.5.4 After adding 1 drop of NaOH observed pH = ______________ 2.5.5 After adding 10 drops of NaOH observed pH = ______________ What is the observed change in the pH of the buffer solution after the addition of: 1 drop of acid? ____________________________________ 10 drops of acid? ____________________________________ 1 drop of base? ____________________________________ 10 drops of base? ____________________________________ Does a solution of water act as a buffer? Explain Compare the effects of acids and bases on a buffer solution and unbuffered water. Lab Manual 37 CHE 1400 EXPERIMENT 7 ACID NEUTRALIZATION BY ANTACID How to stop heart-burn OBJECTIVES 1. To achieve neutralization in a titration reaction. 2. To observe the effects of decreasing hydrogen ion concentration on the indicator phenolphthalein. 3. To calculate the amount of acid absorbed per gram of antacid based on data gathered from titrations. Relates to chapter 7 of chemistry for changing times, 13th Ed. BACKGROUND Stomach acid is a combination of gastric juices and an acid very similar to hydrochloric acid. Sometimes eating rich food or experiencing stress may cause more than the usual amount of stomach acid to be produced. This causes discomfort to the person. Commercial antacids are primarily composed of basic or alkaline compounds and binders to hold the tablet together, but sometimes fillers are also added. The basic compounds react with or neutralize the excess stomach acid, which causes acid stomach or « heart burn ». Some of the basic compounds are hydroxides such as magnesium hydroxide, Mg(OH)2, or aluminum hydroxide, Al(OH)3. The hydroxide ion is released when the hydroxide dissolves in water. The reaction between the hydroxide ion and the acidic hydrogen ion reduces the amount of acid and relieves the discomfort. Mg(OH)2 (s) H2O Mg2+ + 2 OH- OH- + H+ H2O The hydrogen ion doesn’t actually exist by itself. It is always combined with some molecule. Usually that molecule is water and will form ions such as H3O+ or H5O2+. Because we don’t know the exact form of the ion, it is easier to just write hydrogen ion as H+ or H+ (aq). Other basic compounds used, as antacids are carbonates such as calcium carbonate, CaCO3, and sodium carbonate, Na2CO3. These react with the hydrogen ion to form carbonic acid, which quickly dissociates into water and carbon dioxide. The carbon dioxide is a gas and may cause belching which also helps to relieve stomach distress. CO32- + 2 H+ H2CO3 H2O + CO2 This investigation involves doing a back titration. To an excess amount of acid, the antacid will be added. The amount of base necessary to neutralize the solution is equal to the excess acid in the solution. The titration of an excess of reagent added is a back titration. Thus, the less base required to reach neutrality, the more acid was absorbed by the antacid. The reaction is: Lab Manual 38 CHE 1400 H+ from acid + OH- H2O from antacid Phenolphthalein is a compound that is colorless in its acidic form. When it loses a hydrogen ion to get its basic form, it gives a pink color. Since it will lose its hydrogen ion to other bases, it is used as an indicator. Which antacid neutralizes the most acid? It is time for you to find out. WASTE AND ENVIRONMENT Concentrated acids and bases can damage plumbing if not neutralized or diluted. Pour the solutions down the drain with plenty of water. PROCEDURE 1. In a clean Erlenmeyer flask, weigh 1.00 g of liquid antacid named “Maalox 1”. 2. Then, add 50 mL of 0.36 M hydrochloric acid solution to the flask. The quantity of acid added is in excess. Add 2-3 drops of phenolphthalein and a magnetic stirring bar. 3. Place the Erlenmeyer flask on the top of a magnetic stirrer and start stirring the solution for 5 sec. 1. Fill the burette with a 0.1M NaOH solution. Start adding NaOH to the Erlenmeyer flask, quickly in the range of 0-14 mL, and dropwise afterwards. A pink color can be noted and will fade quickly with stirring. Stop adding when one drop causes a pink color that does not fade within 30 seconds: the end point is reached. Record the volume of NaOH added. 2. Repeat steps 1-3 for the second liquid antacid named “Maalox 2”. Flush all solutions down the drain with plenty of water. Lab Manual 39 CHE 1400 Experiment 7 Acid neutralization by antacid Name(s) Date Laboratory Instructor REPORT SHEET Brand name of antacid Mass of antacid (g) Vbase added (mL) nacid added C ACID V ACID added 0.36 M 50 mL Antacid 1 Antacid 2 “Maalox 1” “Maalox 2” 1.00 g 1.00 g __________ mL __________ mL ____________ mmol nbase added C BASE VBASE added ________ mmol ________ mmol nacid neutralized = nacid added - nbase added = ________ mmol ________ mmol ________ mmol/g ________ mmol/g nacid neutralized per gram of antacid Lab Manual 40 CHE 1400 IV. QUESTIONS 1. Read the antacid label to find the basic compounds and write the balanced chemical equations showing the reaction with the hydrogen ion Basic compounds: ______________________________________________ Equation 1: . Equation 2: . 2. What happens to the pH of your stomach if you take more antacid than necessary to neutralise the acid? 3. Which antacid has more neutralizing power per gram? 4. Which antacid appears to be best at relieving acid stomach? 5. In recent commercials some antacids brag that they contain calcium. If calcium is used by the body to form bones and teeth, why does it matter that it is in an antacid? Lab Manual 41 CHE 1400 EXPERIMENT 8 OXIDATION AND REDUCTION Those travelling electrons OBJECTIVES 1. To observe several redox reactions and note the changes of substances from their elemental form to their ionic form and vice-versa. 2. To create a voltaic cell and measure the current it produces. Relates to chapter 8 of “Chemistry for changing times, 13th Ed.”. BACKGROUND Oxidation-reduction reactions are reactions in which electrons are transferred from one element to another. Many everyday reactions are based on oxidation-reduction (redox) reactions. Even rust is produced by a redox reaction: 4 Fe + 3 O2 2 Fe2O3 (rust) One common misconception is that water causes rust. Although objects in water will rust more quickly, water only provides the medium for the ions to travel more quickly. Copper ions and aluminium metal Redox reactions occur because one element has a stronger attraction for electrons than another element. As an example, the copper 2+ ion will take electrons from aluminum metal, producing copper metal and aluminum ions. 3 Cu2+ + 2 Al (s) 2 Al3+ + 3 Cu (s) Due to the presence of a thin and dense layer of aluminium oxide, Al2O3, on the surface of any piece of aluminium exposed to the oxygen of the air, Cu2+ ions cannot react directly with Al (s) as described in the equation above. The solution to this problem is to add Cl- ions to the solution by adding a few drops of NaCl. Chloride ions react easily with aluminium to form AlCl3. In presence of Cl- ions, the reaction is: 3 CuCl2 (aq) + 2 Al (s) Lab Manual 2 AlCl3 (aq) + 3 Cu (s) 42 CHE 1400 Iodine and zinc metal Another redox reaction occurs between zinc and iodine in solution. Zn (s) + Zn2+ + 2 I- I2 (in alcohol) purple colorless Chlorine will react with the iodide ion. 2 I- + Cl2 H+ 2 Cl- + I2 colorless purple Voltaic cell Batteries are based on redox reactions, which are arranged to cause the electrons to flow through an external circuit. Different metals can be used to set up a redox cell, which will produce a voltage between the two metals. This arrangement is called a voltaic cell. A simple voltaic cell can be produced by two metal strips (eg copper and zinc) with different affinities for electrons which are connected by an electron path. Paper towels, which have been soaked into sulfuric acid, H2SO4, can serve as an ionic bridge. The free hydrogen ions will be attracted by the copper strip (cathode) where they will be reduced into hydrogen gas, while free sulfate ions will remain in solution to balance the zinc ions released by the zinc strip (anode). Finally, a voltmeter connects the two metal strips to allow a path for the flow of electrons. The voltage read is the attraction of one metal for the electrons compared to the other metal’s attraction. WASTE AND THE ENVIRONEMENT The solutions in this investigation are not toxic. Wet zinc dust in air can burst into flames. Placing the zinc solutions or damp zinc solids on a metal pad protects from fire damage and sets up conditions for an oxidation reduction reaction which will form zinc oxide. Lab Manual 43 CHE 1400 Zinc oxide has a low toxic hazard rating. PROCEDURE 1. a) Place a small amount of zinc powder in a small beaker. Cover the metal with tincture of iodine solution and set it aside. b) Wait a few minutes. If the brown color of the solution over the zinc has faded, pour the liquid into a second beaker leaving the zinc powder behind. Add several drops of bleach to the solution. Then, add a few drops of 1 M acetic acid. Observe the change of color. 2. a) Place 25 mL of 1 M copper sulfate (CuSO4) in a 50-mL beaker. Roll a «4 by 4» piece of aluminum foil into a roll. Place the aluminium roll in the beaker and add 5 drops of NaCl solution. b) Place 5 mL of 1 M copper sulfate in a second 50-mL beaker and set it aside. c) Wait a few minutes. Remove the aluminium foil from the beaker of copper sulfate. The resulting orange/black solid is copper metal. Observe the difference in the aluminum foil. Compare the color of the two beakers of copper sulfate. 3. Make a voltaic cell using a strip of zinc metals and a strip of copper metal. Cut a piece of paper towel to about the size of a strip. Wet these pieces with 1 M sulfuric acid H2SO4. Place one strip on the bottom, then the piece of wet paper towel and the other strip on top. Call the instructor to measure the voltage produced by this voltaic cell. The aluminum foil may be thrown away in the trash. The zinc solutions from steps 1 and 5 should be placed on a metal pan to dry. The zinc oxide formed can then be buried in a landfill. Lab Manual 44 CHE 1400 Experiment 8 Oxidation and reduction Name(s) Date Laboratory Instructor REPORT SHEET I. COPPER AND ALUMINIUM 1. What does the difference in color in the copper solution indicate? 2. What happened to the aluminium foil? 3. In what form is the copper at the end of the reaction? 4. Could copper be recycled by collecting it on aluminium foil? II. ZINC, IODINE AND CHLORINE 1. In what form is the zinc (Zn(s) or Zn2+(aq)) when the purple color is gone from the solution? 2. Does bleach change the zinc ions back to the metal? 3. In what form is the chlorine (Cl2(aq) or Cl-(aq)) when the purple color appears in the solution? III. Cu/Zn/H2SO4 VOLTAIC CELL voltage: ________________V 1. Would other metal strips work? 2. Would other acidic solutions like HCl work in the voltaic cell instead of H2SO4 ? Lab Manual 45 CHE 1400 EXPERIMENT 9 CLARIFICATION OF WATER Is it clean enough to drink? OBJECTIVES 1. To understand the process of cleaning a particulate matter as one of several steps necessary to make it drinkable. 2. To produce an alternate usable compound from used aluminum cans. Relates to chapter 12 and 14 of “Chemistry for changing times, 13th Ed.”. BACKGROUND Cleaning water enough for it to be usable is a major task for most of the world. As we increase our population it becomes a larger problem. A primary method to clean water is to use a flocculent, a compound that produces a precipitate that will settle to the bottom, thus trapping suspended solids as it settles. This compound increases the speed at which suspended particles will settle and traps some particles, which would not otherwise settle. One example of this type of compound is aluminum hydroxide, a gel-like solid. The aluminum hydroxide will increase the amount of solids which settle out, and the speed at which they settle out, but the time period is often days or weeks. If we could make the aluminum hydroxide from aluminum cans, we also reduce litter. Making aluminium hydroxide from aluminium cans is not usually done as it is too expensive and making new aluminium cans is a better use for the old aluminium cans. Another method of cleaning water is to pass it over activated charcoal. This is considered an advanced or tertiary method. Activated charcoal works by adsorbing the heavy molecules in the water. The molecules are caught on the irregular surface of the activated charcoal. After cleaning the water, the charcoal can be regenerated by heating it to 500-1000 °C with steam or carbon dioxide. In this investigation, the aluminium hydroxide is compared to water that settles without the precipitate and to the original dirty water. The dirty water is produced so that it is a visible “dirty” suspension. The settling is speeded up by centrifuging so that the length of time fits into the lab time. Lab Manual 46 CHE 1400 METHOD Aluminum can be dissolved by the action of a strong base (steps [1]-[2]) to afford the aluminium compound NaAl(OH)4. Next, this compound is converted into aluminium sulphate, Al2(SO4)3, by the addition of sulfuric acid, H2SO4 (step [3]). The second goal of sulfuric acid is to dissolve any remaining aluminium hydroxide, Al(OH)3, formed in step [1]. Finally, aluminium ions present in the resulting solution are reacted with sodium hydrogen carbonate (also called sodium bicarbonate, a cheap base) to produce a large precipitate of aluminium hydroxide, Al(OH)3 (step [4]). unreacted part removed by filtration 2 Al (s) + 6 H2O (l) [1] 2 Al(OH)3 (s) + 3 H2 (g) [2] + 2 NaOH (aq) 2 NaAl(OH)4 (aq) [3] + 4 H2SO4 (aq) Al2(SO4)3 (aq) + Na2SO4 (aq) + 8 H2O (l) [4] + 6 NaHCO3 (s) + dirty water 2 Al(OH)3 (s) + 3 Na2SO4 (aq) + 6 CO2 (g) + dirt + dirt-free water removed by centrifugation Notes: - In step [3], it is not necessary to add a large amount of sulfuric acid. - In step [4], the filtrate from the aluminum sulfate solution contains aluminum ions in an acidic solution. By adding sodium hydrogen carbonate until the solution is basic, hydroxide ions are produced. These ions will precipitate with aluminum ions as aluminum hydroxide (Al(OH)3), a gel like precipitate which will trap large particles. Lab Manual 47 CHE 1400 WASTE AND THE ENVIRONMENT Acidic and basic solutions can damage plumbing unless neutralized or diluted. Aluminum ions are note classed as toxic but it is not good to put metal ions into the water system. PROCEDURE 1- Cut two 1 cm x 3 cm strips of aluminium. 2- Place the strips in an evaporating dish with 15 mL of 6 M sodium hydroxide (NaOH). Make sure the strips are completely covered. 3- Warm the dish with a hotplate in a fume hood. Observe fizzing. Heat gently to prevent boiling. Stir gently to prevent foaming over. 4- Continue until most of the fizzing stops (about 5 min.). Cool the solution. 5- Add 20 mL of 3 M sulfuric acid (H2SO4) and stir well to dissolve as much remaining aluminum hydroxide (Al(OH)3) as possible. 6- Filter the aluminum sulfate solution by suction filtration using the Buchner funnel suction apparatus as shown in Figure 9.1. Wet the filter paper slightly with distilled water to seat the filter before adding the solution. Figure 9.1 Buchner funnel with safety flask 7- Add 10 mL of the filtrate to 10 mL of dirty water in a 100-mL beaker. Add solid sodium hydrogen carbonate (NaHCO3) slowly while mixing until the solution is basic to pH paper (pH 7). Stir well and pour some of the slurry into a small test tube and mark the tube “1”. Lab Manual 48 CHE 1400 8- Although the solids will settle over a period of time, we can speed the process by centrifuging. Place about the same volume of dirty water in a second test tube marked “2”. Balance the centrifuge by placing the test tubes on opposite sides. Centrifuge for two minutes. 9- Place the same volume of dirty water in a third test tube marked “3”. Compare the three test tubes visually. Then, compare the 3 test tubes in a spectrophotometer set at 490 nm. Record the absorbance. The solutions can be flushed down the drain with a lot of water. The solids containing aluminum should be collected to be buried in a toxic waste site. Lab Manual 49 CHE 1400 Experiment 9 Clarification of water Name(s) Date Laboratory Instructor REPORT SHEET I. ALUMINUM HYDROXIDE Compare the volume of aluminum hydroxide obtained to the original volume of aluminum metal. II. USE OF ALUMINUM HYDROXIDE CLARITY OF WATER III. Absorbance of test tube with dirty water ______________ Absorbance of test tube with centrifuged dirty water ______________ Absorbance of test tube with clarified water ______________ QUESTIONS 1. What were the bubbles which formed when aluminum reacted with alkali? 2. Is it easy to see aluminum hydroxide floating in water? Explain. 3. What were the bubbles which formed when sodium hydrogen carbonate was added to the sulfuric acid solution? 4. What happens to the colorful decorations on the aluminum can when you dissolve it? 5. Is the clear water you have prepared now safe to drink? Why or why not? 6. Why isn’t a centrifuge used to clean the city water? Lab Manual 50 CHE 1400 EXPERIMENT 10 SAPONIFICATION from butter to lye soap OBJECTIVES - To understand the process of saponification. - To produce small amounts of soap from butter and sodium hydroxide. - To calculate the average molar mass of soap produced from butter. Relates to chapter 9 of “Chemistry for changing times, 13th Ed.”. BACKGROUND In the times of our grandparents and great grandparents, a wide range of cleaning products was not available. In fact the main cleaning product was lye soap. It was used for a large number of cleaning jobs: bathing, hair washing, clothes washing, and dish washing. It was usually formed into a bar but it could be cut into chips to make it more soluble for dishwashing or clothes washing. Lye soap was made by the reaction of a fat or oil with lye (sodium hydroxide, also called caustic soda). Although many soaps are now made with coconut oil, hog lard was probably more commonly used in the days of our grandparents. Palm oil or whale oil can also be used. Some restaurants have even used the grease collected in the kitchen to make soap to wash the dishes. Potassium hydroxide can be used instead of sodium hydroxide. The soap making process is called saponification. The reaction is: a fat plus sodium hydroxide produces soap plus glycerol. Before the reaction, the oils and fats are tri-esters (three ester groups). The reaction breaks the ester group apart so that a carboxylate ion (RCOO-) and an alcohol (R-OH) are formed. By removing the molecules from water the carboxylate ion combines with the sodium ion to form a solid that is called soap. After the reaction is complete, the soap is separated from the glycerol by “salting out”. O C O O R C O CH2 O R2 C O CH + O R3 C O CH2 1 A fat O C O Carboxylate Ester O R C ONa O R2 C ONa + O R3 C ONa 1 3 NaOH Soap HO CH2 HO CH HO CH2 Glycerol The solution is mixed with a concentrated sodium chloride solution. This electrolyte causes the dispersed soap to coagulate. The soap is then washed several times with concentrated sodium chloride to remove the excess lye. R is a straight hydrocarbon chain Lab Manual 51 CHE 1400 of 12-20 carbon atoms. The three R groups (R1, R2, R3) are often different from each other. In your butter sample, there are three main R groups; linoleate, oleate, and palmitate. Each produces one of the three most abundant soap molecules derived from butter, sodium linoleate (approx. 50%), sodium oleate (approx. 30%), and sodium palmitate (approx. 20%). O CH3(CH2)CH CHCH2CH CH(CH2)7CONa O CH3(CH2)7CH CH(CH2)7CONa Sodium Linoleate Sodium Oleate O CH3(CH2)14CONa Sodium Palmitate WASTE AND ENVIRONMENT The compounds in this investigation are not toxic. Acidic and basic solution can harm plumbing if not neutralized or diluted. PROCEDURE 1. Boil a mixture of 10.0 g of butter and 15 mL of 20% sodium hydroxide (NaOH) in a 100-mL or 150-mL beaker until all of the water is evaporated. Boil the mixture as strongly as possible in order to re move the water as quickly as possible. Take the following precautions to prevent injury caused by the spattering of the soap solution. Vigorous stirring with a long rod (at least 8 in.) will prevent spattering and frothing if the boiling is not too rapid. The long rod is used to keep your hand far from the beaker. Keep your head back and bellow the opening of the beaker. When it is necessary to rest from stirring, remove the burner. And continue to stir until the boiling ceases. Resume stirring before replacing the burner in one hand so that you can remove it intermittently to avoid burning the soap. One purpose of the stirring is to prevent the soap from forming a solid at the bottom of the beaker. A solid layer on the bottom will often char. 2. When all of the water has apparently been removed let the mixture cool slightly. If a waxy solid forms, the process is complete. Let the mixture cool. If a syrup liquid results, the reaction is not complete and heating must be resumed. 3. Pour 20 mL of the concentrated sodium chloride solution into the beaker containing the soap. Break up the soap into small pieces with the stirring rod so that all of the soap is washed free of glycerol and sodium hydroxide. 4. Filter the soap solution with a piece of filter paper in a funnel. Fold a piece of filter paper in half and half again (see Figure 10.1); tear off the outer corner of the fold. Lab Manual 52 CHE 1400 Open the paper up between the first and second sheet, and place the paper in the funnel. Pour the soap solution through the filter paper into a beaker. Figure 10.1 Folding filter paper for gravity filtration 5. Scrape the soap back into the beaker and repeat the washing with each of the other two 20 mL portions of concentrated sodium chloride. 6. Work the soap on a piece of dry filter paper or paper toweling to remove the last part of the wash water. Weigh the soap. 7. Show the soap to your instructor for approval. 8. Add a piece of soap about the size of a pea to 5 mL of water in a test tube and with your thumb sealing the top of the test tube, shake vigorously. Use pH paper to determine the pH of the solution. 9. Wash your hand with a small piece of your soap. 10. Place 5 mL of distilled water in a test tube. Place 4 mL of tap water and 1 mL of magnesium salt solution into another test tube. Put a small piece of soap into each test tube and shake vigorously. Now compare the lathering in both tubes. Look for a gelatinous precipitate in each test tube. The solids can be thrown in the trash. The solutions can be flushed down the drain with plenty of water. Lab Manual 53 CHE 1400 Experiment 10 Saponification Name(s) Date Laboratory Instructor REPORT SHEET I. SOAP WEIGHT: ________________ g Instructor’s approval of soap: ________________ II. QUESTIONS 1. Explain why you washed the soap with salt solution rather than water. 2. What was the pH of soap in water? 3. How do you explain this pH? 4. Describe any observations about washing your hands with a small piece of the soap. 5. In step 10, which aqueous solution allows the soap to lather more freely? 6. Saponification of your butter sample produces mainly 3 soaps: sodium linoleate, sodium oleate and sodium palmitate. Sodium linoleate results from saponification of the ester of the fatty acid named linoleic acid. What are the chemical names of the other 2 fatty acids? Lab Manual 54 CHE 1400 7. The three most abundant soap molecules produced from your butter sample are sodium linoleate, often written as C18H31O2Na; sodium oleate, written as C18H33O2Na; and sodium palmitate, written as C16H31O2Na. Calculate the average molecular weight of your butter sample. Soap component Formula Portion of the component in soap Molar mass (g/mol) Molar mass x Percentage (g/mol) sodium linoleate C18H31O2Na 50% _________ _________ sodium oleate C18H33O2Na 30% _________ _________ sodium palmitate C16H31O2Na 20% _________ _________ Average Molar mass of soap molecule = _________ g/mol The butter would be 3 soap molecules as fatty acids combined on a 3-carbon backbone. To find the average molecular weight of butter, multiply the average soap by 3, subtract 69 for the 3 sodium atoms, which are removed, and add 41 for the 3-carbon backbone. The 3–carbon backbone is glycerol without the 3 OH groups. M butter 3 M soap M Na M CH2 CH CH2 3 M soap 23 41 3 M soap 28 M butter _____________________ g/mol O R C O CH2 O R C O CH + O R C O CH2 3 NaOH Butter O R1 CONa O R2 CONa O R3 CONa Soap 8. Calculate the expected yield of soap nbutter nsoap 3 msoap 3 10.0 g msoap mbutter M butteer 3 M soap msoap 3 mbutter M soap M butter ___________ g / mol ____________ g ___________ g / mol 9. How does your actual weight compare to the calculated yield? Does your soap still have water in it? Lab Manual 55 HO CH2 + HO CH HO CH2 Glycerol CHE 1400 EXPERIMENT 11 Water analysis OBJECTIVE To learn how a water sample is collected and stored. Perform some chemical analyses on a water sample. Relates to chapter 14 of “Chemistry for changing times, 13th Ed.”. APPARATUS AND CHEMICALS conductivitimeter 150-mL beakers 100-mL, graduate cylinders Buffer solutions Vacuum filtration apparatus Analytical balance Hot plates pH meter with electrodes 100 mL, 25-mL pipets Weighing papers Filter paper INTRODUCTION Water is the universal solvent and many contaminants (impurities) are easily dissolved upon its contact. They may give water a bad taste, color, odor, or cloudy appearance (turbidity), and cause hardness, corrosiveness, or staining. They can also damage growing plants and transmit disease. At low levels, impurities generally are not harmful in water. Removing all contaminants would be extremely expensive and in nearly all cases would not provide greater protection of health. At high level (waste water) many of these impurities are treated and removed or rendered harmless. Chemists are concerned with the purity of water but regulatory agencies are concerned with setting standards to protect the environment and public health. One mean of establishing and assuring the purity and safety of water is to meet standards for various contaminants found in water. In this project you be able to get some practice on how water is collected and stored and to perform some routine chemical analysis. A. Sample Collection and Storage Site data should be recorded for all sampling locations. The information generally required includes time, date, grid references of site, weather, temperature, method of collection and information about any local activities that might influence the results. Some of these data may only be applicable for certain classes of water samples. A number of sampling devices are available for taking water samples from small ponds and from different depths in large stratified lakes. The simplest system uses a weighted bottle which is suspended at the required depth. The stopper is then removed by a sharp pull on a separate line. Water samples are especially subject to alteration in chemical composition due to microbiological activity and chemical reactions. Heavily polluted waters can undergo changes in composition within an hour of collection, and most natural waters are affected Lab Manual 56 CHE 1400 to some degree. Some tests (particularly pH and dissolved gases) should, if possible, be carried out in the field. Glass sample containers are frequently recommended for storage since polythene vessels can be porous to gaseous constituents and have been found to absorb phosphorus. In general, to minimize possible bias of results caused by any changes occurring during storage it is important to: 1 - Analyze the samples as soon as possible. Ensure that solution collectors at sites are emptied regularly. 2 - Fill containers to exclude air. 3 - Keep sample cool, but do not freeze. Physical preservation methods Fine filtration If the interest is only in the dissolved fraction then fine filtration, which removes many of the microorganisms, can be applied. It will also remove fine mineral matter and any traces of turbidity which could affect a later analytical stage. An alternative approach is centrifuging which can be used to separate various particle sizes. Temperature reduction Although some microbial activity appears to continue even at 0 °C the rapid cooling of samples after collection is generally to be recommended. Preliminary and general tests For reasons given in the previous section tests on waters should be made as soon as possible after sampling and in some cases in the field. This particularly applies to labile and gaseous constituents. Odor, turbidity, and color Odor can serve as a guide to gross pollution of water. For example, characteristic odors are associated with chlorination plants, untreated sewage and chemical industry effluents. Color in water may be a true color due to dissolved material or an apparent color when suspended material is present. The latter is quite common in natural waters, seen for example when algal blooms impart a greenish tinge. Turbidity may be used as estimate of undissolved substances in the sample. It is generally measured by visual comparison with standards or photometrically, using a neophelometer or spectrophotometer. Turbidity and color control light penetration in lake which in turn affect phytoplankton population. SOLID Total suspended solids (TSS) The finer suspended matter in natural waters is usually of an organic nature representing colloidal matter, which has been flocculated under the influence of bacteria and protozoa. Inorganic suspended matter is chiefly restricted to siliceous material resulting from the erosion of mineral soils. Lab Manual 57 CHE 1400 Total dissolved solids (TDS) It is often convenient to determine the dissolved solids in the filtrate remaining from the TSS determination. Total organic matter (TOM) TOM is all organic matter that can be found in a given water sample Alkalinity (and acidity) The alkalinity of water is its capacity to neutralize a strong acid, and the values obtained will depend on the pH of the titration end-point. In practice it is the bicarbonate, carbonates and hydroxides in solution that largely determine the alkalinity although there are minor contributions from silicates and other anions. Total alkalinity is determined by titration to the equivalent point of carbonic acid which occurs between pH 4.2 and 5.4 depending on the carbon dioxide content of water. It has long been the practice in water analysis to determine solids as dissolved, suspended and organic. Conductivity Conductivity is a property of water governed by the total ionic content. Although it is non-specific and varies with the proportion of species presents, it is often measured, because of its value in characterising waters. It expresses the resistance of 1 cm cube of water to the passage of a current, usually at 25°C (specific resistance). PROCEDURE 1. TSS determination Filtrate 100 mL of your sample and then determine the weight of the solid in the filter paper 2. TDS determination Evaporate the filtrate (liquid) to a small volume (from 100 to 50 mL) Transfer to a weighed 100 mL beaker for evaporation. Dry at 105°C to a constant weight. Cool and weigh. Express the result in mg/L. 3. TOM determination Transfer to a small pre-weighed evaporating beaker 50mL of your sample. Evaporate to dryness at constant temperature and weigh beaker plus contents. Ash in a muffle furnace, leaving at 500 °C for 1 hour. The loss in weight of the residue gives the TOM in the sample. The method is only approximate and estimates of total organic carbon are preferable when organic contents are low. 4. Alkalinity (and acidity) Measure pH of water with a pH meter If pH> 8.3 add 3-4 drops of phenolphthalein indicator and titrate against 0.01 M HCl. Lab Manual 58 CHE 1400 5. Conductivity Add the unfiltered sample into two beakers and bring to the required temperature (preferably 25 °C) by immersion in a water bath. Immerse the electrodes in each beaker that contain water samples. Record the conductivity in the second tube (having used the first as a rinse). Check the sample temperature just after immersion of the electrode. Source: A.P. Rowland & H.M. Grimshaw, in Chemical Analysis of Ecological Materials (S.T. Allen, Editor) 2nd Edition, Oxford, UK: Blackwell, 1989, p. 62. Lab Manual 59 CHE 1400 Experiment 11 Water Analysis (I) – Field Trip Name(s) Date Laboratory Instructor REPORT SHEET 1/2 Observations (in the field) Water temperature (oC) pH Conductivity γ (µmhos/cm or µS/cm) Amount of plants, or living things you observe Think about it and ANSWER these questions in the field Areas with large amount of water insects and underwater plants are warmer because these living things produce heat. Compare your results with other groups that collected water from areas with and without plants/insects, is there a temperature difference? Plants need CO2 from the water as food. When CO2 dissolves in water, an acid is produced, if CO2 is removed from water by the plants. What would happen to the pH in areas where there is plenty of plants? Does this agree with your observations of pH? The conductivity is a measure of the amount of ions in water. Plants need CO 32-, NO3- and PO43ions and sunlight to produce food. Given the conductivity measurements that you have made on the water that you took and combining it with the pH measurements, do you think that there can be life in the area where you took the water? The solubility of ionic compounds in water depends on the temperature. Higher temperatures dissolve more of the compound to produce more ions. Does your measurement of conductivity and temperature respectively, when compared to that of other groups, agree with this observation? Lab Manual 60 CHE 1400 Experiment 11 Water Analysis (II) Name(s) Date Laboratory Instructor REPORT SHEET 2/2 (Water analysis will be performed at the AUI Chemistry laboratory) Part 1 - SOLID Mass of total suspended solids (TSS) in mg ________________ Mass of total dissolved solids (TDS) in mg ________________ Mass of total organic matter (TOM) in mg ________________ Alkalinity (and acidity) pH ________________ If pH is larger than 8.3 continue otherwise skip to conductivity measurements Part 2 – Determination of Alkalinity. Concentration of standard HCl solution: 0.01mol/L Trial 1 Trial 2 0 0 1 Buret reading, initial (mL) 2 Buret reading,final (mL) 2 Buret reading, initial (mL) 3 [HCl] (mol/L) 4 Volume of HCL used (mL) _____________ _____________ 5 Amount of HCL added (mol) _____________ _____________ 6 Amount of OH- in satd solution (mol) _____________ _____________ 7 Volume of sample solution (mL) 8 [OH-], equilibrium (mol/L) Part 3 - Conductivity γ Lab Manual _____________ 0.01 25.0 _____________ _____________ 0.01 25.0 _____________ γ = _____________ µmhos/cm (or µS/cm) 61 APPENDIX I CHE 1400 Common polyatomic ions Lab Manual 62 CHE 1400 Common ions Lab Manual 63 APPENDIX II Lab Manual CHE 1400 64