Manual Chemical Practices

Manual Chemical practices . T.S.U. level by Isaac Misael Lucas

Gmez is licensed under a Creative Commons Reconocimiento-NoComercial 4.0 Internacional License.

LABORATORY MANUAL FOR INTRODUCTION OF PHYSICS - CHEMISTRY

T.S.U.

Collector autor: Lic. ISAAC MISAEL LUCAS GMEZ

LABORATORY MANUAL - CHEMISTRY T.S.U.

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Manual Chemical practices . T.S.U. level by Isaac Misael Lucas Gmez is licensed under a Creative

Commons Reconocimiento-NoComercial 4.0 Internacional License.

Introduction

Chemistry is a discipline based on observation (as are all sciences). In lecture, you will learn the principles and theories that, to date, best explain the observations that have accumulated. The problem is that, if all you have is lecture, and then it is all too easy to forget that these theories apply to the real world. The laboratory experience is, by design, your opportunity to see these principles and theories in practice.
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Table of Contents

Title Experiment Page Number

Laboratory Safety 4

Experiment 1: Density of a Liquid and a Solid 7

Experiment 2: Study of changes of liquid with temperature. 11

Experiment 3: Gas Laws 14

Experiment 4: Colligative properties 23

Experiment 5: Classification of Materials 30 Experiment 6: Materials Science Applied to Household

Appliances. 33


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LABORATORY SAFETY

A.Come to the Laboratory Prepared for Safety: 1. At the beginning of the course, you were given the manual of laboratory experiments. Read the experiment BEFORE you come to class. Make sure you fully understand the experiment before starting the actual work. If you have a question, ask your professor for clarification BEFORE starting the procedure. 2. Do only the experiments that have been assigned by your professor. No unauthorized experiments will be allowed. 3. SAFETY GOGGLES MUST BE WORN AT ALL TIMES IN THE LABORATORY. 4. It is not advisable to wear contact lenses during lab. 5. Do not wear loose clothing to lab. It is a fire hazard. Wear closed shoes. Tie back long hair so it does not fall into chemicals or into a flame from a Bunsen burner. 6. Learn the location and use of the emergency eye-wash fountain, emergency shower, and fire extinguishers. Memorize their locations in the laboratory. Know the location of the exits in the lab.


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7. NO FOOD OR DRINK IS ALLOWED IN THE LABORATORY. Never put anything into your mouth while you are in the laboratory. Wash your hands before leaving lab. 8. Behave in a responsible manner while in lab. Be aware of the other students around you. 9. Keep the lab bench clear of all personal items not needed for the experimental work. Store backpacks, purses, and coats in the storage area provided. B. Handle Chemicals and Equipment in a Safe Manner 1. Double check the label on the container before you remove a chemical. To avoid contamination of the chemical reagents, NEVER insert droppers, pipets or spatulas into the reagent bottles. 2. Take only the quantity of chemical needed for the experiment. Pour or transfer a chemical into a small, clean container from your place. Label the container. Do not take the stock container to your desk. 3. DO NOT RETURN UNUSED CHEMICALS TO THE ORIGINAL STOCK CONTAINERS. You risk contamination of the chemicals. Follow your professors instructions for disposal of unused chemicals. 4. Do not shake laboratory thermometers. Laboratory thermometers respond quickly to the temperature of their environment. Shaking a thermometer is unnecessary and can cause breakage. 5. Clean up spills. Spills of chemicals or water in the work area or on the floor should be cleaned up immediately. Small spills of liquid can be cleaned up with a paper towel. Use Sodium Bicarbonate to neutralize any acid spills. Mercury spills require special attention. Notify your professor if you break a thermometer so that special methods can be used to remove the mercury. 6. Dispose of broken glass in the special containers provided. Do not put broken glass in the wastepaper basket.


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7. Heat only heat-resistant glassware (marked Pyrex or Kimax). Other glassware may shatter when heated. Be very careful of hot objects. Iron or glass looks the same when it is hot as it does at room temperature. 8. Be careful of fires. Small fires can be extinguished by covering them with a watch glass. If a larger fire is involved, a fire extinguisher can be used. If clothing or hair catches on fire the rule is drop and roll to extinguish the flames. 9. Report any injuries that occur in the laboratory to your professor.

LABORATORY SAFETY RULES

Note: Failure to follow safety rules will result in expulsion from this course.

1. Wear approved safety goggles AT ALL TIMES in the laboratory. 2. It is not advisable to wear contact lenses during lab. 3. Do not wear loose clothing to lab. It is a fire hazard. 4. Tie back long hair. It too is a fire hazard. 5. Wear closed shoes to lab. 6. Never put anything into your mouth while in the lab. 7. Immediately wash off any chemicals spilled on your skin or clothes. 8. Keep the lab neat. Return reagent containers and equipment to proper locations. Put any belongings not needed for experimental work on the shelves provided. 9. Clean up all chemical spills or broken glass immediately. You should report Mercury spills to your instructor or the stockroom (do NOT attempt to clean up spilt Mercury). 10. Think about how much chemical you will need before you take it from a stock (reagent) bottle. NEVER return unused chemicals to stock bottles. 11. Dispose of waste chemicals only as instructed. 12. Behave in a responsible manner. 13. Be aware of the location and use of laboratory safety equipment. 14. Immediately report accidents and injuries to your professor.


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15. Do NOT perform unauthorized experiments 16. Thoroughly wash your hands any time you leave the lab. 17. No smoking in or near the Allied Health and Sciences building. I have carefully read all of the safety precautions summarized above and recognize that it is my responsibility to observe them throughout this course.

Student Full Name:

ID number

Signature

Introduction

Experiment 1: Density of a Liquid and a Solid

Number of sessions: 1

Goal Know of meaning of density and how determinate. Purpose (1) To practice the procedures commonly used in a laboratory for take liquids. (2) To learn how to use some of the common laboratory devices. (3) To distinguish between chemical and physical properties. Skills: analytic capacity, team work, responsibility, reasoning.


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In just about any lab manual you look at, you will notice that the first experiment is something like the determination of the density of various objects. The reason for this is quite simple to understand: the author wants you to learn how to use some of the basic pieces of the lab while performing an experiment that is relatively safe. In time, you will be performing experiments that do, out of necessity, have inherent dangers, but before you do, you want to be comfortable with your own laboratory skills through practice. That is really what this experiment is all about. You will be performing a series of relatively simple procedures, but as you do, keep in mind that these are skills and tools you will need in future experiments, so be sure to get any questions that arise answered, and be sure to take many notes and observations for yourself for future reference, especially potential problems and thing to watch out for when there techniques show up again. Remember to refer back to laboratory procedures for any techniques you do not know.

Background: See Basic Laboratory Procedures: pipette, graduated cylinder, analytic balance. MATERIAL.

Quantity Material Characteristics

200ml water

200ml Ethanol

2 Precipitant glass Capacity 100ml

1 Erlenmeyer flask Capacity 250ml

1 Pipette Capacity 20ml

1 Beaker Capacity 500ml

1 Graduated cylinder 20ml



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Procedure: Part I: Density of a Liquid You will measure the density of the same liquid three times to demonstrate the difference in various techniques for measuring the volume. In a clean, dry beaker of an appropriate size, get approximately 30 ml of the unknown liquid and bring it to your desk. Get the dry weight of a second clean and dry container. Using a pipette, put 20.00 ml of the unknown liquid into the second beaker. Determine the mass of the second beaker with the liquid in it. Return the liquid to the first beaker, and dry the second beaker. Using a clean and dry graduated cylinder, measure out 20.0 ml of the unknown liquid and put it into the second beaker. Again, get the mass of the second beaker with the liquid in it.Once again, return the unknown liquid to the original beaker. Measure out 20.00 ml of the unknown liquid using an Erlenmeyer flask. Determine the mass of the beaker with the liquid in it. Dispose of the unknown liquid as instructed. Part II: Density of an Unknown Solid From time to time, a chemist has to be clever enough to find an indirect method to measure some quantity. For example, how would one go about measuring the volume of an unusually shaped solid. Archimedes faced this problem when be had to determine the density of a crown for the king in order to determine whether or not the blacksmith stole some of gold and substituted copper for it. To do so, he used water displacement to determine the volume of the crown, as you will do for this part of the experiment. Get a solid object from your instructor. Determine At mass on an electronic balance. Choose a graduated cylinder of an appropriate size. Fill it approximately half full with water. It is not important to fill it to exactly half, but it is important to determine exactly what the initial volume is. Once you have recorded the volume, carefully lower the solid into the graduated cylinder so as to avoid


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splashing the water or breaking the graduated cylinder. Once you have recorded the final volume, dry the solid and return it. Calculations: Part I: For each of the three trials, determine the mass of the liquid by subtracting the mass of the container from the mass of the container and liquid. Divide the mass of the liquid by the volume (20 mL) to determine the density. Part II: To determine the volume of the object, subtract the volume of the liquid in the graduated cylinder from the volume of the liquid in the graduated cylinder with the object. To get density, divide the mass of the object by its volume. Questions: 1. Why are we measuring the density of the same liquid using three different techniques? 2. How do we determine the volume of an oddly shaped solid? 3. What volume of liquid are we using to determine the density of the oddly shaped solid? 4. Who developed the method of volume by water displacement? 5. What is the quickest method for determining the volume of a liquid? Which is the most accurate? 6. How do your densities compare with the three methods of volume determination from part I for the liquid? 7. How would you measure the volume of a sample of sand?


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8. For each of the following, what technique would you choose for measuring the volume? (a) You want to take 50 mL of a reagent from the area that it is stored to your desk (b) A titration requires 10.00 mL of a reagent measured as accurately as possible (c) A synthesis requires 35 mL of acid

REFERENCES:

Goal: knowledge of the influence of temperature in the physical characteristics of the substances Purpose (1) To practice the procedures commonly used in a laboratory for take liquids. (2) To learn how to use some the thermometer and analytical balance. (3) To learn the influence of temperature in physical phenomenon Skills: analytic capacity, team work, responsibility, reasoning.

Experiment 2: Study of changes of liquid with temperature. Number of sessions: 1


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Introduction. Although all matter has mass and volume, the same mass of different substances have different volumes occupy and noticed that iron or concrete are heavy, while the same amount of rubber or plastic are lightweight. The property that allows us to measure the lightness or heaviness of a substance called density. The higher the density of a body will seem heavier. Density is defined as the ratio between the mass of a body and the volume it occupies. Thus, as in the S.I. mass is measured in kilograms (kg) and volume in cubic meters (m3) density is measured in kilograms per cubic meter (kg / m3). This unit of measurement, however, is little used because it is too small. Most substances have similar water so, using this unit, it would always using very large numbers densities. To avoid this, is often used another measurement unit gram per cubic centimeter (g. / Cc). Background: Investigate the definition and mathematical expression of density. Investigate the definition of specific weight. Describe the difference between specific gravity and density. Material:

Quantity Material Characteristic

1 Pycnometer


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1 bath Maria

1 Heating grill

1 thermometer

1 universal support

1 Caliper three fingers

Cooking Oil

Procedure:

Initially the pycnometer thoroughly washed using distilled water, then let dry and weigh the pycnometer using an analytical balance with greater accuracy. Fill the pycnometer to the mark with vegetable oil, prevent the formation of air bubbles in the pycnometer; then wipe off excess oil from abroad and precede despite the analytical balance the pycnometer with oil. Repeat the process though, but applying heat to reach the following temperatures: 30, 35, 40, 45, 50,55 and 60 C. In each case should reach the desired temperature and immediately proceed to take the weight of the pycnometer in each case.

Questions: 1. Say that oil density determined for each of the temperatures? 2. Construct a graph of the results of density and temperature. 3. What conclusions obtained from the graph of density vs temperature? 4. How is the density of the liquid to the density of gases?


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5. How does the density of a solid, is modified when it is higher or lower volume but does not change the mass?

REFERENCES

Manual

Chemical practices . T.S.U. level by Isaac Misael Lucas Gmez is licensed under a Creative Commons Reconocimiento-NoComercial 4.0 Internacional

License.

Introduction When one thinks of the gas laws in terms of the medical field, respiration immediately springs to mind. After all, respiration allows us to exchange oxygen for use in the cells with carbon dioxide, the bi-product from the cells, and since we are not aquatic animals, our respiration takes place in a gaseous medium. Since respiration takes place in a gaseous medium, it is subject to gas laws. One of the most significant laws governing respiration is Boyles Law, but lets not forget Charles law, without which temperature calculations would be impossible. Lets begin with a review of volume. Volume we all know. A one dimensional structure such as a line has length, a two dimensional structure such as our shadows have area and as three dimensional creatures, we have volume. Volume is an effect of the third dimension, which is where most of us exist, although there may be those who seem to be from dimensions other than our own. Anyway, lets start out with a length, say 1 cm. If we attach another length to the end of the first

Goal: To become familiar with the behavior of gases Purpose (1) To practice the procedures commonly used in a laboratory. (2) To learn Boyles Law (3) To learn Charles law Skills: analytic capacity, team work, responsibility, reasoning.

Experiment 3: Gas Laws Number of sessions: 1


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one, also 1 cm, we have defined a box, which has area. The area is the two sides multiplied together, or, 1 cm2. To add a third dimension, we can place a 1 cm line connected at the same point where the first two lines are connected and at a right angle to each of them. The only way this is possible is to place it perpendicular to the plane defined by the first two lines, thereby defining a box. The volume of the box will be the length of all three sides multiplied together, or 1 cm3. This volume is one cubic centimeter, or 1 cc. It is also, by definition, 1 mL. Thus, 1 cm3 = 1 cc = 1 mL. The mL is the connection to liters, L, which is the metric unit used for volume. Pressure, on the other hand, is a force per unit area. Force is mass (a measure of the quantity of matter) times acceleration, F=ma. The most familiar force in USA country is the pound, lb. Weight is a force. The acceleration is acceleration due to gravity, which is smaller on the moon. The amount of matter we would have on the moon, m, would be constant and thus, our mass would be constant, but because the acceleration constant due to gravity, a, would be smaller on the moon than it is on the earth, our weight would be less on the moon than here on earth. So, if we take a weight and divide that weight by an area, we have force. For instance, suppose there is a person that weighs 180 lb with feet that have a surface area (in shoes) of about 99 in2. Then the pressure that person is exerting on the earth as he or she is standing still is 180 lb/99 in2 = 1.82 lb/in2 = 1.82 psi. (psi stands for pounds per square inch.) Boyles law relates gas volume to pressure. Boyle carefully measured how the volume of a gas changes as he varied the pressure on that gas. He discovered that for a system with a fixed amount of gas (n) and temperature (T), a plot of volume versus pressure gave a straight line with a negative slope. A more common way to state this would be to say that as pressur e increases, volume decreases. We say that for constant n and T, volume is inversely proportional to pressuere. VP= k | n,T

The vertical line above with the subscript n,T is a mathematical symbolism used to remind us that this is true only if the number of moles of the gas n and the temperature T remain constant. A proportionality can be converted into an equation with the introduction of some constant, in our case k, even if we do not know the value of this constant. Thus, or, since k must be constant, for any two


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states 1 and 2 (that is, for any initial state 1 to any final state 2 where we vary either pressure or volume. V = k/P| n,T V1 P2 =V1 P2 Charles decided that he was rather more interested in the relationship between the volume of a gas and the temperature of the gas at constant number of moles n and pressure p. Holding the pressure constant on the cylinder (as simple as not adding any weight to the cylinder, making its pressure equal to atmospheric pressure), he measured the volume of a gas as he heated and cooled the cylinder. Charles discovered on plotting volume and temperature that there was a direct proportionality, that is, as temperature increased, so increased volume. He wrote the corresponding proportionality as the next step, as with Boyle, is to remove the proportionality by adding a constant:

V = rT

where r is some constant, or : V1/T1=r Since r is constant for any two states 1 and 2, we can write V1/T1=V2/T2

Unfortunately, there IS a problem with Charles. See, if you have one of the temperatures set at zero, then the equation becomes undefined. How does one circumvent this problem? Well, Charles did so by extrapolating his data all the way to V=0. Such a condition, where the volume occupied by the gas is zero, is only possible for an ideal gas, and since there are no ideal gases, this becomes a hypothetical limit. It is not possible, and yet, in this impossible situation, Charles noticed something wonderful; no matter what gas, or mixture of gases he used, no matter what pressure he kept the gas, no matter how many moles of gas he started with, these lines all extrapolated to exactly the same temperature; -273.15 C. No matter what he did, he could NEVER reach a temperature in his extrapolations below this temperature; that means that this must be the theoretical limit of the temperature scale. Recall that temperature is directly related to kinetic energy, or motion. Doesnt it make sense, then, that there must be a point where there is no more motion, where the temperature is so cold


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that all motion in a molecule actually stops? And once this state is reached, is it possible to have a lower temperature, since temperature is related to motion of molecules? No, of course not, because we can never have less motion than absolutely no motion at all. This temperature is called absolute zero; it is the coldest temperature theoretically possible, and it corresponds to a state where there is no molecular motion at all; no movement, no vibration, nothing. We can use this fact to get around the undefined equation problem. If we take the centigrade temperature scale, and add to it 273.15, then we get a new temperature scale where the temperature can never go below zero. In fact, since this was a theoretical limit only, we can never reach absolute zero either, so we will always have a positive number for temperature. We call this temperature Kelvin (the corresponding absolute temperature based on the Fahrenheit scale is called Rankine). Whenever working with temperature in the gas laws, you must always convert to Kelvin Background: Investigate the Boyles Law Investigate the Charles law Investigate the absolute scale of temperature. Material:


4 syringe 10ml

1 bath Maria

1 Heating grill

1 thermometer

1 universal support

1 Calipers of tree fingers

wather

Procedure:


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This experiment is broken down into two parts, one to demonstrate Boyles Law, and to demonstrate Charles law. It does not matter which you choose to do first. Wear your eye protection and aprons at all times. Follow all safety guidelines strictly. Boyles Law: Obtain a modified Leur Lock portion of a syringe, its plunger and a ruler. If the syringe has a plunger already partially in it, return it for a syringe which has no plunger in it! The syringe has been modified in two ways; first, no needle is available. Second, the end of the syringe has been sealed off to prevent air from escaping out of or going into the barrel once the plunger is in place. 2. If necessary, place the plunger back into the barrel. Test the syringe for air leaks by pressing the plunger with your thumb as far into the syringe as possible. Carefully listen for any hissing sounds. If you hear any hissing sounds, or if the plunger does not return to the top of the syringe on release at any point in the experiment, the syringe has an air leak and will not work. Report this to your lab supervisor and get a new syringe. 3. Remove the plunger from the syringe and carefully measure the inner diameter of the syringe with a ruler. Record your results in centimeters on the report sheet. Replace the plunger. 4. Using an adjustable clamp, fasten the syringe vertically to a ring stand. The syringe must be attached with the plunger on the top such that a mass can be balanced on top of the plunger. Make sure the syringe is as vertical as possible. 5. Obtain a book with a known mass from the instructor. Record the mass of this object in grams on the report sheet. 6. Carefully balance the object on the flat surface on top of the plunger. You want this object to be balanced such that it is not touching anything and can sit on the syringe without falling. This will require patience. 7. Once the object is balanced, push down on the object slightly to force the plunger down. Release and allow the plunger to rise back up. When the plunger has stopped moving, allow it to sit undisturbed for about 10 seconds. 8. Carefully read the volume of the syringe from the scale on the syringe barrel to the


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nearest mL. Be careful not to let the book fall on you if it should fall off of the syringe! Record this volume in mL on the report sheet. 9. Remove the object from the plunger and allow it to return to the top. If any sign of a leak ever appears throughout this experiment, report it immediately to the lab supervisor. 10. Repeat steps 4 through 8 for at least 5 different masses.

Charles Law: Obtain a sealed syringe with a plunger already in place and a thermometer. Never remove the plunger from this syringe! If your does not have a plunger in it, return it for one that has the plunger already in it and positioned in the middle of the volume range. 1. Read the thermometer and record the temperature in degrees celcius. Caution! Youre your thermometer contains silver colored mercury, remember that mercury is toxic. Be VERY cautious when handling these thermometers to avoid breaking them. Should one break, do NOT attempt to clean the broken thermometer up yourself; report it to your lab supervisor IMMEDIATELY! Record this temperature in the Charles law table. Read the volume on the syringe and record it next to the room temperature. 2. Prepare an ice bath by filling an appropriate sized beaker about 2/3 full of ice and adding water almost to the top of the ice. As this is chilling, begin


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heating a second beaker to boiling. Always be cautious when using an open flame or around hot water! 3. Place the plunger in the ice bath for several minutes, along with the thermometer. When the temperature on the thermometer stabilizes, record the temperature in the table. Read and record the corresponding volume from the syringe. 4. Once boiling, place the thermometer and syringe in the boiling water. CAUTION! Be careful NOT to allow the syringe to get too close to the edge of the beaker, or it will melt! Allow the thermometer and syringe to remain in the boiling water for several minutes. Once the temperature has stabilized, read the temperature and corresponding volume. Record these values in the table. 5. IF your lab supervisor has an additional temperature bath, ask him/her to put your syringe into it (CAUTION! This will be TOO cold for you to do without DIRECT supervision; if she/he is NOT standing nearby, wait until he/she returns before proceeding!). Ask the temperature, and record it while you wait for the syringe to cool. Read the volume of the syringe when you are instructed to do so. Calculations: Boyles Law: Calculate the radius of the plunger from its diameter (radius=diameter/2). For simplicity of calculation, we assume the plunger is flat (even though we know it is not). Therefore, we can calculate the area of the plunger as A= r2, where r is the radius and is about 3.141592654. To calculate the pressure exerted by the mass, simply divide the weight of the object by the area of the plunger. The total pressure on the gas trapped in the syringe is the pressure exerted by the mass (see above) plus the pressure exerted by the atmosphere, which we will assume to be 1 atm. Thus, for total pressure, put in the pressure exerted by the mass plus one. On a clean piece of graph paper, carefully plot the volume of the gas (y axis) versus the total pressure (x axis). Draw the best straight line that you can through the experimental points. Take the inverse of each pressure and fill in the column 1/Pressure. On a clean piece of graph paper, carefully plot the volume of the gas (y axis)


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versus 1/Pressure (x axis). Draw the best straight line that you can through the experimental points. Calculate the slope of this line. This slope is the value of the constant k(=nRT). Charles Law: Create a graph to plot volume versus temperature, but be SURE that the temperature scale extends at least to -600 C, and that the volume scale extends to 0 mL. Carefully plot your data points. Extrapolate the line with a straight edge to V=0, and read the resulting temperature. This is your estimate for absolute zero.

Gas Laws Experiment: Boyle and Charles Report Sheet

Boyles Law: Syringe Inner Diameter: ___________cm Syringe Inner Radius: _____________cm Syringe Inner Area: _______________cm2

Object Mass (g)

Object Mass (kg)

Pressure Exerted by Object

Volume of Gas (mL)

Total Pressure (atm)

1/(Total Pressure) (atm-1)

Charles Law:

Temperature in C Volume in mL


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Questions: 1. What is the function of the objects we are resting on top of the syringes in the Boyles Law portion of this experiment? 2. What must we be careful of in the Charles law portion of this experiment where we are putting the syringe in the boiling water? 3. Does the plot of volume versus total pressure demonstrate the inverse proportionality of volume and pressure? If so, why? If not, what were you expecting to see? 4. From the plot of volume versus 1/pressure, calculate the slope. This is the value for the constant k. 5. Since we know from the ideal gas law that k=nRT, then n=k/RT. Assume that the gas inside the cylinder is at 19oC (approximately room temperature). Calculate the number of moles of gas in the cylinder, n (R=82.06 mL*atm/mol*K). HINT! Dont forget to convert your temperature to Kelvin! 4) What is your estimate for absolute zero from your plot of volume versus temperature? List as many sources of error as you can. 8) Explain, in simple terms, exactly what is happening at absolute zero (assuming it can be reached). Why is this a theoretical limit, rather than a real limit?

REFERENCES


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Introduction. If our body is 90% water, that means we are 10% other stuff. Water can be thought of as the solvent in our bodies (the solvent can be thought of as the carrier; that in which the solute is dissolved), while the rest of the stuff (proteins, lipids, DNA, nutrients, waste, and a plethora of other things) can be thought of as the solute (the active ingredients in a solution; what makes the solution of interest). Typically, the solvent is the compound present in greater amounts, but this is not always the case; it is more generally correct to think of the solute as the active ingredient, that is, the reason for us to pick up the solution in the first place, while the solvent is the carrier for the solute.

Goal: Lear about colligative properties Purpose (1) To gain experience with solubility and colligative properties

Skills: analytic capacity, team work, responsibility, reasoning.

Experiment 4: Colligative properties Number of sessions: 1


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It is not technically correct to speak of heterogeneous or homogeneous solutions. All solutions, by definition, must be homogeneous. To be a true solution, you must have an even distribution of solute throughout every part of the solution. If you have an uneven distribution, then you have a mixture, which is heterogeneous. In fact, the line probably ought to have been disgusting walking mixtures. Mixtures are cloudy in appearance caused by the diffraction of light off of the separate regions of the mixture (called the Tyndall effect). Even a mixture that looks homogenous (or advertised to be so as in the case of milk) is in fact not a solution at all if it is cloudy. Solubility refers to the amount of solute that can be dissolved in a solution at a given temperature (and pressure if the solute is a gas). Notice that it does not speak to how long it takes to dissolve, just the maximum amount. This means that things like stirring, which makes things dissolve faster, will not influence solubility, just how long it takes for the solvent to dissolve. The proof is trying to dissolve excess solute in a solution that has already reached its solubility limit (called a saturated solution). Temperature will influence solubility, as will pressure but only if the solvent is a gas. The strongest influence of solubility is the nature of the solute and the solvent. There is an old rule of thumb when discussing solubility; like dissolves like. Although there are exceptions to this rule, generally speaking it means that polar solutes will dissolve in polar solvents, and non-polar solutes will dissolve in non-polar solvents. This provides interesting insight into substances, with a quick and convenient experiment to test polarity. It also provides insight around the home. If you wanted to remove peanut butter from a container, for example, you know that water will not work. Well, since water will not work, it might be worth your while to try a non-polar solvent first, like cooking oil. Once the peanut butter is gone, the cooking oil can be easily removed by detergent. The presence of a solute in a solution will influence the properties of the solution. That a solute will dissolve in a solvent to any extent means that the interaction between the solvent and solute molecules (or ions) is more energetically favorable than the interactions between the molecules (or ions) of the solute alone. In other words, the presence of the solute will attract solvent to itself and hold onto it strongly. This results in stronger intermolecular interactions in a solution than you would have in the solvent alone. As such, certain properties will change. These changes (freezing point depression, boiling point elevation and vapor pressure depression) depend on the concentration of the solution, but not on the identity of


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the solute. That is, the same concentration of any solute will produce the exact same properties. We call these Colligative properties. Background: Investigate the definition of Colligative properties. Material:


12 test tubes

1 bath Maria

1 Heating grill

1 thermometer

1 universal support

1 Caliper three fingers

water

Sodium chloride

Calcium phosphate

sugar

Naphthalene

Ethylene glycol

Paraffin Oil

potassium sulfate

sodium sulfate

potassium chloride

Procedure: The experiment part made in three parts:

Part A: Solubility of Compounds in Polar and Non-Polar Solvents:


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Water is a polar solvent, hexane is non-polar. In six test tubes, put 1 mL f water into each. In six different test tubes, put 1 mL of hexane (CAUTION! Toxic, highly flammable) into each. Pair the test tubes, one with water and one with hexane, into four pairs. Add equal amounts of each of the following to each of the two test tubes within a pair; (1) Sodium chloride (an ionic compound); the tip of a spatula full (about 0.1 g) (2) Calcium phosphate (CAUTION! Toxic) (an ionic compound); the tip of a spatula full (about 0.1 g) (3) Table sugar (sucrose, a polar covalent compound); the tip of a spatula full (about 0.1 g) (4) Naphthalene (CAUTION! Toxic, flammable, toxic fumes) (a non-polar covalent compound); the tip of a spatula full (about 0.1 g) (5) Ethylene glycol (CAUTION! Toxic) (typical ingredient in antifreeze; a polar covalent liquid); about 3 drops (enough to see) (6) Paraffin Oil (CAUTION! Toxic, flammable) (a non-polar covalent liquid); about 3 drops (enough to see) Agitate each solution by flicking the test tube. Describe what happens in each. Part B: Effect of Temperature on Solubility: Using a clean beaker, begin heating about 50 mL of distilled H2O. Place each of the following into three medium sized test tubes: (1) 2 g potassium sulfate K2SO4 (CAUTION! Toxic) (2) 7 g sodium sulfate Na2SO4 (CAUTION! Toxic) (3) 5 g potassium chloride KCl (CAUTION! Toxic)


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When the water reaches a temperature between 30 and 40oC, add 10 mL of warm water to each of the test tubes. Agitate well, but not all of the solute will dissolve. Continue heating the water. Once you have taken your observations, place the test tubes in the heating water. Continue heating the water, carefully agitating the solutions periodically, until the water is about 100oC (about boiling). Record your observations. Part C: Colligative Properties: We are going to make three solutions of more or less equal concentration. To do so, make each of the following solutions in three different test tubes: (1) Put 20.2 g of potassium nitrate (0.2 mol) and 20 mL of water, and agitate to dissolve. (2) Put 38.4 g of sugar (0.2 mol) and 20 mL of water, and agitate to dissolve. (3) Put 11.14 mL of ethylene glycol (0.2 mL) and 20 mL of water, and agitate to dissolve. Heat each test tube carefully in turn until the liquid begins to boil, take the temperature of the solution with thermometer. Record the temperature when the solution first begins to boil. Calculate the value of i for each substance. Calculations: Solubility

Solute Hexane Water

Sodium Chloride

Calcium Phosphate

Sugar

Naphthalene

Ethylene Glycol


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Paraffin Oil

Observations: Temperature and solubility:

Solute Low Temperature High Temperature

Potassium Nitrate

Sugar

Ethylene Glycol

Observations: Colligative Properties: The equation for boiling point elevation is Tb=iKbm, where m is molality (moles of solute per kilogram of solvent), and for water, the boiling point elevation constant, Kb, is 0.512 C/m. Solute Boiling Point Value of i KNO3

sugar

ethylene glycol


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Questions: 1. What similarities do you notice in the salts used in the effect of temperature on solubility? What differences? 2. What Colligative property are we studying? Do you expect that cooking oil is polar or non-polar? Explain your reasoning. Look at the compounds we used for boiling point elevation and the value of i that you calculated for each. What do you suppose i stands for?

REFERENCES


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Introduction. There are many places to get samples of materials. You can use old appliances, or go to junk yards, flea markets, or various industries. Materials are all around us. Know what the samples you have selected are. Be sure to include fibers like Kevlar, glass wool, fiberglass, and composite materials. Mylar and reflective mylar can be used to give students something to think about in classifying them. Metals have identifying characteristics such as shine, hardness, ductility, and they conduct heat and electricity. Ceramics tend to be hard, but brittle, stiff, and do not conduct heat or electricity as a rule. Polymers are usually flexible, have a low density, are insulators,and burn.

Goal: The student can selected material(s) into one of three categories: Metal, ceramic, or polymer Purpose (1) Selected material(s) into one of three categories: Metal, ceramic, or polymer.


Experiment 5: Classification of Materials Number of sessions: 1


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Composites are combinations of any of the above materials. In some cases, no material by itself (metal, ceramic, or polymer) has the characteristics required for a particular use, so a combination of materials (composite) is used. Background: The students will begin to establish concepts of materials, their characteristics, and how this relates to function. Material:

An assortment of different materials taken from various sources in the environment. Examples include parts of appliances, fabrics, bottle fragments (both glass and plastic), nails, wires, fiberglass, and insulating materials. Be sure to include a few items that are composite materials so students will have to ponder where to place them. It is best to have at least one sample per student.

Procedure:

1. Display the materials on a table or desk in front of the classroom.

2. On the table or desk, set aside space for three areas labeled metals, polymers, and ceramics where students may place an object after they have identified the material. 3. Have students, one at a time, select an object of their choice and place it in the category they feel is appropriate. 4. After students have categorized all objects, select various samples and have the students who classified those objects justify why they were placed in certain categories. 5. Place a randomly selected material(s) into one of three categories: metal, ceramic, or polymer.


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Observations:

Questions: 1. What are characteristic of metals? 2. What are characteristic of polymers? 3. What are characteristic of ceramic?

REFERENCES


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Introduction. There are many places to get samples of materials. You can use old appliances, or go to junk yards, flea markets, or various industries. Materials are all around us. Know what the samples you have selected are. Be sure to include fibers like Kevlar, glass wool, fiberglass, and composite materials. Mylar and reflective mylar can be used to give students something to think about in classifying them. Metals have identifying characteristics such as shine, hardness, ductility, and they conduct heat and electricity. Ceramics tend to be hard, but brittle, stiff, and do not conduct heat or electricity as a rule. Polymers are usually flexible, have a low density, are insulators, and burn.

Goal: This is a project to give the students experience with materials that they are around every day. It provides an opportunity for them to explore, discover, and handle the inner makings of common household items. Purpose Dismantle a small appliance and organize and/or categorize materials from within the appliance into groups of materials, categories of physical properties, or types of materials used in engineered systems.


Experiment 6: Materials Science Applied to Household Appliances. Number of sessions: 1


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Composites are combinations of any of the above materials. In some cases, no material by itself (metal, ceramic, or polymer) has the characteristics required for a particular use, so a combination of materials (composite) is used. Background: Investigated principal characteristics of metal, ceramic and polymer. Material:

Old appliances such as toasters, irons, hair driers, wind-up toys, clocks, curling irons, cameras, mechanical or electrical toys. Screwdrivers, (Most screws can be undone with mini-screwdrivers, but you will want to have some large-handled screwdrivers to loosen hard turning screws.) Pliers Wire cutters Candle Containers (plastic or paper bags) Permanent marker pens Ohm meter or continuity device Safety glasses

Procedure:

Dismantle the appliance using the tools needed to remove the appliances casing and inner parts. 2. Place disassembled parts into containers labeled metals, ceramics, polymers, and composites. 3. Discuss the quantities of materials gathered in each container. Name some reasons certain materials are more commonly used than others. Could there be a better material to use than what is found in


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your appliance? Why do you think the manufacturer decided to use the material currently used in the appliance? 4. Record observations about the disassembled appliances in the laboratory record book. Is there a particular part or mechanism that could be drawn to show special details of this appliance? Additional Activities Electrical Conductivity Using a flashlight bulb, a battery, and three pieces of wire, measuring about 6 in. per wire, set up an electrical continuity device to check if electricity will conduct through some materials (see Figure 4.2). The light will light up if the material is electrically conductive and will remain off if no electricity passes through the material. (This a crude continuity device.) An ohm meter would also be useful to check the electrical resistance of materials. The experience of working with the ohm meter is valuable.

Be careful when working with the ohm meter so as not to destroy some of its internal parts. This would happen by touching the leads to a system that already has a voltage applied to it.


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Thermal Effects Obtain several liters of liquid nitrogen. Dip the dismantled materials in it and see if or how the cold temperature affects their strength. Polymers and composites will be most affected by the temperature; metals and ceramics will experience the least amount of change. Wearing leather gloves, flex the material being tested to observe if cold has changed the strength and flexibility of the material. Heating material will provide valuable information about many materials. Using a burning candle, Bunsen burner, or propane torch, pass each material slowly through the flame, and determine the effect of heat on the material. Most polymers can be identified by burning them and observing their smoke, smelling the fumes (carefully), and observing how it burns in the flame. Use caution because some materials will melt, drip, and splatter hot liquid. Other materials may oxidize and some materials may not be affected at all. Caution: Melt or burn unknown materials only in an area with direct exhaust to the outside. Some materials may burn and produce irritating, choking, and/or toxic fumes. Do not heat containers or electrical devices (i.e., capacitors) that may have a potential to explode. Observations: Record your observations in your journal. Details could include the following Drawings of appliances or specific parts Type of appliances worked on Types of material the major parts of the appliance are made from (What parts are made of metals, ceramics, polymers, composites?)


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Reasons why materials were chosen for specific purposes of the appliance Possible reasons the inventor or manufacturer used some unusual materials to make some parts of the appliance.

REFERENCES


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LABORATORY MANUAL FOR PRINCIPLES BASICS OF CHEMISTRY

T.S.U. ENVIRONMENTAL CHEMISTRY

Collector autor: Lic. ISAAC MISAEL LUCAS GMEZ


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Introduction

Chemistry is a discipline based on observation (as are all sciences). In lecture, you will learn the principles and theories that, to date, best explain the observations that have accumulated. The problem is that, if all you have is lecture, and then it is all too easy to forget that these theories apply to the real world. The laboratory experience is, by design, your opportunity to see these principles and theories in practice.


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Table of Contents

Title Experiment Page Number

Laboratory Safety 4

Experiment 1: Measurements: Length, volume and temperature. 7

Experiment 2: Compound types 13

Experiment 3: Determination of Chemical formulas 19

Experiment 4: Chemical Reactions 24

Experiment 5: Solutions 28

Experiment 6: Synthesis of a Compound 33


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LABORATORY SAFETY

A.Come to the Laboratory Prepared for Safety: 1. At the beginning of the course, you were given the manual of laboratory experiments. Read the experiment BEFORE you come to class. Make sure you fully understand the experiment before starting the actual work. If you have a question, ask your professor for clarification BEFORE starting the procedure. 2. Do only the experiments that have been assigned by your professor. No unauthorized experiments will be allowed. 3. SAFETY GOGGLES MUST BE WORN AT ALL TIMES IN THE LABORATORY. 4. It is not advisable to wear contact lenses during lab. 5. Do not wear loose clothing to lab. It is a fire hazard. Wear closed shoes. Tie back long hair so it does not fall into chemicals or into a flame from a Bunsen burner. 6. Learn the location and use of the emergency eye-wash fountain, emergency shower, and fire extinguishers. Memorize their locations in the laboratory. Know the location of the exits in the lab. 7. NO FOOD OR DRINK IS ALLOWED IN THE LABORATORY. Never put anything into your mouth while you are in the laboratory. Wash your hands before leaving lab. 8. Behave in a responsible manner while in lab. Be aware of the other students around you.


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9. Keep the lab bench clear of all personal items not needed for the experimental work. Store backpacks, purses, and coats in the storage area provided. B. Handle Chemicals and Equipment in a Safe Manner 1. Double check the label on the container before you remove a chemical. To avoid contamination of the chemical reagents, NEVER insert droppers, pipets or spatulas into the reagent bottles. 2. Take only the quantity of chemical needed for the experiment. Pour or transfer a chemical into a small, clean container from your place. Label the container. Do not take the stock container to your desk. 3. DO NOT RETURN UNUSED CHEMICALS TO THE ORIGINAL STOCK CONTAINERS. You risk contamination of the chemicals. Follow your professors instructions for disposal of unused chemicals. 4. Do not shake laboratory thermometers. Laboratory thermometers respond quickly to the temperature of their environment. Shaking a thermometer is unnecessary and can cause breakage. 5. Clean up spills. Spills of chemicals or water in the work area or on the floor should be cleaned up immediately. Small spills of liquid can be cleaned up with a paper towel. Use Sodium Bicarbonate to neutralize any acid spills. Mercury spills require special attention. Notify your professor if you break a thermometer so that special methods can be used to remove the mercury. 6. Dispose of broken glass in the special containers provided. Do not put broken glass in the wastepaper basket. 7. Heat only heat-resistant glassware (marked Pyrex or Kimax). Other glassware may shatter when heated. Be very careful of hot objects. Iron or glass looks the same when it is hot as it does at room temperature. 8. Be careful of fires. Small fires can be extinguished by covering them with a watch glass. If a larger fire is involved, a fire extinguisher can be used. If clothing or hair catches on fire the rule is drop and roll to extinguish the flames. 9. Report any injuries that occur in the laboratory to your professor.


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LABORATORY SAFETY RULES

Note: Failure to follow safety rules will result in expulsion from this course.

1. Wear approved safety goggles AT ALL TIMES in the laboratory. 2. It is not advisable to wear contact lenses during lab. 3. Do not wear loose clothing to lab. It is a fire hazard. 4. Tie back long hair. It too is a fire hazard. 5. Wear closed shoes to lab. 6. Never put anything into your mouth while in the lab. 7. Immediately wash off any chemicals spilled on your skin or clothes. 8. Keep the lab neat. Return reagent containers and equipment to proper locations. Put any belongings not needed for experimental work on the shelves provided. 9. Clean up all chemical spills or broken glass immediately. You should report Mercury spills to your instructor or the stockroom (do NOT attempt to clean up spilt Mercury). 10. Think about how much chemical you will need before you take it from a stock (reagent) bottle. NEVER return unused chemicals to stock bottles. 11. Dispose of waste chemicals only as instructed. 12. Behave in a responsible manner. 13. Be aware of the location and use of laboratory safety equipment. 14. Immediately report accidents and injuries to your professor. 15. Do NOT perform unauthorized experiments 16. Thoroughly wash your hands any time you leave the lab. 17. No smoking in or near the Allied Health and Sciences building. I have carefully read all of the safety precautions summarized above and recognize that it is my responsibility to observe them throughout this course.


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Student Full Name:

ID number

Signature

Introduction The world uses a variety of units to measure length. Scientists use the metric system in which the unit of length is the meter (m). Using appropriate prefixes, one can indicate a length that is greater or less than a meter. A meter stick can be divided into 100 centimeters. Each centimeter could be divided into 10 millimeters. When a meter stick or other measuring device is used, the measurement must be reported as precisely as possible. The number of significant figures that can be included depend on the markings on the device that is used. When a piece of data is

Experiment 1: Measurements: Length, volume and temperature. Number of sessions: 1

Goal Know the correct form to measurement in the laboratory Purpose (1) To practice the common procedures.



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recorded, the next to the last digit reported is the number represented by the smallest increment marked on the measuring instrument. The last digit is an estimate. If the quantity being measured is estimated to fall exactly on a line marked on the measuring device, the last digit of the measured number is a significant zero. If the quantity being measured is seen to fall between two lines, an estimate is made as to the distance the quantity is between the lines. This estimated number becomes the last digit recorded for the measurement. In this experimental procedure, you will use graduated cylinders to determine the volumes of several substances. To read the volume of liquid properly you must avoid parallax. You should set the cylinder on a level surface and bring your eyes to a level even with the top of the liquid. You will notice that the liquid level is not a straight line, but curves down at the center. This curve, called a meniscus, is read at its lowest point (center) to obtain the volume measurement of the liquid. In the graduated cylinder shown, the volume of the liquid can be read as 42.1 mL. (Note: the smallest markings on the cylinder shown are for 1 mL increments. By estimating the volume between the 1-mL markings, the volume can be reported to the tenths (0.1) of a milliliter.)

Temperature can be measured by several different methods. One method uses the fact that when most liquids are heated, their volume increases are almost directly proportional to the change in temperature. A mercury-in-glass thermometer, the type that will be used in this experiment, is constructed of a very small diameter capillary tube connected to a relatively large reservoir of mercury in a bulb. As the mercury is warmed it expands into the capillary tube. The change in height of the mercury column can be calibrated to correspond to a temperature scale. The scale most often used by chemists is the Celsius scale in which the normal freezing point temperature of water is 0C and the normal boiling point temperature of water is 100C.

Background: Research the metric system and English system of measurements. MATERIAL.


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Quantity Material Characteristics

200ml water

2 Precipitant glass Capacity 100ml

1 Erlenmeyer flask Capacity 250ml

1 Pipette Capacity 20ml

1 Beaker Capacity 500ml




Procedure: A. Length Measurements Use the metric scale on a ruler or meter stick to make the measurements indicated on the report sheet (be sure to record the number of significant figures appropriate for the measuring device(s) that you use). In some cases you will need to use a piece of string to determine the distance (around your wrist, for example). B. Measuring Volumes of Liquids 1. A display of graduated cylinders containing different volumes of liquids has been set up for you. Review the information in the Introduction concerning parallax, reading a meniscus, and reporting significant figures. Read and record the volume of the liquid in each graduated cylinder using the number of significant figures appropriate for that cylinder. 2. Using a measuring cup, measure out 1 cup of water as exactly as possible. Use a 250-mL or 500-mL graduated cylinder to measure the volume of the water in mL. Record your answer. Repeat this procedure two more times. Average your data to obtain an experimental value for the number of mL in a cup. Use English-Metric Unit conversion factors found in your textbook to calculate the true (actual) value for the number of mL in 1 cup. Comment on how close your experimental value is to the true value. C. Temperature Measurements


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The bulb of the thermometer must remain in the liquid while the measurement is being made. Beakers of water at different temperatures are available in the laboratory. Measure and record the temperature of each in C. Convert each temperature to Kelvin. D. English-Metric Conversions A selection of product containers is available in the lab. Pick four products and record the volume of the contents of each in both fluid ounces (English system), and Liters (metric system). Based on the volumes given, calculate an English/metric conversion factor in fluid ounces/L from each set of data. Calculations: A. Length measurements Complete the following equalities: 1 cm = ____________ m 1 km = ______________ m 1 mm = _____________ m Make each of the following measurements in centimeters, and then convert the measurements to millimeters and meters.

objects Measured number Calculated numbers (mm)

Calculated numbers (m)

Width of little finger nail

Width of desk top

Length of shoe

Your height

Length of a pencil

Determine the length and width of the sides of the rectangle shown above in centimeters. Have a second person repeat the measurements. Record both sets of data below. Calculate the area of the rectangle each of the sets of data.


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Measurement (first person)

Measurement (second person)

Measurement (Average)

Length, cm

Width, cm

Area, cm2

Would you expect the sets of data above to be identical? Explain why or why not. B. Measuring the Volume of Liquids

Volume in Display Cylinders, (mL)

Cylinder A Cylinder B Cylinder C

Volume of 1 cup of water, (mL)

Trial 1 Trial 1 Trial 1

Average Volume of 1 cup, (mL)

Using conversion factors found in your textbook, calculate the actual number of milliliters in exactly one cup. (One quart is exactly four cups.) Determine the percent difference between your experimental value and the accepted value you just calculated. Show your calculations here. C. Temperature Measurements

Temperature in C Temperature in Kelvin

Ice water

ambiental water

Warm water

D. English-Metric Conversions


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Name of Product Volume in English Units (fluid ounces)

Volume in Metric Units (Liters)

Conversion factor (fluid ounces/L)

1

2

3

4

What do you notice about the conversion factors obtained from each product? Questions: You may need to look at the English-metric conversion table in your textbook to complete some of these problems. SHOW YOUR WORK including all conversion factors. 1. Convert 3.85 10-4 kilometers to (a) micrometers (b) feet (c) centimeters 2. A piece of string is found to be 35.9 meters long. How long is the string in inches? 3. A section of lawn that is 25.5 feet by 75.0 feet needs fertilizer. The fertilizer is sold in 5.00 pound boxes and 1.00 pound of fertilizer is needed for 10.0 square yards of lawn. If each box costs $1.65, how much will it cost to fertilize the lawn? 4. If a gallon of gas costs $1.50, how many cents would a liter of gas cost? (Round your answer to the nearest whole number of cents). 5. Convert 75.4F to Celsius and then to Kelvin.


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REFERENCES:

Introduction. One important thing is categorization of compounds into covalent and ionic. In class, probably discuss these compounds in terms of electrons, wherein ionic compounds transfer electrons and covalent compounds share electrons. How did the early chemists classify compounds, though, when they did not know what electrons were? They used properties, such as solubilities, melting points, and conduction. Solubility helps us to classify compounds as polar or non-polar, because, as a general rule, polar solutes dissolve in polar solvents (like water),

Goal: To examine the difference between ionic and covalent compounds and understand how their properties give rise to this categorization Purpose (1) To learn how differences between ionic and covalent compound Skills: analytic capacity, team work, responsibility, reasoning.

Experiment 2: Compound types Number of sessions: 1


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while non-polar solutes dissolve in non-polar solvents (like oils). Conductivity means whether or not a compound will conduct electricity when it is dissolved in water. We call these electrolytes, which are just like regular electros, but with 1/3 fewer calories. An electrolyte will conduct electricity when dissolved in water, while a nonelectrolyte will not. Finally, ionic compounds tend to have higher melting points than covalent compounds. These are summarized as follows:

Non-polar Covalent

Polar Covalent Ionic

Solubility non-polar solvents

polar solvents

polar solvents (or generally not soluble)

Conductivities Non-electrolytes

Non-electrolytes

Electrolytes (even if apparently not soluble)

Melting points very low low high

Background: Investigate the definition .of non-polar covalent compounds, polar covalent compounds and ionic compounds. Material:


test tubes

multimeter

distilled water

scoopula

Bunsen burner

Procedure:


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Your teacher gives a series of solid compounds. Run the following tests to decide if each compound is ionic, polar covalent or non-polar covalent based on the above table. Begin by taking very careful observations of each compound, and run each of the following tests on each compound. Solubility: You will need two clean test tubes (one of which is dry) for each of the unknown solids. Make sure that these test tubes are cleaned very well, and rinsed very thoroughly with distilled water. Any contamination from tap water or other sources will seriously affect your conductivity experiment. Put about 1 mL of water (a polar solvent) into one series of test tubes, and 1 mL of the non-polar solvent (probably Hexane) in the other series of test tubes. Place just enough of each solid into one water and one non-polar solvent test tube. Agitate each test tube by flicking it several times while holding it such that it does not fly out of your hands. Note whether or not the solid dissolved completely, dissolved partially, or did not appear to dissolve at all. Record your observations. Partially dissolved means that it is apparent that there is not as much of the solid in the test tube remaining as you put in initially, but there is still some solid left. If the amount of solid did not seem to decrease, it is insoluble. To be truly dissolved, the solution must be clear (not necessarily colorless, but clear). If it appears cloudy, then there is still undissolved solid in the test tube deflecting the light (called the Tyndall effect). You may discard the solutions with the non-polar solvent according to the instructions provided in lab. Keep the water solutions for the next step. Conductivity: Whether the solid appears to have dissolved or not, perform this conductivity test on each water solution. You will find a multimeter set up with conductivity available. Rinse the probe tips off very well with distilled water (you need not dry them). Rinse of a clean watch glass with distilled water as well. Pour a little bit of


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the liquid from the test tube onto the watch glass; if solid remains in the test tube, you need not include the solid in the quantity you pour out. Place the probes into the water solution on the watch glass, and give the probe a few seconds to equilibrate. Once equilibration seems to have been reached, record the conductivity for that solution. Discard the solutions according to the instructions given in lab, and remember to clean the probes and watch glass and rinse them all off with distilled water very well. Repeat for each water solution. Melting Points: Here we are not interested in absolute melting point temperatures, but rather, relative melting points. Place a VERY SMALL AMOUNT of each solid (just enough to see it) onto a scoopula, close enough that they can be viewed and heated more or less at the same time, but far enough apart that you can easily remember which is which. Light a Bunsen burner, and, being very careful not to burn yourself, pass the scoopula through the flame several times. Note the order of melting (which melts first, second, etc) and whether or not any appear to burn rather than melt. Record your observations. Remember that the compounds that melt first have the lower melting points. You might not be able to get all of the compounds to melt; these are all very high melting point compounds. Calculations: Based on the results from the procedure, categorize each compound as non-polar covalent, polar covalent, or ionic. Solubility: In the following table, record whether each substance is soluble, partially soluble or insoluble in each of the solvents.

Compound Number Polar Solvent Non-Polar Solvent


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Observations: Conductivity:

Compound Number Conductivity

Observations: Melting Point: In the following table, record which melted first, second, third, etc. in order of melting

Compound Number Order of Melting Observations

Observations:


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Classification: In the following, record for each compound if you believe it to be non-polar covalent, polar covalent or ionic.

Compound Number Compound Type

Questions: 1. Describe briefly how we are determining the order of melting. 2. What test are we using to determine if a compound is an electrolyte or a non-electrolyte? 3. If you have a compound that is not soluble in hexane or water, and is an electrolyte, what kind of compound do you expect you will have? 4. Is it possible to have an ionic compound that is not water soluble?


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5. Is it possible for an electrolyte to be covalent? 6. Is it possible to have an insoluble electrolyte? What does this imply about the term insoluble?

REFERENCES


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Introduction A chemical formula can be interpreted on two levels. On an individual (microscopic) basis, a chemical formula indicates the number of atoms of each element present in one molecule or formula unit of a substance. The subscripts of the formula can represent the numbers of atoms of the various elements present in one unit of the substance. At a macroscopic level, the subscripts in the chemical formula represent the numbers of moles of atoms of the different elements present in one mole of the substance. In an empirical formula, the subscripts represent the smallest whole number ratio of the atoms present in a substance. For most ionic compounds, their chemical formula is an empirical formula.

Goal: To become familiar with empirical formula Purpose (1) To practice the common reactions used in a laboratory. (2) To learn how determinate the basic empirical formula Skills: analytic capacity, team work, responsibility, reasoning.

Experiment 3: Determination of Chemical formulas Number of sessions: 1


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Background: Investigate types of reactions Investigate definition of empirical formula Material:


2 glass dish

1 Bunsen burner

2 beaker 250ml

1 hot plate

1 glass dropper

1 dessicator

1 g Manganese metal

drops Hydrochloric acid

Procedure: PROCEDURE 1. Clean an evaporating dish. Place the evaporating dish on a wire screen suspended on a ring approximately 5 cm above a Bunsen burner flame. Heat the dish for at least 5 minutes. 2. Place the evaporating dish in a dessicator and allow the evaporating dish to cool to room temperature.DO NOT PLACE THE DESSICATORS LID ON THE BENCHTOP. THE GREASE WILL BECOME CONTAMINATED AND NO LONGER SEAL THE DESSICATOR PROPERLY. 3. Record the mass of the cooled evaporating dish (to the nearest 0.001 g). (Note that warm objects cannot be weighed accurately due to the convection currents of the atmospheric gases that are established around warm objects.) 4. Place 0.2 to 0.3 grams of Manganese metal into the evaporating dish. Determine and record the mass of the evaporating dish and metal.


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5. Place approximately 150 mL of water into a 400 mL beaker. Place the beaker on a hot plate in the fume hood and turn the hot plate to High. Later, when the water begins to boil vigorously, turn the heat down. The water should be kept at a low boil throughout the experiment. 6. Place the evaporating dish containing the Manganese on top of the beaker containing the boiling water. CAREFULLY add approximately 20 drops of 6 M Hydrochloric acid to the Manganese metal letting the acid run down the inside surface of the evaporating dish. Gently agitate the evaporating dish to mix the metal and the Hydrochloric acid. Allow the reaction to proceed while frequently agitating the dish. When the reaction appears to be slowing, add 2 to 3 more drops of acid. Continue adding acid and agitating the dish until the Manganese is completely dissolved. 7. When the Manganese has completely dissolved, stop adding acid and allow the solution in the dish to evaporate completely. 8. When all of the liquid in the dish has evaporated, the salt that remains should be a light pink. There should be no brown spots. When you are sure there is no liquid remaining, use beaker tongs to transfer the evaporating dish to the ring and wire screen over a Bunsen burner. 9. Heat the evaporating dish gently. If you see any signs of bubbling, immediately remove the Bunsen burner. Continue to apply and remove heat until there is no bubbling. 10. Heat over a very low flame for an additional 10 minutes. DO NOT GET THE DISH TOO HOT. Watch the Manganese Chloride product. It should stay light pink. If it starts to turn brown, it is getting too hot. 11. Immediately place the evaporating dish in the dessicator to cool to room temperature. 12. Determine the mass of the evaporating dish and Manganese Chloride product. 13. Calculate the mass of Manganese used, the mass of Manganese Chloride product, and the mass of Chlorine that reacted with the Manganese.


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14. From the mass of Manganese and the mass of Chlorine, determine the percent by mass of each element in the product. 15. Each group will write their mass percentages on the board. Determine the average mass percent for each element using the class data. 16. Use the average of the mass percentages for each element to determine the empirical formula of the Manganese Chloride product.

Calculations: DETERMINATION OF CHEMICAL FORMULAS

Mass of Manganese and evaporating dish, (g)

Mass of empty evaporating dish, (g)

Mass of Manganese, (g)

Mass of Manganese Chloride product and evaporating dish, (g)

Mass of empty evaporating dish, (g)

Mass of Manganese Chloride product, (g)

Mass of Chlorine, (g)

Percent by mass of Manganese Percent by mass of Chlorine

Average % Mn Average % Cl

Empirical Formula of the Manganese Chloride product: ___________________________ SHOW CALCULATIONS (Use additional sheets)


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Questions: 1. When 2.50 g of Copper metal reacts with molecular Oxygen, the Copper Oxide product of the reaction has a mass of 2.81 g. What is the empirical formula of the copper oxide product? 2. When 10.8 g of Silver was reacted with Sulfur, 12.4 grams of product was produced (there was only one product). What is the empirical formula of the product?

REFERENCES


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Introduction. In chemical reactions, the atoms of the substances present at the start of the reaction (the reactants) are rearranged into different combinations to produce other substances (the products of the reaction). However, there is no change in the number of each type of atom (mass and atomic identity is conserved). That means that the total number of atoms of each element in the reactants is equal to the total number of atoms of that element present in the products. This principle is used to balance the chemical equation that represents a chemical reaction. As with chemical formulas, chemical reaction equations can be interpreted on a microscopic level in which the coefficients (the numbers in front of each substance) can represent the number of individual units of that substance present. Reaction equations can also be interpreted on a macroscopic level in which the coefficients in the equation represent the mole-to-mole relationships between the reactant and product substances. Background: Investigate the common reactions

Goal: Lear about common reactions Purpose (1) To gain experience with common reactions


Experiment 4: Chemical Reactions Number of sessions: 1


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Material:


6 test tubes

1 test tube gripper

1 test tube rack

Copper(II) Sulfate

Magnesium

Copper

Zinc

25 ml Hydrochloric Acid

Procedure:

A. Metals with Hydrochloric Acid. 1. In three test tubes, each containing about 3 mL of 6-M HCl. A small piece of three different metals will be added separately to the tubes of acid. Copper, Zinc, and Magnesium will be used. Record the appearance of each piece of metal before it is placed in the acid. 2. As each piece of metal is placed in the reaction, carefully observe the metal to determine if bubbles of gas are being formed. If the metal does react with the acid, Hydrogen gas and a soluble metal Chloride compound are formed. For the reaction with Magnesium, feel the tube as the reaction occurs. What do you observe in regard to heat generated? 3. Write a balanced chemical equation for any reaction that occurs (or write no reaction if there is none). INCLUDE PHASE LABELS. B. Zinc and Copper(II) Sulfate 1. Place approximately 3 mL of 1-M Copper(II) Sulfate solution in each of two test tubes. 2. Place a small piece of Zinc in one of the tubes. Keep the other tube as a reference. Observe and record the color of the Copper(II) Sulfate solution.


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3. Place the two test tubes in your test tube rack. Stir the solution containing the Zinc every 15-20 minutes and observe the tubes periodically for about an hour. Write a balanced chemical equation for the reaction that occurred. (Hint: it is a single replacement reaction). INCLUDE PHASE LABELS. 4. Discard the solutions down the drain with lots of water and place any solid remaining in a waste paper basket. DO NOT PUT THE SOLID IN THE SINK.

Calculations: CHEMICAL REACTIONS A. Metals with Hydrochloric Acid Copper Observation____________________________________________________ Balanced Chemical Equation _______________________________________________________________ Zinc Observation____________________________________________________ Balanced Chemical Equation _______________________________________________________________ Magnesium Observation____________________________________________________ Balanced Chemical Equation _______________________________________________________________ What did you observe concerning the heat generated for this experiment? B. Zinc and Copper(II) Sulfate Observation:____________________________________________________ _______________________________________________________________ Balanced Chemical Equation: _____________________________________________________________________________


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Questions: Write balanced chemical equations for each of the following reactions: 1. The decomposition of Diiodine Pentoxide to form Iodine and Oxygen. 2. Silver Nitrate reacting with Potassium Sulfate in a double replacement reaction. 3. The combination of Lithium and Nitrogen to form Lithium Nitride. 4. The decomposition of Potassium Carbonate to form Potassium Oxide and Carbon Dioxide. 5. A single replacement reaction in which Sodium metal reacts with Aluminum Oxide and replaces the Aluminum. 6. The Combustion of C7H1

Manual Chemical Practices

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