Chemistry 120

General Chemistry Laboratory Manual

University of Washington

Department of Chemistry

Summer Quarter 2002

TABLE OF CONTENTS GENERAL INFORMATION.........................................................................................................1 MISSED LAB POLICY FOR 100-LEVEL COURSES..................................................................3 LABORATORY SAFETY..............................................................................................................4 THE LABORATORY NOTEBOOK..............................................................................................6 THE FORMAL REPORT ...............................................................................................................9 EXPERIMENTS ...........................................................................................................................12

HOTPACKS & COLDPACKS................................................................................................12

LINE SPECTRA.......................................................................................................................16

MOLECULAR GEOMETRY ...................................................................................................18

ANALYSIS OF HOUSEHOLD CHEMICALS........................................................................25

CHEMICAL REACTIONS ......................................................................................................28 PART A: REDOX REACTIONS & THE ACTIVITY SERIES ............................................28 PART B: PRECIPITATION REACTIONS..........................................................................32

DETERMINATION OF CHLORIDE CONCENTRATION....................................................36

CATALYSIS ............................................................................................................................39

DETERMINATION OF pH USING VEGETABLE INDICATORS.........................................44

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GENERAL INFORMATION CHEMISTRY 120

DROP/ADD: If you need to drop, add, or move sections go to the UG Stockroom, 271 Bagley Hall. The instructor has no power to do any of these things.

MATERIALS: The Chemistry 120 Laboratory Manual A bound notebook (available at the bookstore) A pair of goggles (available at the bookstore) A calculator is recommended

CHARGE #: In the hallway next to the laboratory entrance you will find a class list posted for your section. You have been assigned a charge number. Locate the station in the laboratory that corresponds to your assigned charge number. This number is your account number with the Chemistry department. You will need this number to check out items from the stockroom. Anything you purchase from the stockroom will be charged to this number.

EQUIPMENT: For each day's experiment, the teaching assistant (TA) will unlock the appropriate drawer. Check the items in the drawer against the equipment list found in this manual. If any items are missing or broken when you check-in (first 20 minutes of lab), obtain a pink slip from your TA. On the pink slip, write the name of the missing or broken item, have your TA sign the slip, then take the slip to the stockroom for a free replacement. If during the lab you break any glassware on the equipment list, take the broken glassware to the stockroom for a replacement.

Occasionally some experiments require special glassware that is not in the drawer, but which you must checkout from the stockroom. If one of these items is broken or lost you will have to purchase a replacement. You may charge these items using your charge number, but all accounts must be settled by the last scheduled lab period. If your account has not been settled by the end of the quarter, you will be billed for any charged items plus an additional $10.00 administrative fee, and a hold will be placed on your transcript until all fees are paid. The next page contains instructions for filling out the green slips that you will use to charge items from the stockroom.

Please leave the glassware clean for the next set of students. If you ever find equipment that is unusable because the previous user did not clean it, please notify your TA. A $2.00 fee is assessed to any student who leaves without cleaning up.

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HOW TO FILL OUT A LOAN-RETURN SLIP (GREEN SLIP) IN ORDER TO

CHECK OUT EQUIPMENT/SUPPLIES FROM THE UNDERGRADUATESTOCKROOM (BAGLEY 271).

No. 10University of Washington Chemistry Department

Date _____20__ Charge/Desk No. ___

Name __________________________

No. 10 LOAN CARD

Date _____20__ Charge/Desk No. ___

Name _________________________

Course __________ Section ______

For quicker retrieval, green slips arefiled by this number, not by your name.

Write down your charge number. It is

posted by your name on the bulletin board

outside your laboratory classroom.

Keep the top half of the slip. Thisportion is given to the attendantwhen returning items checked out.

The bottom half is what we file. It willbe returned to you when you return theitems you have checked out. We usethese slips to keep track of items notreturned. It is your responsibilitythat you get this half of the greenslip back when you return all items.

LIST ITEMS THAT YOUWANT TO CHECK OUTON THIS HALF.

RETURN CARD

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MISSED LAB POLICY FOR 100-LEVEL COURSES For the 100-level chemistry courses you must participate in the labs to pass the class. However, situations sometimes do occur that make it impossible to attend a scheduled lab. Missing a scheduled lab time will result from either an excused absence or an unexcused absence. Either way, you will be assigned to a different lab section if possible. For the Department of Chemistry, what constitutes an “excused absence” is: a serious accident; a death in the family; personal illness requiring a doctor’s care; participation in a University sports event; or, an event of similar gravity that can be verified. A doctor's appointment, a counseling session, a job interview, tests in another class, etc., are not excused absences--these are just scheduling challenges!

To obtain the permission form, take the proof of your situation to the stockroom. Explain what happened and ask the stockroom attendant to sign the form. The stockroom attendant will check to see if there is space available in another lab section (only 24 students are allowed to be in one lab section). If there is space available, the attendant will sign the form giving you permission to attend that section. At the beginning of that lab section, show up to the stockroom to be assigned a space to work in the lab.

If your excuse for missing lab is reasonable and documented, you will be granted an excused absence and receive full credit for performing the experiment. If, however, it is found that your excuse for missing lab is not acceptable, you will receive only 50% of your grade for that experiment. Also, if you show up more than 10 minutes late for the lab, you may be assigned to another lab section and you will receive only 75% of your grade for that experiment.

If you are granted an excused absence but the stockroom is unable to fit you into a different lab section, you will be assigned a special assignment by the Laboratory Coordinator, Dr. Frazier Nyasulu (Bagley 303B, nyasulu@chem.washington.edu). However, if you are not granted an excused absence (or you are more than 10 minutes late) and the stockroom is unable to fit you into another lab section, you will receive a grade of zero for that experiment. You cannot make up the work later in the quarter.

Words of caution:

1) You will be allowed only one absence of any kind from lab during the quarter.

2) If your regular lab section is late in the week, it is much less likely that we can fit you into another lab section.

To have your absence excused, you'll need the following:

q Proof of the situation resulting in your absence: a doctor's note, an accident report, a memorial folder, or similar documentation. The documentation must include a contact name and telephone number.

To attend another laboratory section, you'll need the following:

q A permission form filled out and signed by the stockroom (Bagley 271).

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LABORATORY SAFETY Safety goggles must be worn in the laboratory at all times. Any failure by a student to observe this state health regulation will result in loss of 10% of the available points for that laboratory, and may also result in removal of that student from the laboratory. Eyes are too valuable to risk. Students will not be allowed to work in the laboratory without approved standard laboratory goggles. Standard laboratory goggles that meet all state regulations may be purchased from the bookstore. Safety glasses, etc. are not acceptable. If you already have goggles, the stockroom personnel must first approve them before you can begin working. Because of health regulations, goggles cannot be borrowed from the stockroom.

Dress appropriately for the laboratory. In order to work with chemicals safely, bare skin should be kept to a minimum. This means socks, full shoes, long pants, and long sleeved tops that cover the hips—i.e., no gaps anywhere. The best plan is to wear a lab coat over your clothes, which you can purchase at the stockroom. FYI: cotton clothing (including denim) is particularly susceptible to being eaten by acid solutions. The laboratory is not a good place to wear your favorite clothes. Do not wear clothing so loose or bulky that it hampers your work (this is a good way to break expensive glassware). Long hair should be tied back. Failure to dress appropriately will result in a loss of available lab points, and you will be sent out of lab to acquire the correct clothing. If you do not return in time to complete your work, the absence will be unexcused.

Do not eat, drink, or smoke in the laboratory. Do not even bring these materials into the laboratory.

Wash hands often when working in lab, and always thoroughly before leaving. Do not taste any chemicals. Do not put your hands, pens, or pencils in your mouth while working in the lab.

Keep coats, backpacks and other non-essential materials away from areas where people are working. Lockers are available in hallway.

Hall lockers: You may bring a lock from home and claim an empty hall locker for use during the quarter. These are ideal for storing coats, backpacks, and other bulky items. Also, a locker is the best place to store your goggles (there have been some problems with the goggle drawers). Lockers must be emptied by the end of the quarter - between quarters the locks will be cut off and the locker contents thrown away.

Learn the location and operation of the safety showers, emergency eyewashes and fire extinguishers in the laboratory. In the case of spills onto a person or clothing, the immediate action should be water and lots of it. Do not hesitate to yell for help. Use the safety showers and/or eyewashes and don't worry about the resulting mess. Don't use the safety showers for non-emergencies since they are designed to deliver about 50 gallons of water before shutting off. Report accidents to your instructor. He/She has been certified to administer first aid. If you are not familiar with the operation of the fire extinguishers ask your instructor to explain it to you. The fire extinguishers should only be used for real emergencies since the chemicals they contain can cause considerable damage. In any emergency that requires the assistance of the fire department, aid car or police, send someone to the red emergency phones in the hallways. These are a direct line to the University Safety Division and will be answered 24 hours a day. Stay on the phone until the operator understands the situation.

Become familiar with all of the exits from the laboratory. A repeating siren and flashing of the FIRE indicator is the building evacuation signal. If this alarm goes off while you are in the lab, turn off any open flames, grab your valuables, and leave the building as quickly as possible.

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Never attempt any unauthorized or unassigned experiments. Follow the experimental procedures explicitly, checking and double-checking the identity of all reagents before you use them. There are potentially hazardous combinations of chemicals present in the laboratory. If you have an idea for further investigation, discuss it with your instructor.

Clean up spills immediately. The next person to come along has no way of knowing if the clear liquid or white powder on the lab bench is innocuous or hazardous. Neutralize acid spills with sodium bicarbonate (baking soda) before cleaning them up.

Dispose of chemical reagents and other materials properly. The proper disposal of chemical wastes is essential to the health and safety of University faculty, staff, students and the surrounding community. Chemical wastes must be managed and discarded in the most responsible and environmentally sound method available. The University and Metro expect your cooperation in taking care of the environment. Your laboratory manual will specify how to dispose of chemicals used during the laboratory period. Do not put chemicals into glass boxes or wastebaskets. Only specified non-hazardous water-soluble materials can be rinsed down the drain. Waste containers for other materials will be provided. If you are unsure of how to dispose of a particular material, ask your instructor. Metro requires that any solutions going down the drain be between pH 5.5 and 12. Therefore, neutralize any excess acid and base solution before rinsing it down the drain.

Hazard Identification: As part of the UW Laboratory Safety Manual, each laboratory has a Chemical Hygiene Plan (CHP). This is available to all students in the lab at all times. As part of the CHP, Material Safety Data Sheets (MSDS) must be readily accessible to all students. MSDS are available through the campus computer network on the Lab Safety System (LSS). The computers in the lab have a link to LSS.

Material Safety Data Sheets (MSDS) are provided by the manufacturer or vendor of a chemical. They contain information about physical properties of the chemical and identify any hazards associated with the chemical. They also identify any special handling precautions and protective equipment needed when working with the chemical. You should be familiar with the MSDS before working with any chemical.

Reagents: Read the label (contents and hazards) before using reagents. Take only as much reagent as you need - they are expensive and time consuming to prepare. When taking reagents, transfer the amount you need to a clean beaker or other suitable container for taking the material back to your desk. Replace the cap.

Never return unused reagents to the ir storage containers. If you take more than you need, dispose of the excess in the appropriate manner.

Read chemical labels carefully. Chemicals are rated from 0 to 4 according to the hazard they impose; with 0 representing no hazard and 4 representing high hazard. An example of a hazard diamond label is shown below. Each chemical is rated for health, fire and reactivity. Special warnings are reserved for the 4th diamond.

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3

2 0

Health

Fire

Reactivity

Hexane

Flammable

Do not leave a Bunsen burner or other heated apparatus unattended. The person working next to you may not know what is involved with your setup and may be working with a flammable material. Turn off open flames if you must leave your area. Make sure the gas taps are completely off whenever the Bunsen burner is not lit.

Do not pick up hot objects with your bare hands. Be sure all apparatus is cool before picking it up with your fingers.

Dispose of all broken glassware and other sharp objects into the cardboard glass disposal boxes. Custodial personnel will stop collecting trash after they find broken glass in the trashcans. If this happens, your class will be in charge of emptying the trash cans for the remainder of the quarter!

Chipped glassware and glass apparatus from your drawer may be traded for undamaged items at the stockroom. We can fire-polish chipped glassware so it is usable, but we can’t fix cut hands.

Do not adjust glass tubing connected to rubber stoppers. Severe cuts or puncture wounds may result.

Lubricate rubber tubing. When slipping rubber tubing over connectors, such as filter flasks or aspirators, lubricate with a drop of glycerin (hood) or liquid soap (in lab).

Do not force pipet bulbs onto pipets. Apply just enough pressure to maintain a seal between the pipet and the pipet bulb. Forcing the bulbs may cause the pipet to slip and break, leading to severe cuts or puncture wounds.

Always add acid to water slowly to dilute. If you attempt to add water to concentrated acid the heat of solution may vaporize the water and splash concentrated acid in your face. Sulfuric acid must be diluted particularly slowly since it releases a tremendous amount of heat upon dilution.

THE LABORATORY NOTEBOOK In all research it is essential that good records be kept - otherwise work may be misinterpreted, lost, or unnecessarily repeated. Whether in an industrial or academic laboratory, the investigator is expected to keep understandable notes written in permanent ink in a bound notebook. Observations and numerical data are written in the notebook as the work is in progress. Loose scraps of paper are unacceptable as places to record results. In many industrial laboratories, each day's work is signed by the worker and dated. Frequently the signature of the investigator is witnessed by a colleague. One reason for keeping good records is to have the information available for later use. The notes of the students in this course will be kept in a similar manner, therefore:

• A bound notebook is required. Each experiment should be dated and start on an unused page in the notebook.

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• Records should always be kept in ink. Never erase an entry. In the event that an error is made in recording an observation or data value, the erroneous words or numbers should not be erased. They should instead be crossed out and the correct information written nearby.

• Your TA will check that data is recorded in your notebook. Your notebook must be initialled by the TA before you leave the lab.

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Include the following in your notebook for each experiment:

*The first 3 sections must be done before starting the lab. Any changes should be noted during the lab.*

1. Title: name of the experiment, your name, lab partner's name, Section, and date.

2. Objective: a brief statement of what is to be accomplished.

3. Procedure : this should be a brief statement, flowchart or outline of the steps involved in your experiment. Do not write your procedure verbatim from the manual.

4. Data: This section could consist of qualitative (color changes etc.) and/or quantitative observations. All observations should be recorded in the notebook while the work is in progress. Don't worry about being neat when recording data in your notebook. If it is messy or disorganized, make a neat organized table later for your formal report.

5. Calculations : Preliminary calculations may be made on pages in the notebook and marked "preliminary" at the top. These pages will not be used for the final report, but you may want to refer to them later.

SAMPLE FROM A LABORATORY NOTEBOOK

DETERMINING THE DENSITY OF A SOLID Name: Mary Jones

Lab Partner: Henry Chang CHEM 120 - Section AA

October 5, 1995

OBJECTIVE: The purpose of this experiment is to determine the density of a solid. The density will then be used to determine the identity of the solid.

PROCEDURE: I. Measure the mass of the solid using the balance. Record estimated uncertainties II. Measure the volume of the solid by measuring the amount of water it displaces. Record estimated uncertainties.

DATA The solid was a dull grey sphere about the size of a marble. Mass of solid = 263.8±0.1 g Volume of water before solid added: 20.0 ± 0.5 mL Volume of water after solid added: 44.0 ± 0.5 mL Volume of solid: 24.0 ± 0.7 mL

CALCULATIONS: Density(D) = m/V 263.8 g/24.0 mL = 11.0 g/mL

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THE FORMAL REPORT In addition to keeping a laboratory notebook you are expected to write a formal report of each experiment. The formal report will be a more detailed, organized and critical analysis of your experiments than the notes you keep in your laboratory notebook. Your formal report should include the following: 1. Title of the experiment Include your name, TA's name, partner's name, section & date 2. Method Briefly state the purpose of the experiment or a statement of what is to be

accomplished. Discuss what measurements will be made and how those measurements lead to the results. Include a concise discussion of theory that supports the procedure. Include relevent chemical equations and mathematical formulas (but do not include the complete derivation of the equations). In the end, this section should explain how the procedure is going to answer the experimental question.

3. Data This section contains observations and data organized neatly in a table. Note any

changes from the published procedure. If data from the rest of the class is used it should be organized into a table.

4. Calculations and Graphs Show calculations that lead from data to results. Calculations should be readable. If

calculations are numerous and repetitive only one representative sample must be given. 5. Results This section summerizes any conclusions you can make based upon your data and

calculations. 6. Discussion The discussion section is a detailed evaluation of the procedure and results. Discuss

possible errors that would lead to your results being incorrect or inconsistant. Include in this section the answers to any questions that were asked in the lab manual.

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SAMPLE FORMAL REPORT

DETERMINING THE DENSITY OF A SOLID by Mary Jones

Lab Partner: Henry Chang CHEM 120 - Section AA

October 5, 1995 METHOD: The purpose of the experiment is to determine the density of a solid. The density is then used to determine the identity of the solid. The density is determined by measuring the mass and volume of the sample and subsequently

applying the equation D=m/V. The mass is measured by weighing the sample on a balance. The volume is measured by determining the volume of water that the sample displaces.

By comparing the calculated density of the sample to the densities of various elements listed as

possibilities the identity of the sample can be determined. DATA The solid was a dull grey sphere about the size of a marble. Mass of solid = 263.8±0.1 g Volume: Volume of water before solid added: 20.0 ± 0.5 mL Volume of water after solid added: 44.0 ± 0.5 mL Volume of solid: 24.0 ± 0.7 mL CALCULATIONS: Density(D) = m/V 263.8 g/24.0 mL = 11.0 g/mL RESULTS: The density of the solid was calculated to be 11.0 g/mL. This density is reasonably close to the density of lead (11.34 g/mL).

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DISCUSSION: The lab manual lists the following elements and densities as possiblities to consider: Element Density (g/mL) Copper 8.94 Iron 7.86 Lead 11.34 Nickel 8.9 Silver 10.49 Zn 7.14 From the list of possibilities only lead and silver come close to the density of the sample,

with lead being the closest. To determine between these two possiblities the appearance of the sample can be taken into account as well. The dull grey appearance of the sample is characteristic of lead.

Upon consideration of the possible errors associated with the measurements, the largest

source of error arises from misreading the volume graduations on the graduated cylinder. At best, the uncertainty in the volume measurement is only ± 0.5 mL. This error is compounded by the fact that droplets of water adhered to the side of the graduated cylinder and could have affected the readings. Also, any confusion in reading the meniscus (reading the top instead of the bottom) would have added an additional error. (Misreading the volume by only 1 mL would alter the density calculation dramatically - if the volume was 25 mL the density would have been 10.5 - much closer to silver!)

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EXPERIMENTS

HOTPACKS & COLDPACKS

In this experiment you will observe the temperature changes which occur when calcium chloride and ammonium chloride are dissolved in water. From this data you'll be able to calculate the heat given off or absorbed during the dissolution. Heat is not the same as temperature. Temperature doesn't depend on the quantity of the material it is simply a measure of how hot or cold a sample is. Heat, on the other hand, does depend on the quantity of the material and is a measure of how much energy a sample contains. For example, compare the temperature and the heat content of a cup of freshly brewed coffee and an urn of freshly brewed coffee. The coffee in each case will be approximately the same temperature but an urn full of coffee contains more heat than a cup of coffee simply because there is more coffee. Heat content is not only dependent upon the quantity of a substance it is also dependent upon the identity of the substance. Each subsance has its own specific heat which is defined as the amount of heat necessary to raise the temperature of 1 g of the substance by 1oC. The specific heat of water is 1 calorie/goC. This means it takes 1 calorie of energy to raise the temperature of 1 g of water by 1 degree Celsius. Ethyl alcohol has a specific heat of only 0.59 calories/goC. It requires much less energy to raise the temperature of 1 g of ethyl alcohol by 1oC. The third aspect to consider (along with the quantity of material and the specific heat of the material) in order to calculate the heat given off or absorbed during your dissolution reaction is the temperature change. The calculation is expressed below:

heat given off or absorbed = (specific heat)(mass)(temp. change) The particular reactions in this experiment are those which take place in commercially available hot and cold packs. The packs consist of an outer pouch which contains the solid and an inner pouch containing water. When the inner pouch is broken the reaction takes place with an accompanying release or absorption of heat. A reaction which gives off heat is called an exothermic reaction and a reaction which absorbs heat is called an endothermic reaction. MATERIALS calcium chloride, ammonium chloride, graph paper thermometer, stirring rod, 4 styrofoam cups 100 mL graduated cylinder, spatula

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SPECIAL NOTES/SAFETY CONSIDERATIONS This experiment is somewhat quantitative so be careful when weighing. This means that you don't need exactly 20 g but you do need to record the exact weight of what you have measured. he styrofoam cups should be reused. Do not throw them away! Simply rinse them out. The calcium chloride and ammonium chloride are non-hazardous and can be rinsed down the drain after finishing your experiment. Chemicals are rated from 0 to 4 according to the hazard they impose; with zero representing no hazard and 4 representing high hazard. Each chemical is rated for health, fire and reactivity. Special warnings are reserved for the 4th diamond.

0

1 0

Health

Fire

Reactivity

Calcium Chloride

0

1 0

Health

Fire

Reactivity

Ammonium Chloride

PROCEDURE 1. In your drawer you will find styrofoam cups. Construct a simple calorimeter by nesting 2

styrofoam cups, one inside the other. 2. Add 100 mL of water to the styrofoam cup. Allow the water to stand for several minutes to reach

a stable temperature. Record the temperature. 3. Weigh out approximately 10 g of CaCl2. Be sure to record the exact weight.

4. While your partner holds the cup and the thermometer steady, add all the CaCl2 to the cup and stir

rapidly with a stirring rod. Be careful not to hit the thermometer while stirring. DO NOT use the the thermometer as a stirring rod.

5. After mixing, time-temperature points should be recorded in the notebook. One partner reads the

temperature; the other reads the time and keeps the record. It is best to take temperature readings at as frequent intervals as possible when the temperature changes dramatically, then take readings

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at longer intervals as the temperature change starts to slow down. Make a plot on graph paper to find the maximum or minimum temperature.

6. After recording your data, wash the contents of the cup down the drain with lots of water and clean

the cups for reuse. 7. Rinse off the thermometer and stirring rod. 8. Repeat the above procedure using 10 g NH4Cl.

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CALCULATIONS 1. Calculate the heat gained/lost by the water during each reaction: The heat gained/lost by the water is found by finding the product of the specific heat (1.00 cal/g oC

for water) multiplied by the number of grams of water (100 g water) multiplied by the change in temperature (Tfinal-Tinitial) :

(1.00 cal/goC)( g water)(Tf - Ti)

(where Tf is the highest/lowest temperature reached after the salt is added in degrees Celsius and

Ti is the initial temperature before the salt is added).

2. One calorie is equal to 4.18 joules. What was the heat gained/lost by the water in joules during

each reaction? QUESTIONS 1. Which reaction was exothermic and which was endothermic?

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LINE SPECTRA

When sunlight passes through a raindrop a rainbow of colors called a continuous spectrum, is produced. Likewise, when white light from an incandescent bulb is passed through a prism, it is separated into a rainbow of colors. However, when light from a mercury lamp is passed through a prism we do not see all the colors of a continuous spectrum. Instead, several discrete colored bands of light called spectral lines are observed. When an atom absorbs energy from a flame or electrical discharge, it is excited to a higher energy state. When the electron returns to its original energy state, it emits the energy previously absorbed in the form of one or more photons (packages of light). The photon energy, Ephoton, is related to the wavelength and frequency of the light by the following equation:

Ephoton = h(c/λ) = hν

Where c, is the speed of light traveling through a vacuum (3 x 108 m/sec), h is Planck's constant (6.63 X 10-34 Js/photon), λ is the wavelength of the photon in meters, and ν is the frequency of the photon in s-1 (or Hertz, Hz). Each element has its own characteristic set of spectral lines. We can use this information to identify elements. For example, mercury has three very distinct lines in its visible spectrum: a violet line with a wavelength of 436 nm, a green line with a wavelength of 546 nm, and a yellow line with a wavelength of 580 nm. Also from the above equation we can get the relationship of the wavelength to the energy of the photon. The above equation states that the energy of the photon is inversely proportional to the wavelength. That is, the longer (larger) the wavelength the lower the energy of the photon. In the laboratory you will look through a spectroscope in order to see the spectral lines of mercury. The spectroscope consists of a cardboard box with a slit and ruler at one end and a diffraction grating at the other end (the diffraction grating is our stand-in for a prism). The spectroscope has not been calibrated, that will be your first task. Once the spectroscope is calibrated, you will look at several other lamps and flames to determine their characteristic line spectra. Finally, you will be asked to identify an unknown sample by its line spectrum. Procedure: Part A: Calibration of the Spectroscope All scientific instruments must be calibrated before any quantitative measurements can be made. In order to calibrate your spectroscope, darken the room. Turn on the power supply containing the mercury discharge tube. The tube will glow with a characteristic color. Look through the spectroscope at the tube to observe the line spectrum. The spectrum that you see will consist of three fairly prominent lines of light. Match each spectral line to a line on the ruler. The yellow line is 580 nm, the green line is 546 nm and the violet line is 436 nm. Plot a graph of "ruler mark" vs "wavelength". This graph is the spectroscope calibration line which will enable you to find the wavelengths of unknown spectral lines. Are there any other lines that you can see? There may be some faint ones. Identify those that you can see.

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Part B: Helium spectra Now look through the spectroscope at a helium lamp. Write down the number and colors of lines you observe. Once again match each spectral line to a line on the ruler. Use the calibration graph to determine the wavelength of each line. Part C: Strontium, Lithium, Potassium, Sodium & Barium spectra In the laboratory you will find the salts of each of the above elements. Have your partner dip a Q-tip in water and then in one of the salts. Light a bunsen burner and have your partner hold the Q-tip in the flame. Each salt will impart a distinct color to the flame. Look through the spectroscope at the flame (You'll need to be fairly close). Write down the color and ruler marking for each line. Use the calibration graph to determine the wavelength. Repeat this process for each salt using a clean Q-tip each time to avoid contamination of the salts. Part D: Unknown Identification In the laboratory you will find unknown salts. Determine the identity of one of the unknown salts by observing its line spectra. The unknown will be one of the salts you observed in Part C. Questions: 1. Would your calibration curve work for you neighbor's spectroscope? 2. Sodium emits a very simple line spectrum consisting of two orange lines with wavelengths of

589 nm and 591 nm. What is the frequency of these two spectral lines? What is the energy of these two spectral lines

3. Which has more energy, green light or orange light?

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MOLECULAR GEOMETRY

Molecules are three-dimensional objects. Paper and projection screens are two-dimensional. These simple facts-of-life mean that students of chemistry seldom get to 'see' molecules as they actually are, and consequently often have difficulty understanding why they behave as they do. In this lab you will construct three-dimensional models of a number of simple molecules using plastic balls and 'sticks', and then examine these models to see what molecular properties are a consequence of that shape. You will use a model kit to construct the molecules as they are discussed throughout the text of this exercise.

To obtain a model kit, first obtain a green equipment loan slip from your TA and fill it out completely with the date, your name, charge number, and section in the spaces provided. Also, specify that you would like to check out a model kit. Turn this loan slip along with a photo identification (student ID or driver’s license) at the stockroom window (Bagley 271) and they will loan you a model kit. Keep in mind that you are responsible for keeping track of all the pieces in the model kit and you must return the kit with the correct number of pieces.

For each model you will first draw a Lewis dot structure, including non-bonding electrons. The Lewis dot structure is a two-dimensional representation that shows the arrangement of atoms in a molecule. The Lewis dot structure includes both bonding and non-bonding electrons. In drawing covalent molecules remember that the electrons are shared between two atoms forming a covalent bond. The following rules will help you in drawing your Lewis dot structures. (*Notes to remember before you start: Remember the octet rule! Atoms tend to form bonds so that their electron configuration is the same as that of the noble gas just before or just after the atom in the periodic table. Remember that hydrogen forms only one bond- needs only two electrons.) 1. Polyatomic (more than two atoms) molecules and ions often consist of a central atom

surrounded by more electronegative atoms (hydrogen is the exception). You might look at it as the molecule tries to be as symmetrical as possible. For example, in a molecule such as carbon tetrachloride, CCl4, you find the carbon in the center and the chlorine around it.

2. Write the skeletal structure and connect the central atom to the surrounding atoms with a straight line (the bond).

3. Add up the total number of valence electrons in all the atoms in the molecules. For a polyatomic ion, add or subtract electrons to arrive at the appropriate charge (eg., if the charge is -1 you need to add an electron).

4. Place electrons about the outer atoms so that each (except hydrogen) has an octet. 5. Subtract the number of electrons assigned so far (include two electrons for each bond) from the

total calculated in rule 3. Any electrons that remain are assigned in pairs to the central atom. 6. If a central atom has fewer than eight electrons after step 5, a multiple bond is likely. Move one

or more nonbonding pairs from an outer atom to the space between the atoms to form a double bond.

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Let's do an example, the Lewis dot structure of carbon tetrachloride, CCl4.

Step 1: Choosing the central atom. Carbon is in the center and the chlorines are arranged around it.

Step 2: Writing the skeletal structure

C

Cl

Cl

Cl

Cl

Step 3: Total number of valence electrons: 4 for carbon 7 for each chlorine = 28 total = 32 Step 4: Placing electrons We have used two electrons in each bond so there are 24 electrons left and we place

those around the outer atoms.

C

Cl

Cl

ClCl

Step 5: We have placed all the electrons around the four chlorine atoms and discover that all

the atoms now have a full octet, so we have finished.

Once you have drawn the Lewis dot structure for your molecule, use the model kit to build the molecule. The balls in the model kit are color-coded according to a rather traditional scheme: H is white, C is black (like graphite), N is blue (the color of the sky?), O is red (the color of fire), and Cl is green (the color of the element). There are two varieties of 'sticks': short ones and long ones. The long sticks are used when making double and triple bonds, you will discover why as soon as you try to make a double bond using the short sticks. The short sticks are used for single bonds. The problem with a Lewis dot structure is that it is a two-dimensional representation of a three dimensional object. The model that you build will give you a better idea of the actual shape of the molecule. You will find that the shapes of molecules is dependent on more than the number of atoms.

20

THREE-ATOM MOLECULES Let's consider the examples of CO2 and H2O, which both contain three atoms. Draw the Lewis dot structure for carbon dioxide, CO2 and water, H2O CO2 H2O Using the model kit, build CO2 and H2O. In the space below draw (to the best of your abilities) a "more accurate" representation of the actual shapes of these molecules. CO2 H2O The red oxygen ball contains only two holes since the two lone pairs of electrons occupy the other two spaces. Build another water molecule using a black ball with four holes for the oxygen and attaching two hydrogens to it. Does it make any difference in the shape of the molecule where you put the two hydrogens? Explain why water is not linear.

21

MORE-THAN-THREE-ATOM MOLECULES Next, lets look at a set of similar molecules, ammonia (NH3) and ammonium ion (NH4+). As the book explains, the non-bonding electrons, even though they are 'invisible', play a crucial role in determining the shape of the molecule. Draw the Lewis dot structures for the two molecules. ammonia (NH3) ammonium ion (NH4+) Use the model set to construct ammonia and ammonium ion. Draw the shapes below. In drawing the shapes of the molecules make every attempt to give a three dimensional perspective to your drawing. ammonia (NH3) ammonium ion (NH4+) In constructing ammonia, you could use 'a stick' to represent the non-bonding electrons, but usually these important electrons are simply ignored when building models of molecules. Ammonium ion is a polyatomic ion, a charged particle containing two or more covalently bonded atoms. When ammonia is dissolved in water, it reacts in water solution such that it picks up a hydrogen ion, H+. The lone pair of electrons on ammonia and the hydrogen ion unite to form a covalent bond. If you have used "a stick" to represent the non-bonding electrons you will see that you just need to add the hydrogen ball (representing H+) to build the ammonium ion. A very important non-planar geometry is the tetrahedron. This is the geometry of molecules like NH4+, which have bonds that point toward the opposite corners of a cube. Since carbon forms four bonds the structure around carbon atoms is tetrahedral when only single bonds are formed. Draw the Lewis dot structure for methane (CH4) and build the molecule. CH4 Lewis dot: CH4 Shape:

22

POLARITY OF MOLECULES An important characteristic that plays a role in the behavior of molecules is the distribution of electrons. When the electrons are evenly distributed between two atoms, as in fluorine, F2, the bond is called non-polar. When two, dissimilar, atoms are bonded, the electrons migrate toward the atom with the higher electronegativity (attraction for electrons) and the bond is polar. In HCl, for example, the stronger effective nuclear charge from the chlorine nucleus pulls the electrons in the bond closer to the chlorine nucleus and further from the hydrogen nucleus. The net effect is a slightly more negative charge near the chlorine, which is balanced by a slightly more positive charge near the hydrogen. Remember that these molecules are neutral overall. All bonds that involve more than one kind of atom behave this way, some to a greater extent than others. Molecules that contain the highly electronegative atoms (oxygen, fluorine, chlorine) will have a greater separation of positive and negative charge. The little 'extra' charges are indicated by writing 'δ-' near the electronegative atoms, and 'δ+' near the 'electropositive' ones. For a molecule with more than two atoms the polarity of each bond and the arrangement of the bonds needs to be considered. Since there are two atoms involved in the polar bond, the phrase 'dipole moment' is also used ('dipole' means that there are two 'poles', or charges: 'δ-' and 'δ+'). Almost every bond in a molecule has some polarity, which is often indicated by drawing a little arrow next to the bond, pointing from the '+' atom to the '-' atom. (The '+' sign is created by drawing a line through the tail of the arrow: +-->). The overall effect of these individual polar bonds depends on the symmetry of the molecule. If the molecule has sufficient symmetry the individual polar bonds cancel each other, leaving a molecule that appears to have a uniform charge distribution. If, on the other hand, the molecule has some sort of irregular shape or if the molecule is not symmetrical, the polar bonds will not cancel, leaving the molecule with a net, or permanent, dipole moment. Build the following molecules: chloromethane, CH3Cl carbon tetrachloride, CCl4 Using the information from the models you have built, explain why chloromethane has a permanent dipole, while carbon tetrachloride does not. Examine the other models you have built so far, which of these models would have a permanent dipole?

23

ISOMERS Another thing to watch for when you are building a model is whether or not you get a different molecule when you interchange some of the atoms. As an example, consider molecules with the formula C2H2Cl2. Build a model for the following structural formula for C2H2Cl2.

C=C

H

Cl

H

Cl How many different ways can you arrange the hydrogens and chlorines to get different molecules? Save your original model to compare to make sure the molecules are really different. Draw below the structural formulas for the molecules you believe are different. Different molecules constructed from the same set of atoms or, put another way, different molecules with the same molecular formula, are called isomers . You will run into isomers again, if and when you take organic chemistry. Double bonds (as you will see when you build the model) always keep the atoms 'locked in place'. Molecules whose atoms are connected by single bonds tend to be rather 'floppy'. This is because single bonds, unlike double bonds, allow the atoms to rotate. The rate at which such rotation occurs is almost unbelievable: the two CH3 groups in the ethane molecule CH3-CH3 rotate around the

bond that joins them at a rate of approximately 107 times a second! Construct a model of ethane, CH3-CH3, and compare this model with one of the models with a double bond from above. Notice that the double bond prevents the free rotation around the carbon-carbon bond. Draw the structure for ethane, CH3-CH3

24

POLYMERS A polymer is a large molecule which is made out of a series of connected repeating units. In this last exercise each member of the lab section will make a monomer (one of the repeating units) and you will all "react" with each other to form a polymer. First construct the monomer, ethylene (CH2=CH2 ). Connect your monomer to a neighbor student's monomer by breaking one of your double bonds and connecting it to your neighbors ethylene. You will notice that in order for you to bond to your neighbor, they will have to break a double bond - which they can use to bond to their neighbor. This is a chain reaction. You have made polyethylene! Draw the monomer for polyethylene Is this an a-a-a-a-a-a or a-b-a-b-a-b-a-b type of monomer? Polyethylene has up to 50,000 repeating units, how many repeating units were you able to make in lab?

25

ANALYSIS OF HOUSEHOLD CHEMICALS

Imagine finding an unlabeled white powder on your kitchen table. Looking around the kitchen you realize that the white powder could be one of a myriad of chemicals. (You wouldn't want to mistake calcium sulfate also known as plaster of paris for sugar--BLEECH.) With this type of problem in mind, several chemists at Drexel University (1) developed an analysis scheme to identify household chemicals. They included: 1. sugar (sucrose), 2. salt (NaCl), 3. baking soda (sodium bicarbonate), 4. cornstarch, 5. washing soda (sodium sulfate), 6. plaster of paris (calcium sulfate), 7. epsum salts (magnesium sulfate), 8. chalk (calcium carbonate) The analysis scheme includes the following tests: 1. Solubility in water 2. Reaction with iodine (starch and iodine react to form a blue solution) 3. Reaction with vinegar (carbonates will react with acetic acid releasing bubbles of carbon dioxide) 4. Reaction with phenolphthalein indicator (basic solutions turn the indicator pink) 5. Precipitation with sodium hydroxide solution (hydroxide ions react with magnesium ions to produce

a precipitate) 6. Conductivity (soluble ionic compounds conduct electricity) A flowchart on the next page organizes these tests for you. You will use this flowchart to analyze and identify several unknown white powders. Reference: (1) Sally Solomon, Annamaria Fülep-Poszmik, and Alan Lee, Journal of Chemical Education, 68,

328 (1991).

26

starch, CaSO4, CaCO3, NaHCO3,NaCl, sucrose, MgSO4, Na2CO3

1. solubility in waterInsoluble soluble

starch, CaSO4 CaCO3

NaHCO3, NaClsucrose, MgSO4, Na2CO3

2. reaction with iodine

starch

blue brown

CaCO3, CaSO4

4. reaction with phenolphthalein indicator

intense pink colorlessfaint pink

Na2CO3 NaHCO3 sucrose,NaCl,MgSO43. reaction

with vinegarbubbles no bubbles

CaCO3 CaSO4 5. reaction with sodium hydroxide

formsppt.no ppt.

MgSO4sucrose,NaCl

NaCl sucrose

6. Conductivity test

conducts does not conduct

27

MATERIALS Test tubes & holder 10 mL graduated cylinder Spatula Disposable pipettes small beaker conductivity tester

CHEMICAL TEST SOLUTIONS Iodine solution ("tincture of iodine") Phenolphthalein ("Exlax" solution) Acetic Acid ("Vinegar") Sodium hydroxide solution ("Drano" solution)

SAFETY CONSIDERATIONS Although all the chemicals are household materials the waste for this experiment will be collected. PROCEDURE

Step 1. Take a small amount (pea-sized) of the solid and place it in a test tube. Add about 5 mL of water. Agitate the test tube for three to five minutes. Does the solid dissolve readily? If insoluble, go to step 2 below; if soluble, go to step 4 below. For your unknown, list the possible compounds.

Step 2. Add two drops of the iodine solution to the test tube. Is the solution blue or brown? If the

solution is brown, go to step 3; if blue, stop here. For your unknown, list the possible compounds.

Step 3. Place a small amount of the solid in a clean test tube. Add about 1 mL of white vinegar

(solution of acetic acid). Are bubbles produced? No more analysis is required on this sample. For your unknown, list the possible compounds.

Step 4. To the solution from step 1, add 2 drops of phenolphthalein solution. Does the solution turn

pink? If the solution is pink, no more analyses is needed. If the solution is colorless, go to step 5. For your unknown, list the possible compounds.

Step 5. Place a small amount of the solid in a clean test tube. Add about 5 mL of water. After the

solid dissolves, add several drops of 0.25M sodium hydroxide solution using a disposable pipet. Did a precipitate (solid) form? If no precipitate formed, go to step 6; if a precipitate formed, no more analysis is needed. For your unknown, list the possible compounds.

Step 6. Place a small amount of the solid in a beaker and add about 5 mL water. Dip the electrodes

of a conductivity tester in the solution. Does the solution conduct electricity? For your unknown, list the possible compounds.

28

CHEMICAL REACTIONS

PART A: REDOX REACTIONS & THE ACTIVITY SERIES One of the most important characteristics of a metal is its activity (its reactivity), or its ability to lose electrons and become an ion. Metals range widely in activity, from vigorously reactive cesium, potassium and sodium to quite inactive platinum, gold and silver. In this experiment you will rank some metals according to their activities, from most active to least active, to produce an activity series. To do the experiment you will observe whether a redox reaction occurs between a metal and the cation of another metal. For example, if copper metal is placed into a solution of silver nitrate a redox reaction occurs: crystals of silver metal appear and the solution turns blue because of copper ions appearing in the solution. The reaction can be written:

Cu (s) + 2AgNO3(aq) => Cu(NO3)2 (aq) + 2Ag (s) In this example copper is more active than silver because it donates its electrons to silver ions. In order to show exactly where electrons are lost and gained the equation can be divided into two half reactions:

Cu => Cu2+ + 2 electrons (copper loses electrons, it is oxidized) Ag+ + 1 electron => Ag (silver ion gains an electron, it is reduced)

What would you expect if instead of copper metal placed in silver nitrate solution, silver metal was placed in copper nitrate solution? No reaction occurs because silver is less active than copper. With this in mind we can rank the activity of copper and silver as:

Copper Silver

Acids may also enter into redox reactions with active metals to produce hydrogen gas and a salt. For example, aluminum reacts with hydrochloric acid:

2Al (s) + 6HCl (aq) => 2AlCl3 (aq) + 3H2 (g)

Silver on the other hand does not react with Hydrochloric acid: Ag (s) + HCl (aq) => no reaction

Metals which are more active than hydrogen (like aluminum) will displace hydrogen from acids, and those less active (like silver) are unable to do so. With this in mind we can rank the activity of aluminum, hydrogen and silver as: Aluminum

Hydrogen Silver

In this experiment you will rank the activity of zinc, magnesium, copper and hydrogen. Materials zinc, magnesium and copper metal, 0.5M solutions of CuSO4, ZnSO4, MgSO4, HCl small test tubes test tube rack

29

Safety: Chemicals are rated from 0 to 4 according to the hazard they impose; with zero representing no hazard and 4 representing high hazard. Each chemical is rated for health, fire and reactivity. Special warnings are reserved for the 4th diamond. In this experiment avoid contact with corrosive hydrochloric acid. All wastes should be placed in the waste container provided. Do not pour any solutions into the drain.

Health

Fire

Reactivity Health

Fire

Reactivity0

0

0

Copper

0

2

2

Magnesium

Health

Fire

Reactivity0

1

2

Zinc

30

0

2 0

Health

Fire

Reactivity

Copper sulfate 0.5M

0

1 1

Health

Fire

Reactivity

zinc sulfate 0.5M••••

0

3 2

Health

Fire

Reactivity

Hydrochloric Acid 0.5M

0

1 0

Health

Fire

Reactivity

magnesium sulfate 0.5M••••

Corrosive

31

Procedure Obtain small pieces of copper, magnesium and zinc metal. In separate test tubes treat small pieces of each metal with 2 to 3 mL of the various solutions as outlined below. a) Zn and MgSO4 b) Zn and CuSO4 c) Zn and HCl d) Cu and MgSO4 e) Cu and ZnSO4 f) Cu and HCl g) Mg and CuSO4 h) Mg and ZnSO4 i) Mg and HCl Results After 5 to 10 minutes observe the color of each solution and examine each metal surface for evidence of any deposit. Write a balanced equation for any reaction detected and then rank copper, zinc, hydrogen and magnesium according to their activity (from the most active to the least active).

32

CHEMICAL REACTIONS

PART B: PRECIPITATION REACTIONS Ionic solids are made up of individual ions held together by electrostatic forces. When such solids are dissolved in water they dissociate and the ions separate and become essentially independent of one another. Identification of the ions present in a solution can often be made because many ions have sufficiently distinguishable properties, thus they can be characterized chemically. In this experiment you will determine the behavior of a number of different kinds of ions then use this information in analysing an unknown substance. You will study the following ions: Negative ions or Anions: CO32- Carbonate (from 0.1M sodium carbonate solution)

OH- Hydroxide (from 6 M sodium hydroxide solution) I- Iodide (from 0.1M sodium iodide solution) SO42- Sulfate (from 0.1M sodium sulfate solution) Positive ions or Cations Ba2+ Barium (from 0.1M barium nitrate solution) Cu2+ Coppper (from 0.1M copper nitrate solution) Pb2+ Lead (from 0.1M lead nitrate solution) Ni2+ Nickel (from 0.1M nickel nitrate solution) Each of these anions will be mixed with each of these cations and you will observe whether or not a precipitation reaction takes place (a reaction in which a solid is formed).

33

Safety: Chemicals are rated from 0 to 4 according to the hazard they impose; with zero representing no hazard and 4 representing high hazard. Each chemical is rated for health, fire and reactivity. Special warnings are reserved for the 4th diamond. In this experiment avoid contact with caustic sodium hydroxide solution Barium nitrate, copper nitrate, lead nitrate and nickel nitrate are toxic by ingestion. All wastes should be placed in the waste container provided. Do not pour any solutions into the drain.

Health

Fire

Reactivity

1

0

1

Sodium Carbonate 0.1M

Health

Fire

Reactivity

1

0

1

Sodium Iodide 0.1M

Health

Fire

Reactivity

3

0

2

Sodium Hydroxide 6M

Caustic

Health

Fire

Reactivity

3

0

3

Nickel Nitrate 0.1M

Toxic

34

0

0 0

Health

Fire

Reactivity

Sodium Sulfate 0.1M

0

3 3

Health

Fire

Reactivity

Barium Nitrate 0.1M••••

Toxic

0

3 3

Health

Fire

Reactivity

Lead Nitrate 0.1M

0

1 3

Health

Fire

Reactivity

Copper Nitrate 0.1M••••

Oxidizer

Toxic

35

Procedure You will be given a 24 well plate that is arranged into six columns (numbered 1 through 6) and four rows (lettered A thru D). Fill each of the four wells of column 1 half full with sodium carbonate solution. Likewise fill each of the four wells in column 2 with sodium hydroxide solution (Caution! Caustic! Avoid contact with skin eyes or clothing.) Fill each of the wells of column 3 with sodium iodide solution, and each of the wells of column 4 with sodium sulfate. Now in row A add barium nitrate solution into each of the six wells. In row B add copper nitrate solution to each of the six wells, in row C add lead nitrate solution to each of the wells and in row D add nickel nitrate to each of the wells. You can record your observations on a grid like the one shown below. Once you have completed the above reactions, you can test an unknown anion solution. Use column 6 for your unknown. By mixing your unknown with each of the cations you should be able to learn its identity.

CO32- OH- I- SO42- unknown anion

Ba2+

Cu2+

Pb2+

Ni2+

Results Include in your results section a balanced equation for each precipitation reaction.

36

DETERMINATION OF CHLORIDE CONCENTRATION

A common laboratory procedure to determine the concentration of a solution of ions is called a titration. The chloride concentration in river or sea water can be determined by titrating a sample with a solution of silver nitrate. In the titration, chloride ions react with silver ions to form an insoluble white precipitate of silver chloride.

Cl-(aq) + Ag+(aq) => AgCl (s) The balanced chemical equation tells us the number of moles of chloride ions required to react completely with a given number of moles of silver ions. To perform this titration, a known volume of chloride solution is placed in a flask using a pipet. Several drops of potassium chromate indicator are then added to the flask. The indicator imparts a yellow color to the solution. Silver nitrate solution of a known concentration is then slowly dispensed from a buret into the flask until there are just enough silver ions present to completely react with the all the chloride ions. With the addition of one more drop of silver nitrate the silver ions now react with chromate ions to form an orange-red precipitate of silver chromate. The appearance of this orange-red precipitate signals the endpoint of the titration.

CrO42-(aq) + 2Ag+(aq) => Ag2CrO4 (s)

At the endpoint the number of chloride ions originally present exactly equals the number of silver ions added. If the volume of the chloride sample was carefully measured and if the volume of silver nitrate added and its concentration are known, it is possible to calculate the chloride concentration.

Molarity of chloride solution =

(Molarity of silver nitrate solution ) (Vol. of silver nitrate solution )Vol. of chloride solution

For example, imagine that 10.00 mL of sea water are measured into an erlenmeyer. From a buret 9.30 mL of a 0.5040 M solution of silver nitrate were added to the sea water in order to reach the endpoint. From this titration data the chloride concentration in the sea water can be calculated:

Molarity of chloride solution =

(0. 5040M silver nitrate ) (9.30 mL of silver nitrate )

10.00 mL . of chloride solution= 0. 4687 M

In this experiment various local water samples will be analyzed for chloride ions. Water samples were collected from the following sites: #1 Alki Beach: This is sea water from Puget Sound

#2 Harbor Island: This water is a mixture of sea water from Elliot Bay and fresh water from the Duwamish River

#3 Duwamish River at the First Ave. Bridge A map showing these sites will be posted in the lab. Each pair of students will be assigned to analyze water from two of these sites. Students are then expected to share their data with the rest of their section.

37

Materials 50 mL erlenmeyer flask (2) 50 mL beaker 10 mL pipet and bulb 10 mL buret (from stockroom) Silver nitrate solution of known concentration Chloride solution of unknown concentration Potassium Chromate indicator (0.1M) Special Notes/Safety Considerations The silver nitrate solution will stain skin and clothing. The potassium chromate solution is toxic if ingested. All titration waste must be collected in the waste container provided. This lab manual contains hazard ratings for the chemicals you will be using. Chemicals are rated for health, fire and reactivity from 0 to 4 according to the hazard they impose; with zero representing no hazard and 4 representing high hazard. Special warnings are noted in the 4th diamond.

Health

Fire

Reactivity Health

Fire

Reactivity2

0

0

Silver Nitrate

0

0

0

Chloride Solution

Health

Fire

Reactivity4

0

1

Potassium Chromate Indicator

Toxic

Procedure 1. Obtain a green equipment loan slip from your TA and fill it out completely with the date, your

name, charge number, and section in the spaces provided. Specify that you wold like to check out a 10 ml buret. Exchange this slip for a buret at the stockroom window (Bagely 271). Use a buret clamp to secure a 10 mL buret to a ringstand. Thoroughly rinse the 10 mL buret with deionized water. The deionized water should drain evenly from the inside surfaces of the buret and leave no droplets of water behind.

2. Obtain about 25 mL of silver nitrate solution in a small beaker. Note the concentration of the

solution. 3. Fill the buret with silver nitrate solution. Open the stopcock long enough to fill the tip with

solution and remove any air bubbles. Add solution to the top of the buret, or drain the solution through the stopcock, to bring the level close to 0 mL on the scale. It is not necessary that the initial level be exactly 0 mL, but it must be at or below the start of the scale.

38

4. Record the initial level of the buret, keeping your eye level with the bottom of the meniscus. Estimate the last digit .

5. Obtain approximately 30 mL of chloride solution from one of the sites you have been assigned

to analyze. Rinse the 10 mL pipet with two small portions of the chloride solution then use the pipet to accurately transfer 10 mL portions of this solution into each of two clean 50 mL Erlenmeyer flasks.

6. Add several drops of potassium chromate indicator to each flask. 7. Titrate the chloride solution to the endpoint, the first permanent appearance of an orange-red

precipitate. You can anticipate the approaching endpoint by observing how long the small trace of red lingers before disappearing. As you approach the endpoint, continue the addition of silver nitrate solution slowly - eventually drop by drop - mixing carefully after each addition. With practice, it is possible to determine the endpoint with the precision of one drop.

8. Record the final volume in the buret. 9. Repeat the titration on the other 10 mL chloride sample. Check the volume of silver nitrate

you have left in the buret. You may wish to refill the buret with more silver nitrate solution. 10. If the volumes of silver nitrate delivered for the two trials are not within reasonable agreement,

repeat the titration. 11. Repeat the titration procedure for the water from the other site that you have been assigned to

analyze. Calculations From your titration data calculate the chloride concentrations. Write your results on the chalkboard & notify your TA. Results Each of the five sites will have been analyzed by someone in your lab section. The results will be posted on the chalkboard in the laboratory during your assigned lab time and then subsequently posted on the Chem 120 bulletin board. In your lab write-up you should account for the variation in concentrations. Questions 1. What would happen to the color if you were to add a few drops of potassium chloride solution

to the final solution after one of your titrations? Explain your answer. 2. Are there any anions (other than Cl-) which might interfere with your analysis in this experiment?

Is there an easy way to check whether a particular anion would affect your titration? 3. Can you speculate how far up the Duwamish River the tides comes?

39

CATALYSIS Substances which influence the rate of a chemical reaction without undergoing any permanent change themselves are called catalysts. You're probably familiar with catalysts already. For example, since 1976, every new car and truck sold in the United States has a "catalytic converter". This catalyst (finely divided palladium, platinum or rhodium on a ceramic support) is designed to reduce the emission of the smog-causing compounds carbon monoxide, nitrogen oxides and hydrocarbons and transform them into less harmful compounds. For example, the conversion of carbon monoxide to carbon dioxide is normally quite slow in an automobile engine. The catalytic converter is able to take this slow reaction and speed it up.

2 CO +O2 => 2CO2

Catalysis is also occurring up in our stratospheric ozone. Stratospheric ozone protects the earth from dangerous ultraviolet radiation. Since 1970 scientists have noted a dramatic decrease in stratospheric ozone, particularly over the South Pole. The destruction of ozone is catalyzed by chlorine atoms. When a chlorine atom encounters an ozone molecule the following reactions occur:

Cl + O3 => ClO + O2 ClO + O => Cl +O2

Notice that in the first step chlorine atoms are reactants but in the second step they are products (and vice versa for the ClO). In this way the chlorine is regenerated to react with more ozone molecules. On the average, a single chlorine atom may be involved in destroying as many as 100,000 ozone molecules before it is carried back to the lower atmosphere. One of the sources of chlorine atoms in the stratosphere comes from the degradation of manmade chloroflurocarbons (CFC's). In response to the problem of stratospheric ozone depletion approximately 100 nations have agreed to ban the use of CFC's by the year 2000. Whether in catalytic converters or the stratosphere we've seen that relatively simple molecules, atoms or ions can function as catalysts. In biological systems, however, the catalysts are often very large, complex protein molecules called enzymes that may have molecular weights from several thousand to a million or more. The exact mechanism for the action of catalysts is not completely understood in all cases. However, it is clear that catalysts provide an alternate path or series of steps by which the reaction can take place so that some slower steps in the uncatalysed reaction are bypassed. Some reactions are catalyzed by more than a single substance. The decomposition of hydrogen peroxide to produce water and oxygen gas proceeds slowly at room temperature, but the rate of this reaction can be markedly increased by the presence of iron ions (Fe3+). In biological systems the ubiquitous enzyme, catalase, also catalyses the decomposition of hydrogen peroxide. Catalase is an enormous protein molecule with a molecular weight of 250,000, but the small portion of it that functions in catalysis appears to involve the four iron ions which are present in each molecule.

2 H2O2 Fe3+

→ 2 H2O + O2 In this experiment you will observe the effect of a catalyst by determining the rate at which hydrogen peroxide decomposes in the presence of a catalyst. The reaction rate is monitored by the measurement of the volume of water displaced by the oxygen gas formed.

40

MATERIALS Solutions: hydrogen peroxide solution 3% 1.0 M ferric chloride solution 0.5 M ferric chloride solution catalase solution From you drawer: pipet bulb 10 mL & 100 mL graduated cylinder small beaker pinch clamp, oxygen generator, tubing assembly, 500 mL Florence flask SAFETY CONSIDERATIONS Hydrogen peroxide solution can irritate skin and bleach clothing so avoid contact. Waste solutions may be rinsed down the drain. This lab manual contains hazard ratings for the chemicals you will be using. Chemicals are rated from 0 to 4 according to the hazard they impose; with zero representing no hazard and 4 representing high hazard. Each chemical is rated for health, fire and reactivity. Special warnings are reserved for the 4th diamond.

Health

Fire

Reactivity

0

0

1

Hydrogen Peroxide solution

Health

Fire

Reactivity

1

0

1

Ferric Chloride solution

41

PROCEDURE The assembled apparatus is shown below.

Oxygen generator 500 mL Florence flask

100 mL graduated cylinder

Ferric chloride solution

Hydrogen peroxide solution

water

pinch clamp

1. Fill a florence flask with water and attach the tubing assembly so that the long piece of tubing in

the flask extends nearly to the bottom and the short piece is above the surface of the water. 2. Place the end of the tubing assembly without the rubber stopper into the 100 mL graduated

cylinder. 3. Fill the tubing that connects the Florence flask and the graduated cylinder with water. You can

do this by using the pipet bulb to apply pressure from the other end of the tubing assembly (the one with the stopper). This forces water from the Florence flask into the graduated cylinder. At the end of this step the water levels in the Florence flask and the graduated cylinder should be equal and there should be no air bubbles in the tubing that connects them.

4. Close the tubing between the Florence flask and the cylinder with the pinch clamp and then

discard the water in the cylinder. 5. Obtain about 20 mL of hydrogen peroxide solution in a beaker (enough for all three runs). Use

a 10 mL graduated cylinder to measure 5 mL of hydrogen peroxide solution and pour it into the oxygen generator. Make sure that none of this solution gets into the inner tube.

6. Put 0.5 mL of the 1.0 M ferric chloride solution into the inner tube of the oxygen generator.

(The plastic pipets provided with the solutions have graduated stems. The 0.5 mL mark is about halfway up the stem. Make sure the pipet goes back to the right bottle.)

7. Connect the oxygen generator to the Florence flask by attaching the stoppered end of the tubing

assembly.

42

8. Open the pinch clamp. A few drops of water should flow into the graduated cylinder. A steady stream of water or continual dripping indicates an air leak. An air leak can be eliminated by making sure all of the stoppers are tight and trying the procedure again. When you are sure there are no air leaks, shake the oxygen generator flask gently (but thoroughly), so that the two solutions mix. The pinch clamp should remain open during this operation. The gas evolved will force an equivalent volume of water into the graduated cylinder.

9. Continue to gently shake the oxygen generator. Record the volume of water collected in the

graduated cylinder at 15 second intervals until the volume stops increasing. Record the final volume of water in the graduated cylinder.

10. Repeat the experiment using 0.5 mL of the 0.5 M ferric chloride solution in the place of 1M

ferric chloride. 11. Repeat the experiment again using 0.5 mL of catalase solution in the place of 1M ferric chloride. RESULTS 1. Plot the results of each run as milliliters of water displaced vs. time and use this data to compare

the rate of reaction (at the start of the reaction) for different concentrations of Fe3+ catalyst and catalase.

CALCULATIONS 1. The volume of water displaced is a measure of how much oxygen gas was produced by the

reaction. Use the ideal gas law (PV=nRT) to determine how many moles of oxygen was produced. Assume the pressure to be 1.0 atmospheres and the temperature to be 293K (20oC):

n= PV/RT

2. Using the number of moles of oxygen produced and the balanced equation for the reaction

calculate the number of moles of hydrogen peroxide in a 5 mL sample. 3. Using the number of moles and the molecular weight of hydrogen peroxide calculate the number

of grams of hydrogen peroxide in a 5 mL sample:

g H2O2 = (moles H2O2) x (molecular weight H2O2) 4. The concentration of the peroxide solution is listed as 3% by the manufacturer (3g H2O2/100

mL solution). What was the % concentration of your sample?

43

QUESTIONS 1. What is the effect of the Fe3+ concentration on the initial rate? on the final volume of O2? 2. Even though no catalyst is "used up", why should the rate depend on catalyst concentration? 3. What concentration of ferric chloride would be equivalent to the catalase solution in efficiency?

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DETERMINATION OF pH USING VEGETABLE INDICATORS

An acid may be defined as a substance which is capable of donating a proton; conversely, a base is a substance which accepts protons. An example of this is shown below:

HCl + H2O H3O+ + Cl-

In this example, hydrochloric acid, HCl, is the acid and donates a proton (H+) to the base, H2O, water. Hydrochloric acid is a strong acid which means that it is a good proton donor. When hydrochloric acid is mixed with water the above reaction essentially goes to completion with all the HCl being converted to H3O+, hydronium ion, and Cl-, chloride ion. Acetic acid (which is found in vinegar) is a weak acid:

CH3COOH + H2O ⁄ CH3COO- + H3O+

A weak acid is not a good proton donor so when the same volume of acetic acid is mixed with water, few hydronium ions are formed. If equal molar amounts of acetic acid and hydrochloric acid are put into water, the solution of acetic acid has a low concentration of hydronium ions while the solution of hydrochloric acid has a high concentration of hydronium ions. To measure the strength of an acid we measure the concentration of H3O+. The scale used to denote acid strength is pH. The pH of a solution is equal to -log[H3O+]. Since this is a negative log, high concentration of H3O+ corresponds to a low pH. Thus a 0.1 molar (0.1 M) solution of HCl, a strong acid, has a pH of -log[1x10-1] = 1. A 0.1 M solution of acetic acid, a weak acid, has a pH of -log[1x10-3] = 3 because of the 0.1 moles of acetic acid present in each liter of solution, only 0.001 moles or 1/100 of the total amount is dissociated at any point in time. In pure water, there is a slight dissociation of the water molecules into hydronium ions and hydroxide ions:

2H2O ⁄ H3O+ + OH-

In pure water the concentration of hydronium ions is 1x10-7 moles per liter and because every time a H3O+ is formed an OH- is formed, the concentration of OH- is also 1x10-7 moles per liter. Pure water has a pH of -log[1x10-7] = 7. Since the concentration of OH- = H3O+, pure water is neither acidic nor basic but is considered neutral. Most familiar bases are ionic compounds which contain the hydroxide ion, such as NaOH or KOH. These bases accept protons from acids to form water:

H3O+ + OH- ⁄ 2H2O

Again, a strong base such as NaOH will completely dissociate in water to give a high concentration of hydroxide ions. Some bases do not themselves contain hydroxide ions but form hydroxide ions when in aqueous solutions through the dissociation of water. Ammonia is an example of such a base:

NH3 + H20 ⁄ NH4+ + OH-

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The pH system is also used to measure the strength of a base. How can this be done when pH is determined by the concentration hydronium ions and a base usually involves hydroxide ions? Remember that the concentrations of hydronium ion and hydroxide ion in pure water is equal to 1x10-7 moles per liter. The product of the concentration of hydronium ion and hydroxide ion is called the ion product of water.

[H3O+][OH-] = [1x10-7M][1x10-7M] = 1x10-14M

The ion product of water is a constant (at constant temperature) and is the same for all aqueous solutions. Since this number is constant, when the concentration of hydronium ion increases, the concentration of hydroxide ion must decrease and vice versa. When a base is present the concentration of hydroxide ion increases and the concentration of hydronium ion decreases. Since the pH is -log[H3O+], when the concentration of hydronium ion decreases, the pH increases. Just as there are strong and weak acids, there are also strong and weak bases. In a 0.1 molar solution of NaOH, a strong base which is completely dissociated, [OH-] = 0.1 = 10-1. Since the product of [OH-] and [H3O+] must be 10-14 for all aqueous solutions, [H3O+] must equal 10-14/10-1 or 10-13. The pH of a 0.1 molar solution of NaOH thus is -log[10-13] = 13. In a 1 molar solution of NH4OH, a weak base, only about 1/100 of the ammonium hydroxide is dissociated at any point in time so [OH-] = 0.1/100 = 10-3. The pH of a 1 molar NH4OH solution is -log[10-14/10-3] = -log[10-11] = 11. How does one determine the pH of a solution? One of the earliest methods for determining the pH of a solution was to use chemical compounds derived from plants which change color with the pH of the solution. These substances are called indicators. One such indicator, litmus, will turn blue in base and red in acid. Some indicators have a wide range of color changes. In this experiment, you will extract colored substances from cabbage and use them as an indicator. First, you will determine the color the indicator will be at a specific pH using the buffer solutions of known pH provided for you in the lab. Then you will take various household substances and common laboratory solutions and determine the pH of their aqueous solutions. MATERIALS supplied in lab: cabbage leaves, ringstand, wire gauze pad, Bunsen burner, buffer solutions, pH paper, household substances from your desk: 250 mL beaker, 100 mL graduated cylinder, test tubes and rack, eyedropper SPECIAL NOTES/SAFETY CONSIDERATIONS Be careful not to spill the solutions on yourself or your clothing. In case of spills, wash your skin thoroughly with water and clean up the lab bench. Use caution with the household substances - familiarity tends to breed carelessness. Metro asks that any solutions that go down the drain be between pH 5.5 and 12, so upon completing this lab combine all your solutions into one beaker and adjust the pH accordingly.

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0

3 2

Health

Fire

Reactivity

0.1 M Ammonium Hydroxide

Avoid contact

2

2 2

Health

Fire

Reactivity

0.1 M Acetic Acid

Corrosive

0

2 0

Health

Fire

Reactivity

Buffers: pH 2 & 3

0

0 0

Health

Fire

Reactivity

Buffers: pH 4, 5 6, 7 &8

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0

2 0

Health

Fire

Reactivity

Buffers: pH 9,10 & 11

0

3 2

Health

Fire

Reactivity

0.1 M HCl

Corrosive!

PROCEDURE 1. Fold a leaf of red cabbage into a 250 mL beaker. Add approximately 50 mL of distilled water. 2. Attach an iron ring to a ringstand and place a piece of wire gauze on the ring. Place the beaker of

cabbage and water on the gauze pad. 3. Place a Bunsen burner under the set-up and boil the cabbage in water gently until the red-purple

color of the cabbage is extracted into the water. This step usually takes about 15 minutes. Add more water if it starts to dry out.

4. Turn off the burner and allow the mixture to cool for at least ten minutes. 5. Set up an array of buffer solutions of the different pH's provided. Label 10 small test tubes to

correspond with pH's of the buffer solutions (pH's 2-11). Fill each tube approximately 1/2 full (~5mL) with the appropriate buffer from the stock solution squeeze bottle.

6. Arrange the tubes in increasing pH values in your test tube rack. Add several drops of the

cabbage indicator to each tube with an eye dropper. You should get a nice array of colors from the vegetable dye. Note these colors in your lab notebook. It's okay to add more of the cabbage indicator to intensify the color, but be sure to add the same amount to each test tube.

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7. Test the pH of the household products and laboratory reagents available in the lab by placing a 5mL sample in a test tube and adding cabbage indicator. Solid samples such as antacid tablets should be prepared by crushing about 1/4 of the tablet and dissolving it in ~ 5mL of deionized water. Compare the resulting colors to your buffer array to approximate the pH.

QUESTIONS 1. List the household products in order of increasing pH. Does this order indicate increasing acidity

or increasing basicity? 2. Calculate the concentration of H3O+ and OH- in each sample.

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