Polarity of Chemical Substances Lab Report

Chemistry 120



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ISBN 10:1-64043-130-6
ISBN 13:978-1-64043-130-0

Table of Contents

Periodic Table iv

Laboratory Safety v

Common Lab Equipment viii

Experiment 1: Laboratory Techniques 1

Experiment 2: Heat Stability of Ionic and Covalent Compounds 11

Experiment 3: Specific Heat of an Unknown Metal 15

Experiment 4: Chemical Reactions 21

Experiment 5: Single Displacement Reactions 25

Experiment 6: Analysis of Potassium Chlorate 29

Experiment 7: Emission of Light 33

Experiment 8: Molecular Models 37

Experiment 9: Polarity of Chemical Substances 43

Experiment 10: Neutralization: Titration I 53

Experiment 11: Neutralization: Titration II 57

Experiment 12: Analysis of Hydrates 61

Experiment 13: Analysis of Iron Nails 65

Appendix: Nomenclature Worksheets 69






Periodic Table of the Elements





1.008 2

Atomic number 6

13 14 15 16 17




Lithium 6.941



Beryllium 9.012

Symbol Name





Boron 10.81



Carbon 12.01



Nitrogen 14.01



Oxygen 16.00



Fluorine 19.00



Neon 20.18

11 12

Na Mg

Atomic mass (weight)


13 14 15 16

Al Si P S

17 18

Cl Ar

Sodium 22.99

Magnesium 24.31

3 4 5 6 7 8 9 10 11 12

Aluminum 26.98

Silicon 28.09

Phosphorus 30.97

Sulfur 32.06

Chlorine 35.45

Argon 39.95

19 20 21 22 23

K Ca Sc Ti V

24 25 26

Cr Mn Fe

27 28

Co Ni

29 30

Cu Zn

31 32 33

Ga Ge As

34 35 36

Se Br Kr

Potassium 39.10



Rubidium 85.47



Cesium 132.9

Calcium 40.08



Strontium 87.62





Scandium 44.96



Yttrium 88.91


Lanthanide Series

Titanium 47.87



Zirconium 91.22



Hafnium 178.5

Vanadium 50.94



Niobium 92.91



Tantalum 180.9

Chromium 52.00



Molybdenum 95.96



Tungsten 183.8

Manganese 54.94



Technetium (97)



Rhenium 186.2

Iron 55.85



Ruthenium 101.1



Osium 190.2

Cobalt 58.93



Rhodium 102.9



Iridium 192.2

Nickel 58.69



Palladium 106.4



Platinum 195.1

Copper 63.55









Zinc 65.38



Cadmium 112.4



Mercury 200.6

Gallium 69.72



Indium 114.8



Thallium 204.4

Germanium 72.63



Tin 118.7





Arsenic 74.92







Bismuth 209.0

Selenium 78.96







Polonium (209)

Bromine 79.90



Iodine 126.9



Astatine (210)

Krypton 83.80







Radon (222)

87 88

Fr Ra


Actinide Series





































































Cerium 140.1







61 62 63 64 65 66

Pm Sm Eu Gd Tb Dy

Promethium Samarium Europium Gadolinium Terbium Dysprosium (145) 150.4 152.0 157.3 158.9 162.5

93 94 95 96 97 98

Np Pu Bk Cf

Neptunium Plutonium Berkelium Californium

(237) (244) (247) (251)



Holmium 164.9








Thulium 168.9







Ytterbium 173.0







Lanthanum 138.9







Praseodymium Neodymium











Am Cm

Americium Curium











Lutetium 175.0






Noble Gas

Alkali Metal

Alkaline Earth Metal



Other Metal

Transition Metal



Values in parentheses are the mass numbers of the most stable isotope.

Laboratory Safety

orking in a chemistry lab is a hands-on experience similar to cooking in your kitchen at home. You will

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use different instruments to measure, transfer, and mix things. You may experience a pleasant smell, but an offensive odor would not be out of the question either. You may spill something and need to clean it up. You often will heat or chill a substance. You may need to use a sharp tool or accidentally break a piece of glassware.

All of these actions are accompanied by hazards. You can help minimize, control, and/or manage the hazards of working in a laboratory by adhering to the safety practices described in the guidelines below. Of course, the major difference between lab work and cooking is the use of chemicals, broadly defined as substances that are not approved for consumption. Safe handling of chemicals

is discussed after the general guidelines for laboratory behavior.

General Guidelines for Laboratory Safety

Personal Conduct and Preparation

  1. Always conduct yourself in a professional manner. Horseplay and pranks are strictly prohibited in the laboratory.
  2. Carefully follow all written and verbal instructions. If you are uncertain about how to proceed, ask your instructor.
  3. Prepare for the experiment by carefully reading all assigned background material and procedures before entering the laboratory.
  4. Never work alone in a laboratory. Should an accident occur, it is imperative that help is nearby.
  5. Perform only the experiments and procedures described by your instructor. Do not attempt to perform any unauthorized operation in the laboratory space.
  6. Never eat or drink anything in the laboratory.
  7. Never apply cosmetics, handle contact lenses, or use a cell phone in the laboratory.
  8. Practice good housekeeping while performing a lab. Do not bring unnecessary materials (book bags, notebooks, purses, cell phones, etc.) into the workspace, and keep your assigned area tidy. Clean up dirty glassware and materials as you go along.

Attire and Personal Protection

  1. Always wear shoes that completely cover your feet.
  2. Do not wear loose or baggy clothing.
  3. Long hair and any dangling jewelry should be secured before working in the laboratory.
  4. Always wear chemical splash goggles while in the laboratory.
  5. Wear appropriate chemical resistant gloves when handling hazardous substances and when advised by your instructor. If chemicals are spilled on your gloves, remove them immediately and wash your hands.

Always remove gloves and wash your hands before leaving the laboratory.

Laboratory Operations

  1. When using a sharp instrument, always carry it with the sharp end pointed down and away from you and others. Use caution with sharp instruments, and always cut or puncture items in a direction away from yourself and others.
  2. Carry glassware with two hands and in a vertical position to prevent inadvertently bumping or breaking the glass.
  3. Examine glassware before each use. Never use dirty, cracked, or chipped glassware.
  4. Exercise caution with fire and other heat sources.

Never leave an open flame or heated material unattended. Ensure that the area is free of flammable materials before lighting a flame.

  1. If heating a test tube, always point the open end upward and away from yourself and others.



  1. Hot liquids and steam can cause severe burns. Use caution when handling all heated substances.
  2. Remember that objects usually look the same whether they are hot or cold. Check the temperature of glassware, hot plates, and other materials before picking them up. Use caution if it is necessary to move a hot object.

Accidents and Emergencies

  1. Know the location and operation of all emergency exits and equipment, including: first aid kits, eyewash stations, safety showers, fire alarms, fire extinguishers, and fire blankets.
  2. All spills and accidents should be immediately reported to the instructor.
  3. All broken glass should be disposed of in a properly labeled container, not the trash can. Do not pick up broken glassware with your hands. Your instructor will use a broom and dustpan to retrieve the pieces of glass.
  4. If a chemical splashes in your eyes, immediately flush your eyes with water from the eyewash station for at least 15 minutes. Be sure to hold your eyes open while flushing with water.
  5. If a small amount of a hazardous chemical splashes on your skin, hold the exposed area under running water for at least 15 minutes.
  6. If a large amount of chemical splashes onto your skin or clothing, proceed immediately to the safety shower and wash with water for at least 20 minutes. Remove contaminated clothing as quickly as possible while standing under the shower.

Chemical Safety

There is always some level of risk associated with the use of any chemical. In general, we may think about two major categories of chemical hazards: physical and health. Physical hazards are potential dangers to your physical safety posed by a chemical. For example, many organic chemicals are flammable and therefore increase the risk of fire in the laboratory. Health hazards are associated with acute or chronic biological conditions that potentially result from exposure to a chemical. Ingestion of lead, for example, is known to cause developmental delays and memory loss. The experimental procedures in this text were designed to minimize the use of and potential exposure to chemicals known to present serious hazards to health and physical safety. When possible, traditional solvents and reagents have been replaced with less hazardous alternatives. Nevertheless, you should minimize your risk of exposure to all laboratory chemicals by observing the following general guidelines at all times.

General Guidelines for Handling Chemicals

  1. Always double check the label to ensure that you are using the appropriate chemical.
  2. Remove only the quantity of chemical necessary to achieve the task from the original container. Securely replace the cap as soon as you have removed the chemical from the container.
  3. Always use a scoop or spatula to remove solids from a container. Never handle chemicals with your hands.
  4. Clean all scoops, spatulas, and glassware as soon as possible after you are finished using them.
  5. Dispose of all chemicals in the appropriate waste container, as directed by your instructor.
  6. Never return unused chemicals to the original container. Unused portions should be disposed of in the appropriate waste container.
  7. Keep chemicals under the fume hood or within another properly ventilated space as instructed. When working in a fume hood, the materials should be placed at least six inches from the front sash of the hood.
  8. When transporting a chemical, always hold the container securely with two hands, and proceed slowly and carefully around the laboratory.
  9. While working with chemicals, do not touch your face, eyes, mouth, nose, hair, other body parts, or personal items such as a cell phone. Wash your hands after completing your work and before leaving the laboratory.
  10. You will be instructed in the safe handling of acids. If you are required to dilute an acid sample, always pour the acid slowly into the water. Do not add water to a sample of concentrated acid.
  11. Never taste or smell a chemical by holding it directly under your nose to inhale the vapors. If necessary, you will be instructed on the proper technique for smelling a chemical.
  12. Always wear proper personal protective equipment, including proper attire, chemical splash goggles, and disposable gloves. Never assume that personal protective equipment will safeguard you from every type of chemical exposure. Follow proper emergency procedures in the event of an accident or spill. Always remove gloves and wash your hands before exiting the laboratory.
  13. Be aware of the hazards associated with each chemical you use, and follow any recommended special handling practices.

vi Chemistry 120

The final guideline raises an important question, especially if you have little or no experience in a chemistry lab. You have a right and a responsibility to understand the hazards associated with the chemicals you use in the laboratory. After all, you must first be aware of any hazards if you are going to be able to take action to minimize or control the risk. But how do you find this information?

There are a number of text-based and online reference materials that provide information about the physical properties and hazards associated with many common substances. Your instructor may require that you regularly refer to one or more of these resources. By far, the most common and uniform source of information on a chemical’s known characteristics is the Safety Data Sheet (SDS) (formerly known as the Material Safety Data Sheet or MSDS), which each manufacturer of a hazardous chemical is required by law to produce and make available to users. Your institution should maintain a copy of the SDS for each chemical in the laboratory. Manufacturers also generally make their SDSs available on their website. Throughout this text, each laboratory procedure highlights the most important safety information from the SDSs of chemicals used during the lab. It is therefore critical that you are familiar with certain information and terminology found on an SDS.

The United States has adopted the Globally Harmonized

System of Classification and Labeling of Chemicals (GHS) as the standard format for communicating chemical hazard information on an SDS. Under this system, the SDS is divided into 16 sections, each of which includes a different category of information about the chemical. Hazards associated with the chemical are identified in Section 2 and communicated in the following ways:

7 Pictograms: There are nine graphic symbols used to communicate major categories of hazards. A pictogram will be present if a chemical presents one or more hazards found within the associated category. Some hazards fall into more than one category and therefore may be represented by more than one pictogram. In such cases, the pictogram that carries the higher warning is used. The nine pictograms and associated hazards are summarized on page xiv. Definitions of select terms used to communicate chemical hazards are provided in the table on page xv.

7 Signal Words: There are two signal words used to communicate the overall level of hazard associated with a chemical. The word “Danger” signifies a more severe hazard level, and the word “Warning” is used when the hazards are less severe.

7 Hazard Statements: Hazard statements summarize specific hazards associated with the chemical. Some examples are: “Causes skin irritation”; “Harmful if swallowed or inhaled”; “Extremely flammable liquid and vapor.”

7 Precautionary Statements: Precautionary statements summarize actions to prevent or minimize adverse effects of handling a chemical. First aid information is also provided in the precautionary statements. Some examples are: “Keep away from heat/sparks/open flames/hot surfaces”; “Wear protective gloves/eye protection/face protection”; and “IF INHALED— Remove victim to fresh air and keep at rest in a position comfortable for breathing.”

This text uses the GHS pictograms to visually summarize hazards associated with any chemicals used in the laboratory procedures. In addition, “Safety Notes” inserted in the text highlight specific hazards and suggest appropriate precautions.

There are several other sections on an SDS that students using the manual may also find useful. Section 3 describes the composition of a substance, including the molecular weight. Sections 4 and 5 provide detailed information on first aid and firefighting measures, respectively. Section 9 summarizes the known physical and chemical properties of the substance, including appearance, melting point, boiling point, density, etc. The remaining sections include additional handling, storage, disposal, accident response, and other regulatory information.

Safety is always the top priority in any laboratory. By following the guidelines discussed in this section and any guidance from your instructor, you can avoid accidents and minimize the risk of lab work. If you are ever in doubt about the safest way to proceed, do not hesitate to ask! Work hard, have fun, and always stay safe!

Cuyamaca College Laboratory Safety vii

word image 231 Common Laboratory Equipment

word image 232 Wire triangle e

Clay triangl

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Evaporating dish

Utility clamp

word image 234 Test tubes


Ring stand

Crucible and cover

Ring support

Crucible and lid

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Mortar and pestle

word image 236 Wire gauze

Iron ring

word image 237 Watch glass


Cork stoppers

Crucible tongs


Test tube holder

Rubber stoppers

Erlenmeyer flask

Filter flask


Volumetric pipet

Bunsen burner

word image 278 Test tube brush

Pasteur pipet


Buchner funnel

Stirring rod

Graduated cylinder


Powder funnel



Three-pronged clamp


Long-stem funnel Test tube rack

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viii Chemistry 120


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Laurie LeBlanc & Rosana Pedroza

Laboratory Techniques

Inaccurate experimental results are often a consequence of poor lab techniques. The knowledge and use of good lab techniques will provide an understanding of experimental procedures and assure good results during each laboratory period. They will also prevent loss of product, chemical contamination, spilled chemicals and help the student to minimize waste.

In this lab, students will become acquainted with common laboratory equipment and use it in a variety of laboratory techniques. Students will investigate the use and precision of laboratory balances and volu- metric glassware. The Bunsen burner will be used to fire-polish glass rod and the techniques of evaporation and filtration will be used to distinguish between physical changes and chemical changes of matter. An unknown precipitate will also be identified as part of the filtration.

A: Mass Determination

The accurate determination of mass is one of the most fundamental techniques for chemistry students. Mass is a direct measure of the amount of matter in a sample of substance, that is, a direct indication of the number of atoms or molecules the sample contains.

There are several types of balances used in a typical chemistry laboratory. These balances differ in their construction, appearance, operation and in the level of precision they provide in mass measurements. In this laboratory, a top-loading balance will be used that has a precision of 0.01 grams (g) as well as an analytical balance with a precision of 0.0001 g (see Figure 1.1).

Note the use of the word “balance.” A balance (such as we’re using in this lab) measures mass by balancing the sample of interest against a known (internal) mass, effectively cancelling out the force of gravity. A scale, on the other hand, measures “weight” (the force of gravity along with mass). It is a balance that is routinely used in the chemistry lab, not a scale.


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Top-loading Balance Analytical Balance

Figure 1.1

Follow the instructions below when using any kind of balance:

7 Use the brush next to each balance to clean the pan (and enclosure) before and after each use.

7 Be sure that the balance gives a zero reading when the pan is empty. This can be done by using the “Tare” or “Zero” button on the balance.

7 Balances can be damaged by moisture. Do not pour liquids on the pan; clean up spills.

7 No chemical reagents should be weighed directly on the balance pan. Use a weighing tray, weighing paper or glass vessel.

7 Materials should be placed gently on the balance pan to avoid damage to the balance.

7 For accurate mass determinations, the object to be weighed must be at room temperature. If a warm object is placed on the balance pan, the air around it becomes heated. The warm air rises and the motion of the air may be detected by the balance, giving masses that are significantly different than the true value.

7 Procedures in a lab exercise may be written in such terms as “weigh 0.5 grams of substance to the nearest milli- gram.” This does not mean that exactly 0.500 g of substance is needed. Rather, the statement means that an amount between 0.450 and 0.549 g should be obtained. Unless a procedure states explicitly than an exact amount is to be weighed, you should not waste time trying to obtain an exact amount. But always record the mass actually measured to the precision of the balance used.

2 Chemistry 120

Experimental Procedure

  1. Examine the two different types of balances provided in the laboratory. Starting with one of the balances, record the model, manufacturer and balance number in your notebook. Next, determine and record the mass of an empty weighing tray. Place a rubber stopper in the weighing tray and record the mass of both. Determine the mass of the rubber stopper by subtracting the mass of the weighing tray from the combined mass of weighing tray and stopper. This work should be clearly shown in your notebook next to your mass data. Be sure to include units on your mass.
  2. Repeat this procedure using the second type of balance.
  3. Finally, calculate the average of your two stopper masses in your notebook.

B: Volume Determination

Most of the glassware in the lab has been marked by the manufacturer to indicate the volume contained by the glass- ware when filled to a certain level. The graduations etched or painted onto the glassware by the manufacturer differ greatly in the precision they indicate, depending on the type of the glassware and its intended use. For example, bea- kers and Erlenmeyer flasks are marked with very approximate values that serve as a rough guide to the volume in the container. Other pieces of glassware, notably burets, pipets and graduated cylinders are marked more precisely. The temperature (usually 20°C or room temperature) is specified with volumetric glassware since the volume of a liquid changes with temperature (see Figure 1.2).



100mL A

TD 20 C











250mL 1


TC 20 C

Graduated Cylinder

Volumetric Flask

Figure 1.2

Graduated Pipet

Cuyamaca College Experiment 1 7 Laboratory Techniques

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Graduated cylinder


Eye position is level with the surface

of the liquid

Figure 1.3 A graduated cylinder is used to measure the volume of a liquid. The proper measurement is taken at the bottom of the meniscus, or curve, of the liquid in the cylinder.

The accuracy and precision of volume measurements depend on the type and size of glassware used. The most common apparatus for routine deter- mination of liquid volumes is the graduated cylinder. Although this container does not permit as precise a determination of volume as do other devices, it is often sufficient for use in the lab. The graduated cylinders that you will be using in this experiment have numerical markings that identify the calibration lines. By determining the level of the liquid in the graduated cylinder with respect to these lines, you can determine the volume of the liquid.

When making volume measurements using a graduated cylinder, read the level of the liquid at the bottom of the meniscus. The meniscus is the curved surface at the top of the column of the liquid (see Figure 1.3).

Follow these instructions when reading volume:

7 Determine the smallest graduation of volume that can be read with each cylinder.

7 Read the bottom of the meniscus (use black and white cards to help you see the scale).

7 Always read at eye level.

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Experimental Procedure

  1. For the graduated cylinder provided, record the size (maximum amount of liquid that it can hold) in your notebook.
  2. Record the smallest graduation of the cylinder scale.
  3. Record to what decimal place a volume can be read from the cylinder. (Hint: Read between the lines.)
  4. Finally, record the volume of colored water to the correct precision. Be sure to include units.

C: Laboratory Burners

Almost all laboratory burners used today are modifications of a design by the German chemist Robert Bunsen. In Bunsen’s fundamental design, gas and air are premixed by admitting the gas at relatively high velocity from a jet in the base of the burner. This rapidly moving stream of gas causes air to be drawn into the barrel from side ports and to mix with the gas before entering the combustion zone at the top of the burner. The temperature of a Bunsen burner flame is approximately 700°C (or 1292°F).

The burner is connected to a gas cock by a short length of rubber or plastic tubing. With some burners the gas cock is turned to the fully on position when the burner is in use, and the amount of gas admitted to the burner is controlled by adjusting a needle valve in the base of the burner. In burners that do not have this needle valve, the gas flow is regulated by partly opening or closing the gas cock. With either type of burner the gas should always be turned off at the gas cock when the burner is not in use (to avoid possible dangerous leakage).

4 Chemistry 120

Experimental Procedure

  1. Operation of the Burner. Examine the construction of your burner (Figure 1.4) and familiarize yourself with its operation. A burner is usually lighted with the air inlet ports nearly closed. The ports are closed by rotating the barrel of the burner in a clockwise direction. After the gas has been turned on and lighted, the size and quality of the flame is adjusted by admitting air and regulating the flow of gas. Air is admitted by rotating the barrel; gas is regulated with the needle valve, if present, or the gas cock. Insufficient air will cause a luminous yellow, smoky flame; too much air will cause the flame to be noisy and possibly blow out.

A Bunsen burner flame that is satisfactory for most purposes is shown in Figure 1.5; such a flame is said to be “nonluminous.” Note that the hottest region is immediately above the bright blue cone of a well-adjusted flame.

Outer Cone

Inner Blue Cone


Air Vents

Gas Inlet

Figure 1.5 Bunsen Burner Flame

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Needle Valve

Figure 1.4 Bunsen Burner

  1. Practice lighting and adjusting your Bunsen burner, making any necessary notes in your notebook.

Cuyamaca College Experiment 1 7 Laboratory Techniques


D: Glassworking

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In laboratory work it is often necessary to fabricate simple items of equipment, making use of glass rod and tubing. In working with glass rod, improper techniques may result not only in an unsatisfactory apparatus but also in severe cuts and burns. Therefore, the instructions below should be studied carefully. Prepare two stirring rods using 18 centimeter (cm) lengths of 6 millimeter (mm) glass rod. These rods will be used in future experiments. After they have been completed and cooled, store them in your locker.

Whenever glass is cut it must be fire- polished in order to avoid personal injury. Dispose of broken and non-usable glass in the glass waste container provided.


Experimental Procedure

  1. Scoring the glass tubing: Mark the rod with the edge of your file at the point where it is to be cut. Grasp the tubing on either side of the mark and hold it in position on the laboratory table. Hold the file by the handle and, pressing the edge of the file firmly against the glass at right angles to the tubing, make a scratch on the tubing by pushing the file away from you. If the file is in good condition a single stroke should suffice. Several strokes may be required if the file is dull, but if more than one stroke is needed, all must follow the same path so that only one scratch mark is present on the tubing. The scratch need not be very deep or very long, but it should be clearly defined.
  2. Cutting the glass tubing: Grasp the tubing with your thumbs together directly opposite the scratch mark. Now apply pressure with the thumbs as though bending the ends toward your body and the center part of the rod away from your body, while at the same time exerting a slight pull on the tubing. A straight, clean break should result. Use the flat side of your file to remove any sharp projections from the ends of the cut tubing. After cutting glass in this way, the ends of the cut glass, although clean and flat, are still very sharp and must be fire-polished in order to avoid personal injury.
  3. Fire-Polishing Glass. Fire polishing is the process of removing the sharp edges of glass by heating the glass rod in a burner flame. While continuously rotating the rod, heat the end in the hottest part of the flame until the sharp edges are smooth. When the fire-polishing is completed, remember that the glass is hot even though it looks cool! Place the hot glass rod on the benchtop to cool.

6 Chemistry 120

E: Evaporation

Experimental Procedure

  1. Preparing the water bath. Prepare the water bath setup (refer to diagram below). Add 2-3 boiling chips to the beaker. Use a 250- or 400-mL beaker for the water bath, whichever is a better fit for your evaporating dish. The hottest part of the Bunsen burner flame should be in contact with the bottom of your beaker. Set your ring support to the correct height before beginning to heat your water bath! Tap water may be used in the water bath only.
  2. Preparing the solution. Cover the bottom of a standard test tube (150-mm in length) with a small amount of sodium chloride. Fill the test tube about one-quarter full with deionized (dI) water and stir with your stirring rod until all the sodium chloride is dissolved. Add more dI water if needed to dissolve completely.
  3. Evaporating the solution. Pour the sodium chloride solution into an evaporating dish. Place the evaporating dish on the water bath as shown below. The

water bath should be heated to a boil and then kept constant (replenish water in the water bath from your water bottle, if neces- sary) until all the water in the evaporating dish is evaporated and a residue is left behind.

  1. After all necessary observations have been recorded in your notebook, dissolve the residue with tap water and flush the salt solution down the sink.
  2. Before taking the water bath apart, diagram and label each part of the setup including the ring stand, Bunsen burner, evaporating dish,

ring support, wire gauze, beaker (indicate size used) and boiling chips. (See Figure 1.6.)

Evaporating Dish

Medium Ring Beaker

Medium Ring Wire Gauze

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Figure 1.6 Water Bath Setup

Cuyamaca College Experiment 1 7 Laboratory Techniques


F: Filtration

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The process of separating suspended insoluble solids from liquids by means of filters is called filtration. Insoluble sol- ids, called precipitates, are formed during some chemical reactions. In the laboratory these precipitates are generally separated from the solutions by filtering them out on a paper filter. The liquid that passes through the filter paper is the filtrate; the solid precipitate remains on the filter paper.

Experimental Procedure

  1. Forming a precipitate. Fill a test tube about one-quarter full of lead (II) nitrate solution. Fill a second tube about one-half full of sodium iodide solution. These solutions contain lead (II) nitrate and sodium iodide, each dis- solved in water. Pour the lead (II) nitrate solution into a 100 mL beaker. Slowly pour the sodium iodide solution into the beaker, stir and observe the results. The chemical reaction that occurs forms sodium nitrate and lead

(II) iodide. One of these products is a yellow precipitate.

  1. Preparing the filter. Prepare a filter using filter paper in the following way: First, fold a circle of filter paper in half. Fold in half again:

Step 1

Fold the circular filter paper in half.

Step 2

Fold the semicircle in half to yield a quarter-circle.

Next, open out into a cone. Tear off one corner of the outside folded edge. The top of the cone that touches the glass funnel should be torn. Your instructor will demonstrate this.

  1. Filtering the precipitate: Fit the opened filter paper cone into a short-stemmed funnel, placing the torn edge next to the glass. Wet the paper with deionized water and press the top edge of the paper against the funnel, forming a seal. The funnel will then be placed into a 250 mL Erlenmeyer flask. Next, stir the mixture of prod- ucts in your small beaker, and slowly pour it down the stirring rod into the filter paper in the funnel as dem- onstrated by your instructor. Do not overfill the paper filter cone.
  2. Identification of the precipitate. After the filtration is completed, compare the filtered precipitate with the sam- ples of solid sodium nitrate and lead (II) iodide provided by the instructor to determine which of these is the residue/precipitate on the filter paper.
  3. Chemical disposal: Use the tweezers provided to remove the filter paper with the precipitate and transfer it into the waste jar provided. Pour the filtrate into the waste bottle provided.

8 Chemistry 120

Answer each of the following questions completely. Type your work, do not handwrite.

Use your own words for all answers except for definitions in #1. Use complete sentences with proper grammar and spelling for credit.

Number (and letter, if applicable) your answers clearly. Please make them easy to find.



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  1. Define the following terms using a SCIENTIFIC source. Cite all sources of this information for #1, including the lab procedure itself.
(a) mass(d) meniscus(g) evaporation(j) systematic error
(b) accuracy(e) combustion(h) precipitate(k) random error
(c) precision(f) filtration(i) filtrate
  1. (a) What is the difference between a balance and a scale?
    1. Which of the two will you be using in Experiment 1?
    2. What is the base unit of your mass measurements?
  2. (a) What is the approximate temperature of the flame on a Bunsen burner?
    1. Which part(s) of Experiment 1 will require the use of a Bunsen burner?
  3. (a) What is the size of glass rod that you will be fire-polishing?
    1. Summarize the procedure you will use to achieve smooth ends on your glass rods.
  4. Diagram the following pieces of glassware by hand:
  5. beaker
  6. volumetric flask/stopper
  7. test tube
  8. What is the base unit for your volume measurements?
  9. Diagram and label each part of the laboratory setup (by hand) for the filtration in Part F that shows the solution being filtered as completely as possible.
  10. List all chemical reagents (solids and liquids) you will be using in Experiment 1.
  11. Indicate which portions of Experiment 1 involve (a) physical change and (b) chemical change. Explain your reasoning for each change you include.

Cuyamaca College Experiment 1 7 Laboratory Techniques 9

**For questions below, answer in complete sentences. Use proper capitalization and punctuation. Each answer must make sense in order to be given credit.



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  1. (a) Discuss the two different types of balances used in Experiment 1; explain how they vary with respect to construc- tion. Use your two measurements of the stopper mass as part of your answer.

(b) What kind of error (systematic or random) would you encounter if you were to use a top-loading balance instead of an analytical balance? Clearly explain all answers.

  1. (a) Why is it important to use the same balance throughout the course of an entire experiment?
  2. What kind of error (systematic or random) would you encounter if you used different balances?
  3. Would this affect precision or accuracy? Clearly explain all answers.
  4. If you needed to measure 38.2 mL of a liquid, which of the following would be the best-sized graduated cylinder to use for this purpose? Explain your choice.

10-mL cylinder, scaled to 0.1 mL 50-mL cylinder, scaled to 1 mL 100-mL cylinder, scaled to 1 mL

  1. What kind of error, systematic or random, would you encounter if you read the volume of a liquid in a graduated cylinder above the meniscus? Explain.
  2. Diagram and label each part of the Bunsen burner including the flame. As part of your diagram indicate where the hottest part of the burner flame is.
  3. Why is air mixed with gas in the barrel of the burner before the gas is burned?
  4. How would you adjust a burner that (a) has a yellow and smoky flame?

(b) is noisy with a tendency to blow itself out?

  1. (a) Why are glass rods always fire-polished after cutting?
  2. Why is it important to rotate the glass rod while heating?
  3. In Experiment 1, did you observe a chemical or physical change in (a) Part E? (b) Part F of Experiment 1? Explain your answer for credit.
  4. Give the name and formula of the solid residue remaining after evaporation in Part E.
  5. Give the name, formula and color of the precipitate recovered by filtration in Part F.
  6. Give the name and formula of the two compounds that must be present in the filtrate in Part F.

10 Chemistry 120



Laurie LeBlanc

Heat Stability of Ionic and Covalent Compounds

A = hydrogen B

= carbon

= Cl-

= Na+

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figure 2.1 (A) In a covalent compound like propane (C3H8), the formula indicates the actual number of each atom present in a molecule of the substance. (B) In an ionic

compound like sodium chloride (NaCl), the formula indicates the ratios of the ions present in a crystal. Each Na~ ion interacts with many nearby Cl! ions.

Chemical compounds can be broadly classified as ionic and covalent. An ionic bond results from the transfer of one or more electrons from one atom or group of atoms to another. Ionic compounds are generally composed of a combination of metal and nonmetal. A covalent (or molecular) bond results from the sharing of one or more pairs of electrons between atoms. Covalent compounds are usually composed of nonmetals only.

The molecular stability of ionic and covalent compounds when heated will be observed in this lab by exposing the compounds to the heat of a Bunsen burner (approximately 700°C or 1292°F). Three unknown substances will also be heated in an attempt to determine their chemical makeup (i.e., are they ionic or covalent? metal-containing? nonmetal-containing?).


Experimental Procedure

The following compounds will be available in the laboratory:

aluminum oxide ethanol (C2H6O) magnesium sulfate

ammonium carbonate glycerol (C3H8O3) sucrose (C12H22O11)

potassium chloride sodium chlorate sodium acetate

Unknowns A, B and C

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  1. As preparation for lab, create a chart in your notebook at the end of your procedure summary. The chart will contain the following column headings: (1) Compound Name, (2) Compound Formula, and (3) whether it is Ionic or Covalent. Complete each of the three columns for all nine known compounds. Your instructor will check this work.
  2. Use a pea-sized quantity or smaller of the solid samples, or the equivalent of one drop for liquid samples. Place the sample in a clean, dry deflagration spoon and heat over a Bunsen burner in the hottest part of the flame. Observe and record any changes that you observe on heating. Pull the spoon out of the flame occasionally to see if the compound burns on its own. The ability of a compound to burn on its own is called “ignition.” Note this property if you observe it. Also note all observations if there is a residue after strong heating. You may have to do more than one sample of different sizes for a material to be sure of your observations; for example, if the sample burns away too quickly to observe and record changes, try again with a larger sample. Some samples are difficult to burn off and may require extra heating. If necessary, start again with a very small amount of compound in order to decrease the total time needed for heating.
  3. Each sample that has residue remaining must be heated to red hot for a full 5 minutes at a minimum. To reach this heat, the bottom of the deflagration spoon should be held in the hottest part of the Bunsen burner flame. In some cases if the residue is stuck to the bottom of the spoon, it can be tipped sideways to allow the flame to make contact with the sample – this is particularly true of solid, carbon-containing compounds.
  4. After heating each sample, in addition to recording all of your observations in the notebook, indicate how much material is lost upon completion of heating using one of these descriptors: “all is lost,” “some is lost,” or “none is lost.”
  5. *Clean the deflagration spoon between each compound to avoid contamination.
  6. Heat the unknown compounds last. In your notebook, along with your observations, ** record which known substance each unknown most closely resembles in its heat behavior.

7 Use small quantities of compounds. Do not waste chemicals.

7 Do not contaminate reagent containers. Use a clean, DRY spatula or scoopula to remove chemicals. Immediately replace the cover on containers after removing your sample.

7 Sodium chlorate (NaClO3) is a strong oxidizer. Do not throw it in the trash because it can start a fire. Dispose of any leftover or spilled sodium chlorate in the sink and flush with water.


12 Chemistry 120


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  1. For the 9 known compounds construct and complete a table with the following information included:

*Use one of these qualitative descriptions: all, some, none

Relative Amount of Material Name of Compound Chemical Formula Ionic or Covalent Lost After Heating*

  1. (a) Which compounds appeared to have undergone no material loss on strong heating?
    1. Are they ionic or covalent?
    2. Do they contain polyatomic ions?
    3. Do they contain metals?
  2. (a) Which compounds underwent some (but not total) loss of material on strong heating?
  3. Are they ionic or covalent?
  4. Do they contain polyatomic ions?
  5. Do they contain metals?
  6. (a) Which compounds underwent complete material loss on strong heating?
  7. Are they ionic or covalent? (Pay close attention to the formula for each compound here.)
  8. Do they contain polyatomic ions?
  9. Do they contain metals?
  10. What generalizations can you make about the heat stability of ionic versus covalent compounds?
  11. Did you observe any exceptions to the generalization in question 5 in this experiment?
  12. What generalizations can you make about the types of bonds (ionic or covalent) in each unknown substance?

Give your reasoning.

  1. Which unknown(s) must contain metals? Why?
  2. Which unknown(s) must contain all nonmetals? Why?
  3. (a) Even though residue remained, a student heated the sample at red hot for only 3 minutes instead of a full 5 minutes. Is this a systematic or random error? How might this error affect his results?

(b) Suppose a student heats a sample for a full 5 minutes but NOT in the hottest part of the flame? Is this a systematic or random error? How might this error affect her results?

Cuyamaca College Lab 2 7 Heat Stability of Iconic and Covalent Compounds 13

Laurie LeBlanc



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Specific Heat of an Unknown Metal

The science of calorimetry allows us to measure the temperature change when a quantity of heat is trans- ferred and then use the temperature change along with other information to calculate the amount of heat transferred.

Temperature and heat are related to each other mathematically by the specific heat (C) of a substance, defined as the quantity of heat needed to increase the temperature of 1.00 gram of a substance by one degree Celsius. The units of specific heat are J/gºC. The mathematical relationship is as follows:

q m C  t

The specific heat of a substance relates to its capacity to absorb heat energy. The higher the specific heat of a substance, the more energy is required to change its temperature. The specific heat of metals generally varies with their atomic masses. You will see this relationship later when the data in Table 3.1 is graphed. (Note that atomic mass is the independent variable and specific heat is the dependent variable in your graph; this is because we are examining how specific heat changes as a function of atomic mass.)

In this experiment, we will use calorimetry to determine the specific heat of an unknown metal. Heat energy is transferred from a hot metal to water until the metal and the water have reached the same temperature. This transfer is done in an insulated container called a calorimeter to minimize heat losses to the surround- ings. We then make the assumption that all the heat lost by the metal (qm) is absorbed by the water and is equal to the heat gained by the water, (+qw). Since we know the specific heat of water, we have all the variables needed to calculate qw using the reasoning:

qw = mw  Cw  tw

qw = qm

qw = qm = mm  Cm  tm

This relationship can be used to calculate Cm of a metal because both mm and ∆tm can be measured. See the following sample problem.


Table 3.1 Specific Heat of Selected Metals

Name of Metal Atomic Mass, amu Specific Heat, J/gºC

Sample Problem

A metal sample weighing 68.3820 g was heated to 99.0C, then quickly transferred into a Styrofoam calorimeter containing 62.5515 g of deionized water at a temperature of 18.0C. The temperature of the water in the Styrofoam cup increased and stabilized at 20.6ºC. Calculate the specific heat of the metal (Cm) and identify the metal. (Note that the specific heat of water is 4.184 J/gC.)

tw  20.6  18.0  2.6C

qw  mw  Cw  tw  (62.5515g)( 4.184 J/gC)( 2.6C)

 680 J (heat absorbed by the water)

Then, since all of the heat absorbed by the water came from the hot metal, we can say that:

qw  qm  mm  Cm  tm

tm 20.6C  99.0C  78.4C

680 J  (68.3820 g)(Cm)(78.4C)

Cm  0.13 J/gC

Refer to Table 3.1 and determine whether the unknown metal is lead or gold.

16 Chemistry 120

Experimental Procedure

  1. Record the code letter of your metal cylinder. Using a top-loading balance, weigh your dry metal cylinder.
  2. Carefully slide the cylinder into your largest test tube. Place a thermometer beside it.
  3. Preparing the water bath. Attach the test tube to a ring stand with a two-fingered clamp and place it into an empty 400 mL beaker. Prior to heating, be sure the height of the beaker is adjusted so the hottest part of the burner flame will be on the bottom of the beaker. Stabilize the beaker with a ring support as shown in Figure 3.1.
  4. Fill the beaker with tap water so the height of the water in the beaker is about two inches higher than the top of the metal sample. Add 2-3 boiling chips to the water bath.
  5. Heating the metal cylinder. Begin heating and continue with the next steps. As you are working, check the water and note when it starts to boil. Turn down the burner but keep the water gently boiling so that the metal cylinder will heat up.
  6. Preparing the calorimeter. Next, put together your calorimeter in the following way: Nest two, dry Styrofoam cups together, cover and weigh them. Record the mass of the calorimeter.
  7. Take another metal sample similar to the one you are heating. It is a “stand in” for the metal sample you are heating and will not be used during the experiment. Put this “stand-in” into the calorimeter and add enough dI water to cover the metal by no more than

one-half inch.

  1. Remove the “stand-in,” weigh and record the mass of the calorimeter and the water that you added.
  2. Take the cover for the Styrofoam cup and insert a thermometer through the hole. The

nested cups with the cover and thermom- eter are your “calorimeter.” This will be demonstrated by your instructor. If you set the calorimeter on the bench top it might fall over and break the thermometer so put the whole setup into a beaker to stabilize it.

  1. Measure and record the temperature of the water in the Styrofoam cup (this is your

Metal Cylinder

Thermometer Clamp


2-Fingered Clamp

Large Test Tube

Medium Ring


initial temperature for the water). Leave the thermometer in the calorimeter until you are ready to transfer the hot metal.

  1. After the water in the beaker has been boiling for at least 10 minutes and the temperature inside the test tube with the metal has been stable for 5 minutes, record the high temperature for the metal (this is your initial temperature for the metal). Remove the thermometer from the test tube and set it aside so it does not get mixed up with the thermometer used

in the calorimeter.

figure 3.1

Medium Ring

Wire Guaze

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Cuyamaca College Lab 3 n Specific Heat of an Unknown Model


  1. Transferring the metal cylinder to the calorimeter. It is important that the transfer of the metal cylinder to the calorimeter takes place quickly and carefully (a) to minimize heat loss to the surroundings and (b) to avoid splashing and subsequent loss of water from the calorimeter. Remove the cover and thermometer from the calorimeter. Loosen the test tube clamp on the ring stand; lift the clamp and test tube out of the water bath. Then quickly and carefully slide the metal into the water in the calorimeter being careful not to splash.
  2. Immediately, put the cover with the thermometer back on the Styrofoam cup. Stir gently for 2–3 minutes while monitoring the temperature. Record the highest temperature reached by the water/metal in the calorimeter. (This will be the final temperature for both the metal and water.)
  3. Repeat the experiment a second time.

18 Chemistry 120

  1. What is the code letter of your metal sample? The mass of your metal sample?

**The calculations for number 2–6a should be done twice, one for each run.

  1. Calculate the mass of the water in your calorimeter.


  1. Calculate the change in temperature (∆tw) of the water in your calorimeter after the addition of the heated metal.

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  1. Calculate the change in temperature (∆tm) of the metal after it was added to the calorimeter. (Note: This should be a negative value.)
  2. Calculate the heat gained by the water (qw).
  3. (a) Using the qw from problem 5, calculate the specific heat of your unknown metal. (Remember that qw  qm.)

(b) Calculate the average specific heat for your unknown metal.

  1. Considering the appearance and properties of your metal along with the information in Table 3.1, what do you think the identity of your unknown metal is?
  2. What is the theoretical specific heat of your metal?
  3. Using the following equation, calculate the % error of your value for specific heat and comment on your results.

(Theoretical value Experimental Value )

% Error 

Theoretical value


Use the data presented in Table 3.1 to answer questions 1014 and prepare a graph that shows the relationship between the atomic mass and the specific heat of the eight metals listed. Make the graph following the guidelines in the graphing study aid provided by your instructor.

  1. What is the independent variable for your graph? What is the dependent variable?
  2. In Table 3.1, what is the range of values for atomic mass?
  3. In Table 3.1, what is the range of values for specific heat?
  4. Plot the data in Table 3.1 on the graph paper provided. Include the following:
    1. A title (use y vs. x format)
    2. Placement of the independent and dependent variables on the appropriate axes
    3. Correct increments for each axis (be sure to use the correct scale)
    4. Label each axis, including units
    5. Plot the data points
    6. Connect with a smooth curve
  5. Using your graph, summarize the relationship between the atomic mass of metal atoms and the specific heat of a metal.

Cuyamaca College Lab 3 n Specific Heat of an Unknown Model 19



Laurie LeBlanc

Chemical Reactions

This laboratory will familiarize you with one category of double displacement reactions. The general equation for a double displacement reaction is: AX BY Ò AY BX

You will be learning about precipitation reactions in this experiment. These are reactions in which a solid or “precipitate” is formed when two aqueous solutions are combined. Evidence of this type of reac- tion is the formation of a cloudy (opaque) solution and is produced when two clear solutions are mixed.

**Note that the word “clear” means “transparent” NOT colorless!

Using your observations of the reactions between known compounds, you should be able to identify two unknowns.

Experimental Procedure

  1. Solution preparation. The solutions to be tested are: barium nitrate

magnesium nitrate manganese(II) nitrate silver nitrate

Measure 1 mL of barium nitrate into a standard test tube. Repeat with the remaining three metal nitrates. **Use a grease pencil to mark the 1 mL level on your remaining test tubes (you’ll need a total of 16 for the known solutions). This will prevent you from having to use a graduated cylinder for every measurement, thus simplifying the procedure.

figure 4.1 A cloudy solution indicates the formation of a precipitate.

  1. Each metal nitrate will be mixed with about 1 mL of each of the following four solutions in four test tubes:

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ammonium chloride sodium hydroxide sodium phosphate sodium sulfate

Once again, mark with the grease pencil at the 2 mL level. Repeat on all remaining test tubes.


  1. Use a vortexer to mix each pair of reagents well. Use the centrifuge to separate and observe the precipitate when a cloudy solution results. Record your observations. Note the color and texture of any precipitates formed as instructed.

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  1. Record the manufacturer and model of both the (1) vortexer and the (2) centrifuge in your notebook.
  2. Your observations will be very useful since they are going to be used to identify your unknowns. You will be assigned two unknowns. Unknown A will contain one of the above metals ions (barium, magnesium, manganese or silver). Unknown B will contain two of these metal ions. Test about 1 mL of each unknown with about 1 mL of the four compounds used previously:

ammonium chloride sodium hydroxide sodium phosphate sodium sulfate

7 Balance the centrifuge before using.

7 Silver nitrate (AgNO3) may cause temporary stains on skin.

7 All waste should be placed in the waste container under the hood.


22 Chemistry 120


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  1. For each set of two reactants in Experiment 4 (except unknowns), write a balanced chemical equation for each reaction that should occur using your solubility table – NOT your observations. If both products are in the aqueous phase (aq), write “no reaction.” (Remember that in this case all species are spectator ions!) You must indicate physical states of all species in your written equations.
  2. Do all your balanced equations in question #1 agree with your experimental observations during the lab? If not, explain.
  3. (a) What two unknowns were you assigned?
  4. Identify the metal ion present in Unknown A. Give your reasoning.
  5. Identify the two metal ions present in Unknown B. Give the reasoning behind your conclusions. In your explanation, (1) indicate which ions are definitely present; (2) indicate which ions are definitely absent. Cite specific tests (or reactions) that reinforce these conclusions.

Cuyamaca College Lab 4 7 Chemical Reactions 23

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Laurie LeBlanc

Single Displacement Reactions

The chemical reactivity of certain elements can be determined by using a series of single displacement reactions. Using this method of ranking, an activity series can be created. The general equation for a single displacement reaction is:

A  BC(aq) Ò B  AC(aq)

The reactivity of an element is related to its tendency to lose or gain electrons, that is, to be oxidized or reduced. The more reactive element will be oxidized and the less reactive element will be reduced. In the above general equation, element A is the more active element and replaces element B from the compound BC. But if element B were more active than element A, no reaction would occur. This type of reaction is also known as an oxidation-reduction (redox) reaction.

In this lab, you will perform a series of single displacement reactions to determine the relative reac- tivity of hydrogen and five metals. From the results of these reactions, you will be able to create your own activity series.

Experimental Procedure

    1. Place six clean test tubes in a rack and label them 1-6. To each, add about 4 mL of the solutions listed below in step 2.
    2. Obtain three pieces of sheet zinc, two of copper, and one of lead. Clean the metal pieces with fine sandpaper to expose fresh metal surfaces before putting them in the appropriate test tubes. Add the metals to the test tubes with the solutions as listed.

Tube 1: silver nitrate + copper strip Tube 2: copper(II) nitrate + lead strip Tube 3: lead(II) nitrate + zinc strip Tube 4: magnesium sulfate + zinc strip Tube 5: sulfuric acid + copper strip Tube 6: sulfuric acid + zinc strip


    1. Observe the contents of each tube carefully and record any evidence of chemical reaction.

With some of the combinations used in these experiments, the reactions may be slow. If you see no immediate evidence of reaction, set the tube aside and allow it to stand for at least 15 minutes, then reexamine it.

Evidence of reaction will be either evolution of a gas (bubbles) or appearance of a metallic deposit on the surface of the metal strip. Metals deposited from a solution are often black or gray (in the case of copper, very dark reddish brown) and bear little resemblance to commercially prepared metals.

    1. Dispose of the contents of each test tube into the “heavy metals waste” container. Do not allow the metal strips to go into the sink.
    2. In your notebook, for each test tube write your observations and the balanced, conventional equation for each reaction observed. If no reaction occurred, write “no reaction.” Also, for each test tube indicate which of the two elements is more reactive.

26 Chemistry 120


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  1. Complete a table like the one below and indicate the order of reactivity for the elements in each test tube:
Greater Activity (element symbol)
Lesser Activity (element symbol)
  1. Using your results from the table in #1 above, arrange all five of the metals (excluding hydrogen) in an activity series, listing the most active first.
  2. On the basis of the reactions observed in the six test tubes, explain why the position of hydrogen cannot be fixed exactly with respect to all of the other elements listed in the activity series in question 2.
  3. What additional test(s) would be needed to establish the exact position of hydrogen in the activity series of the elements listed in question 5?
  4. On the basis of the evidence developed in this experiment:
    1. Would silver react with dilute sulfuric acid? Why or why not?
    2. Would magnesium react with dilute sulfuric acid? Why or why not?
  5. (a) Using an online or library source, define “heavy metals.”

(b) Which of the metals used in this experiment would be classified as heavy metals?

Cuyamaca College Experiment 5 n Single Displacement Reactions 27


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Laurie LeBlanc

Analysis of Potassium Chlorate

The goal of this experiment is to perform a quantitative analysis of a chemical reaction. The reaction is the decomposition of potassium chlorate. When potassium chlorate is heated over a Bunsen burner, it decomposes to potassium chloride and elemental oxygen, according to the following equation:

potassium chlorate Ò potassium chloride + oxygen

Solid potassium chlorate contains potassium, chlorine and oxygen. Upon heating potassium chlorate, oxygen escapes as a gas, and a solid residue, potassium chloride, is left behind. The experimental percentage of oxygen in the original sample of potassium chlorate can be calculated using data collected when this decomposition reaction is performed quantitatively. Its accuracy can be determined by comparing it to the theoretical percent of oxygen calculated using the formula of potassium chlorate.

From the experimental data, you will be able to calculate the oxygen lost from your starting sample of potassium chlorate upon heating. Using that information, you will determine your experimental percent oxygen and ultimately calculate the error in that value using the following equation:

% Error  (|theoretical %O  experimental %O|/theoretical %O)  100 (**The difference in the numerator is an absolute value so the % error will be positive.)

Experimental Procedure

A: Quantitative Analysis of KClO3

1 Cleaning the crucible. Place a clean, dry crucible (uncovered) on a triangle and heat for 5 minutes at the maximum flame temperature. The tip of the sharply-defined inner blue cone of the flame should touch and heat the crucible bottom to redness. Allow the crucible to cool to room tempera- ture. Handle only with crucible tongs or paper towel from now on to avoid contamination or added weight from the oils on your hands.

Figure 6.1


  1. Take the mass of the cooled crucible and its cover; record into your notebook. Add 1 to 1.5 g of potassium chlorate to your crucible and replace the lid. Record the mass. The difference of these two masses will be the weight of your potassium chlorate starting sample.

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  1. Decomposing potassium chlorate. Place the partially covered crucible on the triangle as shown in Figure 6.1 and heat gently until no more changes occur. This will help you to avoid the loss of molten material from the crucible. The lid can now be placed completely over the crucible. You will then heat for an additional five minutes at red hot to ensure that all oxygen is driven off. (Be sure that the sample is heated to a high enough temperature by making sure that the crucible is a dull red color during this period. If not, adjust your Bunsen burner.)
  2. When finished, allow the crucible to cool for two minutes in the triangle, then grasp the crucible just below the cover with the concave part of the tongs and very carefully transfer it to the benchtop. Allow to cool completely to room temperature and weigh. Do NOT remove the lid from the crucible to avoid losing material that may have adhered to the inside of the lid.
  3. After weighing, reheat the first sample (without removing the crucible lid) for an additional 5 minutes at the maximum flame temperature (bottom of the crucible heated to a dull red color); cool and reweigh. If the residue is at constant weight, the last two masses should be in agreement. If the mass of the residue decreased by more than 0.05 g between these two mass measurements, repeat the heating and weighing until two successive masses agree within 0.05 g.

B: Qualitative Examination of Residue

  1. After your sample residue has reached a constant mass, number and label three test tubes.
  2. Preparation of three test tubes. Put a pea-sized quantity of potassium chloride into tube 1 and the same amount of potassium chlorate into tube 2. Add about 10 mL of dI water to each of these two tubes; stopper and shake to dissolve the salts. For test tube 3, add a small amount of your crucible residue (enough to cover the tip of your spatula). You may need to scrape the residue out of your crucible to get enough. Add 3 mL dI water to the test tube; stopper and shake. If the sample doesn’t dissolve within a minute, too much residue was added. Add more water.
  3. Testing with silver nitrate. Test the solution in each of the three tubes in the following way: Add 5 drops of 6 M nitric acid (which should serve to completely dissolve any undissolved solid) to each tube and then add 5 drops of 0.1 M silver nitrate solution to each. Mix thoroughly and vortex. Record your observations. This procedure using nitric acid and silver nitrate is a general test for chloride ions. The formation of a white precipitate is a positive test and indicates the presence of chloride ions. Be sure to record complete observa- tions for each of the 3 test tubes.

7 Potassium chlorate is a strong oxidizing agent and can cause fires if put in a trash can. Dispose of the excess in a sink and flush with water.

7 This experiment is particularly hazardous because of particles that may fly out of the crucible while heating. Be sure to wear safety glasses at all times!

7 Dispose of solutions and precipitates containing silver in the heavy metal waste container provided.


30 Chemistry 120

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22.09 g (2nd heating)

Mass of crucible/lid and residue

1.25 g

22.11 g (1st heating)

Mass of KClO3

Mass of crucible/lid and residue

21.32 g

Mass of crucible and lid

While performing Experiment 6, after two successive heatings, a student collected the following data:

Using the information shown above, answer problems 1-6 showing all calculations and work for credit.

  1. Will this student have to do a third heating? Include a calculation as proof of your answer.
  2. Calculate the mass of the residue left behind in the crucible after the second heating.
  3. Calculate the mass lost during heating.
  4. Calculate the experimental % oxygen in the potassium chlorate sample.
  5. Calculate the theoretical % oxygen in potassium chlorate.
  6. Calculate the % error in the experimental % oxygen determination.
  7. Write the balanced equation for the decomposition of potassium chlorate; include the states of all species in the reaction.
  8. (a) After each heating, why should you wait until your crucible is room temperature before weighing it?

(b) If you weigh a warm crucible, would this be considered a systematic or random error?

  1. Potassium chlorate is a strong oxidizing agent like the sodium chlorate used in Experiment 2. How should the student dispose of excess potassium chlorate?
  2. Potassium chlorate and potassium chloride look very similar but react very differently with silver nitrate.
    1. Write the balanced equations for the reaction of silver nitrate with potassium chlorate and the reaction of silver nitrate with potassium chloride. Indicate the states of all species using a solubility table.
    2. How might a student verify that the residue left in his crucible at the end of Experiment 6 has changed from potassium chlorate to potassium chloride?

Cuyamaca College Experiment 6 7 Analysis of Potassium Chlorate 31

**For questions below, answer the following questions showing all calculations and work for credit.

Part A

  1. What is the mass of your original sample (KClO3)?
  2. Calculate the mass of oxygen lost during heating.
  3. Calculate your experimental % oxygen in the potassium chlorate sample.
  4. Calculate the theoretical % oxygen in the potassium chlorate sample.
  5. Calculate the % error in your experimental % oxygen determination.



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  1. If you had forgotten to read the label on the jar carefully and put potassium chloride in the crucible instead of potassium chlorate, how would your results have been affected?
  2. (a) If your sample of potassium chlorate was contaminated with potassium chloride, how would your % oxygen change?

(b) How, specifically, might this error occur in the lab?

  1. (a) If some of your potassium chlorate popped out of the crucible while you were heating it, how would your % oxygen have been affected?

(b) What systematic error might a student have done to cause this to happen?

Part B

  1. In Part B, what observations did you make that led you to believe your residue was potassium chloride? Explain fully for credit.
  2. A student hasn’t properly washed her test tubes (1-3) used for Part B of Experiment 6 and they are contaminated with sodium chloride (NaCl). How will this negatively affect her results for this portion of the lab?

32 Chemistry 120



Laurie LeBlanc

Emission of Light

Light is a form of electromagnetic radiation. Electromagnetic energy includes radio waves, microwaves, infrared rays, visible light, ultraviolet rays, X-rays and gamma rays (see Figure 7.1). Visible light includes energy that the human eye is capable of detecting. White light is visible light that comes from the sun or an incandescent light bulb.


Energy Frequency


400 nanometers 700 nanometers

Cosmic and gamma rays



Visible to human eye


Heat Radio waves

figure 7.1 Electromagnetic spectrum.

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When white light is passed through a prism (diffraction grating and droplets of rain serve the same purpose), the light is separated into the colors of the rainbow: red, orange, yellow, green, blue, indigo and violet (ROY G BIV). White light is produced by the combination of all of these individual colors in the visible spectrum. Red light is the lowest in energy and violet light is the highest.

The current model of the structure of the atom includes the presence of a nucleus surrounded by electrons. Experimental evidence has led to the hypothesis that electrons in an atom exist at only certain allowable locations away from the nucleus, corresponding to certain energy levels.

Electrons can absorb energy from a flame or electric discharge, but only in packets that contain the exact amount of energy necessary to allow the electron to move farther from the nucleus into a higher energy level. When this occurs, the electron is said to be in an “excited state.” When the electron returns to a lower energy level, it emits the previously absorbed energy in the form of packets of light called photons, each carrying an amount of energy equal to the differences between two energy levels.


figure 7.2 Flame Test for K.

This phenomenon can be observed as the emission of light from heated objects (such as the tungsten filament in incandescent light bulbs), flames (such as the flame tests we will do in the lab) and emission tubes (as we will observe). An emission tube is a sealed glass tube with metal plates at both ends. The emission tube is filled with a pure, elemental gas. By applying a high voltage to the metal plates, the electrons of the gas can be excited and give off light. The emission tube is similar in de- sign to neon lights (although the glass of neon light is painted to produce a variety of colors) and fluorescent lights (although the inside of the glass of fluorescent lights is coated with a fluorescent material).

In an atom, many excited states are possible due to the existence of a number of energy levels and possible electron transitions. Therefore, when light is emitted by a large collection of excited atoms, a variety of energies are released as the electrons return to their lowest energy. When these photons pass through a prism or a diffraction grating, they produce a line spectrum, where each line corresponds to photons of a particular energy.

Each element exhibits its own characteristic line spectrum due to differences in the energy levels of atoms of different elements. When ele- ments are heated in a flame, the characteristic color produced and seen by the naked eye is due to the combined effect of the individual spectral lines. Since hydrogen produces four spectral lines in the visible region – you may

only be able to see three through your diffraction grating – one line is violet, one line is turquoise and one spectral

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line is red, the color of the hydrogen emission tube seen by the naked eye is purple (the sum of all the spectral colors).

Experimental Procedure

Part I (to be done in a darkened lab)

White Light Emission

  1. Obtain a diffraction slide and flashlight. Observe the white light through the diffraction slide and note the continuous visible spectrum.
  2. Draw the spectrum with colored pencils in your lab notebook.

Emission Tubes

  1. Note and draw the color of the hydrogen emission tube at the narrowest part of the tube as it appears to the naked eye.
  2. Look at the tube through a diffraction slide. Note the three spectral lines.
  3. Draw their approximate positions and color in your lab notebook. Use black between the emission lines where there is no color.
  4. Label the drawing with the appropriate element name.
  5. Repeat the procedure with the emission tubes of mercury, neon and helium.

34 Chemistry 120

Part II (to be done in a lighted lab)

Elemental Flame Tests

  1. Obtain a platinum wire. Be sure that the tip of the wire is coiled in the shape of a small circle. When not in use, the rod should be kept wire up in your test tube rack. (Care should be taken to avoid bending the wire. It is very soft.)
  2. In order to thoroughly clean the wire, dip the tip of the wire into a 6 M HCl solution (being careful not to break it) and heat the wire in the hottest part of the flame of a Bunsen burner until a blue color remains constant.
  3. Place 3 to 5 drops of each of the following saturated solutions into individually, labeled test tubes: lithium nitrate, copper(II) chloride, potassium chloride, barium chloride, strontium nitrate, calcium chloride and sodium chloride. (You will be observing the characteristic color emitted of the above metal ions when they are burned in a flame.)
  4. Dip the loop of the cleaned wire into the first test tube and then hold the loop in the hottest part of the flame. Note the color of the resulting flame. (It is the INITIAL color that you are looking for.) Draw the flame color carefully in your lab notebook. Label with the element name.
  5. Clean the platinum wire loop with hydrochloric acid, burn it off, and continue observing all known solutions in the flame. You will use the drawings of your known solutions to identify your unknown solutions.
  6. It may be necessary to periodically replace the 6M HCl(aq) if it appears to be contaminated.
  7. Obtain samples of two unknown metal ion solutions, which will be assigned to you, and record their codes in your lab notebook. Perform flame tests on both unknowns, drawing the flame colors of each carefully in your notebook. Determine the identities of your two unknowns using your observations from the flame tests of the previous metal ions.
  8. Dispose of all waste in the proper waste container. Return your CLEAN platinum wire to your instructor.

Cuyamaca College Experiment 7 n Emission of Light 35

**This lab report is to be done entirely in your lab notebook.



  1. (a) Explain what causes light to be emitted (on an atomic level) from the emission tubes in Part I.

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(b) How does the process differ when light is emitted from the heated aqueous metallic solutions in Part II?

  1. (a) Which unknowns were you assigned in Part II?

(b) What were the metals present in your unknowns?

  1. Research the phenomenon of light emission and explain three practical applications. To receive credit, you must explain clearly how each works to the best of your ability. Make your writing understandable; define any technical

terms used. Cite all sources used.

36 Chemistry 120


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Laurie LeBlanc

Molecular Models

This is a dry lab and should be done in pencil. The lab is intended to be done independently by each student with the help of a molecular model kit so that the various geometries and their angles can be clearly visualized.

Any resonance structures should be included in the Lewis Dot Structure box ONLY. Also be sure to enclose your Lewis Dot structure in brackets with a charge if structure is for a polyatomic ion.


SPECIESValence Electrons

Available for Bonding

Lewis Dot StructureElectron Pair Geometry3-D DrawingMolecular GeometryBond Angles
Ammonia, NH3
Hydronium ion, H3O


Chemistry 120

SPECIESValence Electrons

Available for Bonding

Lewis Dot StructureElectron Pair Geometry3-D DrawingMolecular GeometryBond Angles
Ammonium ion

Ozone, O3
Sulfur Dioxide


Cuyamaca College

Experiment 8 n Molecular Models

SPECIESValence Electrons

Available for Bonding

Lewis Dot StructureElectron Pair Geometry3-D DrawingMolecular GeometryBond Angles
Carbon Disulfide
Boron Trichloride
Hydrogen Selenide, H2Se
Beryllium Fluoride, BeF2


Chemistry 120

SPECIESValence Electrons

Available for Bonding

Lewis Dot StructureElectron Pair Geometry3-D DrawingMolecular GeometryBond Angles
Chlorite ion
Nitrate ion
Sulfite ion
Cyanide ion, CN


Cuyamaca College

Experiment 8 n Molecular Models



Laurie LeBlanc

Polarity of Chemical Substances

In this experiment, solubility properties will be used to determine the polarity of a group of compounds. Two solvents, water and hexane, will be mixed with a variety of covalent solutes. Depending on whether each solute is soluble in the given solvent, it can be discovered whether each is polar, nonpolar, neither or both. Two ionic compounds will also be tested for solubility.

Experimental Procedure

Lab Period 1

The following substances will be available for testing:

acetone, C3H6O biphenyl, C12H10

butyl acetate, C6H12O2 glycerol, C3H8O3 iodine

magnesium oxide sodium bromide

  1. Test the solubility of each of the above substances first in hexane and then in water (the test tubes MUST be dry for the tests in hexane). For all substances use a small amount (pea-sized quantity of solid or the equivalent of about 3 drops if a liquid) per 2 mL of solvent.
  2. Mix solutions well using a vortexer. The evidence of solubility is CLARITY. If two substances are soluble, once mixed you will observe only one phase; if insoluble, you will observe at least two phases (see Figure 9.1). Record your observations for each mixture as

well as if each is soluble or insoluble.

  1. Three commercial substances (motor oil, aspirin and antifreeze) will also be tested. Determine the solubility of each of these in hexane and then in water.


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figure 9.1 (A) Soluble. (B) Insoluble


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  1. Place iodine solutions or any leftover/spilled iodine in the waste container labeled “IODINE.”
  2. Place ALL OTHER solutions in the 2nd waste container provided. Nothing goes in the sink!


Lab Period 2: Using ChemDraw

ChemBIODraw (commonly known as ChemDraw to chemists) is a program designed for drawing molecules and reactions. The software is easy to use, but the more time you spend exploring available options and practices using the program, the easier it will be to properly complete your future lab reports. In this assignment, you will be asked to complete the drawings of 11 molecular structures. You will need to turn in a copy of your work with your lab report for Experiment 9. It should be placed at the end of your Analysis and before the Conclusion of the lab report.

Part A: User’s Guide & Tutorials

Once you have logged on and are in the ChemDraw program, go to the Help menu and click on ChemBIODraw Basics. You will find a good comprehensive overview of the icons on your ChemDraw toolbar in chapter 6 in the User’s Guide. Alternatively, under the heading “Tutorials,” Tutorial 1 will give you some basic information about drawing simple structures. Look over this information to familiarize yourself with the basic use of the toolbar in drawing simple structures, both two- and three-dimensional.

Part B: Experiment 9 Structures

For the listed substances, you will do a drawing correctly depicting all angles in the molecule, show all element sym- bols that are involved in each and every bond, lone pair electrons (so that all elements have an octet), and show dimensionality where appropriate (tetrahedral electron pair geometry). Label each molecule with its name so that it is easy to find.

Iodine, Water, Hexane, Acetone, Biphenyl, Ethylene Glycol, Glycerol, Butyl Acetate, Aspirin

Title this document “Experiment 9 Molecular Structures” and print it before leaving lab.

44 Chemistry 120


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  1. Prior to the start of Period 1 lab, create and complete a table in your notebook (data and observations section) identifying the (column 1) chemical name, (column 2) chemical formula, and (column 3) Lewis dot structure of all compounds (including commercial substances) tested and both solvents (water and hexane). Use a scientific source to look up each structure. Add two more columns to the table to be completed when this is done.

Here are some helpful hints for your structures:

n To approximate the structure of motor oil (a mixture of various hydrocarbons), use the structure for decane (C10H22). For hexane (C6H14) and decane, the carbons are connected in a straight chain.

n For acetone (C3H6O), connect the carbons in a straight chain, and attach the oxygen to the middle carbon.

n To draw glycerol (C3H8O3) and antifreeze (ethylene glycol, C2H6O2), connect the carbons in a straight chain, and attach one oxygen to each carbon.

n Note that the chemical name for aspirin is acetylsalicylic acid.

  1. Now go back to structures completed (EXCEPT glycerol, antifreeze, butyl acetate, and aspirin) and make a prediction as to whether they should be polar, nonpolar or ionic (column 4) based on the number and types of bonds found in each. (Calculate  electronegativity of each bond to make your determination); finally (column 5) predict solubility (should each dissolve in water or hexane?).
  2. Prior to the start of Period 2 lab, complete the following. Download ChemDraw, using the following link: http://sitelicense.cambridgesoft.com/sitelicense.cfm?sid=768

n Use your student email account (firstname.lastname@gcccd.edu) to download the program.

n Go through the ChemDraw tutorial located under the “Help” menu. In your notebook, create a procedure for using ChemDraw that you can refer to during the lab period on Day 2.

Cuyamaca College Experiment 9 n Polarity of Chemical Substances* 45


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  1. Fill in worksheets 1, 2, and 3 on the following pages. You may use pencil for these worksheets only. In the structure box, draw each structure clearly showing every bond present in the molecule; to the best of your ability for each molecule that contains both polar and nonpolar bonds, judge whether the molecule should be polar, nonpolar, or both. Also note that there are two compounds that are ionic. Which solvent do you predict they should more readily

dissolve in: hexane or water?

  1. A student tests the solubility of an unknown substance. The substance dissolves in water (that is, it acts polar). The unknown substance could be ionic or covalent. Give three experimental methods the student could use to determine whether the unknown substance is ionic or covalent.

46 Chemistry 120

Substance Formula Structure (show all bonds, ∆en, label P/NP)

Predicted Experimental Polarity Solubility Solubility Polarity

Acetone C3H6O

Iodine I2



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Cuyamaca College

Experiment 9 n Polarity of Chemical Substances*

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Cuyamaca College

Experiment 9 n Polarity of Chemical Substances*

Substance Formula Structure (show all bonds, ∆en, label P/NP)Predicted Experimental
Polarity Solubility Solubility Polarity
Butyl AcetateC6H12O2
Sodium BromideNaBrNa–Br (No structure)
Magnesium OxideMgOMg–O

(No structure)

Substance Formula Structure (show all bonds, ∆en, label P/NP)

Predicted Experimental Polarity Solubility Solubility Polarity

Acetylsalicylic Acid



Ethylene Glycol (Antifreeze)


Motor Oil

Mixture of Hydrocarbons

C–C and C–H only (No structure)

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Cuyamaca College

Experiment 9 n Polarity of Chemical Substances*



Laurie LeBlanc

Neutralization: Titration I

The reaction of an acid and a base to form a salt and water is known as neutralization. In this experiment potassium hydrogen phthalate (HKC8H4O4, abbreviated “KHP”) is used as the acid. Despite its complex formula, we see that the reaction of KHP with sodium hydroxide is very simple. One mole of KHP reacts with one mole of NaOH.

HKC8H4O4  NaOH à NaKC8H4O4  H2O

You will use the volumetric analysis technique of titration in this experiment. Titration is the process of measuring the volume of one reagent required to react with a measured volume or mass of another reagent. In this experiment we will determine the molarity of a base (NaOH) solution from data obtained by titrating KHP with the base solution. (This is called “standardizing” the base.) The base solution is

figure 10.1

added from a buret to a flask containing a weighed sample of KHP dissolved in water. From the mass of KHP used, we calculate the moles of base needed to neutralize this number of moles of KHP since one mole of NaOH reacts with one mole of KHP. We can then calculate the molarity of the base solution from the titration volume and the number of moles of NaOH in that volume.

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In the titration, the point of neutralization, called the end-point (or equivalence point), is observed when an indicator, placed in the solution being titrated, changes color. The indicator selected is one that changes color when the stoichiometric quantity of base (according to the chemical equation) has been added to the acid. The indicator used in this titration is phenolphthalein, an organic acid. Phenolphtha- lein is colorless in acid solution but changes to pink when the solution becomes slightly basic. When the number of moles of sodium hydroxide added is equal to the number of moles of KHP originally present, the reaction is complete. The next drop of sodium hydroxide added changes the indicator from colorless to pink.


Experimental Procedure

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  1. Preparation of acid solution. Obtain some solid KHP in a test tube or vial. Weigh two samples of KHP into 125 mL Erlenmeyer flasks, numbered for identification. First weigh the flask, then add KHP to the flask by tapping the test tube or vial until 1.0000 to 1.2000 g has been added; record the mass of the flask and the KHP. In a similar manner, weigh another sample of KHP into the second flask.
  2. To each flask add approximately 30 mL of distilled water. If some KHP is sticking to the walls of the flask, rinse it down with water from a wash bottle. Swirl the flasks until all the KHP is dissolved (warm the solutions if necessary to dissolve all KHP).
  3. Preparation of buret. Obtain a buret and rinse it three times with dI water as shown by your instructor.
  4. Obtain about 100 mL sodium hydroxide of unknown molarity in a clean, dry 250 mL beaker.

** Keep your base solution covered with parafilm when not in use.

  1. After rinsing your buret with dI water three times, rinse it with a 5 to 10 mL portion of the base, making sure that the NaOH runs through the buret tip. Discard the NaOH waste in a labeled, waste container. Next, fill the buret with the base, making sure that the tip is completely filled and contains no air bubbles.
  2. Adjust the level of the liquid in the buret so that the bottom of the meniscus is somewhere below the 0.00 mL mark. Record the initial buret reading. (Read and record all buret volumes to the nearest 0.01 mL.)
  3. Add 3 drops of phenolphthalein solution to each 125 mL flask containing KHP and water. Place the first (Sample 1) on a piece of white paper under the buret, extending the tip of the buret into the flask.
  4. Titrating the KHP solution. Titrate the KHP by adding base until the end-point is reached. The titration is conducted by swirling the solution in the flask with the right hand (if you are right-handed) while manipulat- ing the stopcock with the left. As base is added, you will observe a pink color caused by localized high base concentration. Toward the end point, the color flashes throughout the solution, remaining for a longer time. When this occurs, add the base drop by drop until the end point is reached, as indicated by the first drop of base that causes a faint pink color to remain in the entire solution for at least 1 minute. Read and record the final buret reading.
  5. Determine the approximate amount of base to perform a second titration and refill the buret if necessary.
  6. Repeat the titration with Sample 2.
  7. Calculate the molarity of the base in each sample. If these molarities differ by more than 0.004 M, titrate a third sample.
  8. When you are finished with the titrations, empty and rinse the buret at least twice (including the tip) with tap water and three times with deionized water.

54 Chemistry 120

  1. What is the name of the indicator used in this experiment?
  2. Determine the color the indicator will turn in the following solutions:
    1. a strongly acidic solution
    2. a strongly basic solution
    3. your solution at the end point in Experiment 10
  3. What is the (a) chemical name, (b) chemical formula and (c) molar mass of KHP?



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  1. If you had added 50 mL of water to a sample of KHP instead of 30 mL, would the titration of that sample then have required more, less, or the same amount of base? Explain.
  2. A student weighed out 1.1090 g of KHP. How many moles of KHP did she weigh out?
  3. A titration required 20.35 mL of 0.1635 M NaOH solution. How many moles of NaOH were in the solution?
  4. A student weighed a sample of KHP and found it weighed 1.500 g. Titration of this KHP required 18.76 mL of base (NaOH). Calculate the molarity of the base.

Cuyamaca College Experiment 10 n Neutralization: Titration I 55

  1. Calculate the number of moles of KHP in each of your samples.
  2. Using the information in problem 1, calculate the molarity of your base in each sample.


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  1. What is the average molarity of your base, using your two best titrations? (Be sure to record this information in your notebook because you’ll need it for Experiment 11.)

56 Chemistry 120


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Laurie LeBlanc

Neutralization: Titration II

In this experiment, two sets of titrations will be done. In the first, the molarity of an unknown concen- tration of hydrochloric acid will be determined using your standardized sodium hydroxide from Experi- ment 10. The reaction that takes place is:

HCl(aq)  NaOH(aq) Ò H2O(l)  NaCl(aq)

In the second set of titrations, you will determine the mass percentage of acetic acid in commercial vinegar by titrating the acid with your standardized sodium hydroxide solution according to the follow- ing equation:

HC2H3O2(aq)  NaOH(aq) Ò H2O(l)  NaC2H3O2(aq)

Notice that both acids react with sodium hydroxide on a 1:1 mole basis. This relationship will be used to determine the concentrations of hydrochloric acid and acetic acid in a series of titrations.

Experimental Procedure

A: Titration of an Unknown Hydrochloric Acid Solution

  1. In your notebook, record the average molarity of sodium hydroxide you calculated in Experiment

10. This is the sodium hydroxide you will be using in Experiment 11.

  1. Obtain a 40 mL sample of hydrochloric acid of unknown molarity in a clean, dry 50 ml beaker as directed by your instructor. Record the code of the unknown acid.
  2. With a volumetric pipet, transfer a 10.00 mL sample of the acid to a labeled, clean, but not necessarily dry, 125-mL Erlenmeyer flask. Pipet a duplicate 10.00 mL sample into a second, labeled flask.
  3. You will need about 150 mL of the standardized sodium hydroxide whose average molarity you calculated in Experiment 10. Keep the beaker containing the base covered with parafilm when not in use.
  4. Clean and set up a buret. Rinse the buret with a 5 to 10 mL portion of the base, running the base through the buret tip. Fill the buret with the base, making sure that the tip is completely filled and contains no air bubbles. Adjust the level of the liquid in the buret so that the bottom of the meniscus is below the 0.00 mL mark and record the initial buret reading to the hundredths place.


  1. Add three drops of phenolphthalein and about 25 mL of distilled water to the flask containing the 10.00 mL of hydrochloric acid solution. Place this flask on a piece of white paper under the buret and lower the buret tip into the flask.
  2. Titrate the acid by adding base until the endpoint is reached as you did in Experiment 10. Record the final buret reading when the titration is complete.
  3. Refill the buret and repeat the titration. The volume of base used for both titrations should differ by no more than 0.20 mL. If they differ by more than this amount, titrate a third sample.

B: Acetic Acid Content of Vinegar

  1. Obtain about 40 mL of vinegar in a clean, dry 50 mL beaker. Record the % acidity from the vinegar label into your notebook.
  2. Clean and rinse the pipet with vinegar before pipeting the vinegar samples. Pipet 10.00 mL of vinegar into your two, labeled 125 mL Erlenmeyer flasks.
  3. Titrate the duplicate 10.00 mL samples of vinegar with sodium hydroxide using exactly the same procedure outlined in Part A.
  4. When you are finished with the titrations, empty the buret and rinse it and the pipet with tap water and three times with deionized water.

58 Chemistry 120

Part A. Molarity of Unknown Hydrochloric Acid

  1. What is the code letter of your unknown acid?


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  1. What was the average molarity of the sodium hydroxide solution from Experiment 10? 3 Calculate the moles of acid used to neutralize the sodium hydroxide in each sample. 4 Calculate the molarity of the acid in each sample.
  2. What is the average molarity of your unknown acid?

Part B. Acetic Acid Content of Vinegar

  1. Calculate the moles of acetic acid used to neutralize the sodium hydroxide in each sample.
  2. What is the molarity of acetic acid in each sample of vinegar?
  3. What is the average molarity of acetic acid in the vinegar?
  4. Calculate the grams of acetic acid per liter from the average molarity.
  5. Calculate the mass percent acetic acid in the vinegar sample. (The density of vinegar is 1.005 g/mL.)
  6. What is the percent error for your % acidity in question 10? Discuss your results.

Cuyamaca College Experiment 11 n Neutralization Titration II 59



Laurie LeBlanc

Analysis of Hydrates

A hydrate is a compound that contains a salt weakly bound to a number of water molecules. Because the water is attached to the salt by a weak bond, it can be driven off by heating. In this lab, you will analyze hydrates both qualitatively and quantitatively. A hydrated salt can be converted to the anhydrous form (without water) by heating:

Hydrated salt Ò Anhydrous salt  water

Using this technique, (a) you will qualitatively compare the liquid driven off from the hydrate, copper(II) sulfate pentahydrate (CuSO4 • 5H2O) to pure water. (b) You will next heat an unknown hydrate and, by gravimetric analysis, determine the percentage of water in the hydrate and its empirical formula.

Experimental Procedure

A: Qualitative Analysis of a Hydrate

    1. Take a strip of paper provided and fold it in half lengthwise to make a trough to deliver the hydrate, CuSO4 • 5H2O. Weigh out about 4 grams of the hydrate and load it carefully and evenly into the trough.
    2. Refer to Figure 12.1 when setting up the device to collect liquid driven off from your hydrate while heating. Begin by using a two-fingered clamp to hold your large test tube (25200 mm). Attach the clamp to a ring stand and adjust the tube so that it is about 20° above horizontal. Carefully deliver the hydrate into the test tube, making sure that it sits on the bottom of the tube – not on the sides of the tube.
    3. Now adjust the tube so that it is about 20° below horizontal and the mouth of the tube should rest on another large test tube resting in an Erlenmeyer flask as shown. The rubber clamp should be close to the opening of the

upper test tube to avoid melting it while heating the hydrate.

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figure 12.1


    1. Slowly begin to heat the hydrate with a Bunsen burner, rotating the flame around the bottom of the test tube where the hydrate rests. Do not heat too quickly or too hot to avoid decomposing the hydrate into CuO, which is black. The blue hydrate should slowly fade in color from blue to white. By moving the flame up and down the tube, you will be able to drive the liquid into the lower test tube.

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    1. Allow the liquid to cool. Using a strip of cobalt chloride test paper, dot one end with a drop of the liquid from the hydrate. Note any color change. Now do the same with deionized water on the other end of the test paper. Note any color change. Do they both behave the same? Attach the test paper to your notebook with tape.
    2. Place the anhydrous residue left behind in the test tube on a clean, dry watch glass. Using a spatula, divide it into two parts. Add about 4 drops of the liquid from the hydrate onto the residue. Note any changes. Next add about 4 drops of deionized water onto the other half of the residue. Note any changes. Do they both behave the same?

B: Quantitative Analysis of an Unknown Hydrate

    1. Record the code of your unknown hydrate as well as the formula for the salt portion of the compound in your notebook.
    2. Using a metal triangle balanced on a small, metal ring attached to a ring stand (as in Experiment 6), heat a clean, dry crucible to red hot for five minutes to thoroughly clean it. Handling it only with crucible tongs (to avoid contamination), weigh the crucible and lid. Place between 2 and 3 grams of the unknown into the crucible. Record the mass.
    3. Place the covered crucible into the triangle with the lid slightly ajar to allow water vapor to escape easily as the sample is heated. Begin heating gently and continue until no more changes occur. At this point, heat until red hot and continue heating for a full five minutes.
    4. Cover completely and cool to room temperature. Weigh the crucible, lid and sample.
    5. Once again, heat the crucible at red hot for five minutes. Cool and weigh again. The last two masses should be within 0.05 g. If not, repeat the heating and weighing steps until two successive masses are within 0.05 g.

All waste for Parts A & B should be placed in the designated waste container.


62 Chemistry 120

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A 2.914 gram sample of gypsum, a hydrated salt of calcium sulfate, is heated in a crucible until a constant mass is reached. The mass of the anhydrous CaSO4 salt is 2.304 grams. Use this data to answer the following questions. Be sure to show all work for credit.

  1. Calculate the percent by mass of water in the hydrate described above.
  2. (a) How many moles of water were removed when the hydrate was heated?

(b) How many moles of anhydrous calcium sulfate (residue) remain in the crucible?

  1. (a) Calculate the empirical formula of the hydrate.

(b) Name the hydrate.

  1. Define all terms in bold in the introduction of Experiment 12.

Cuyamaca College Experiment 12 n Analysis of Hydrates

Part A



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  1. Were you able to prove that the liquid in copper(II) sulfate pentahydrate was most likely water? Give two qualitative pieces of evidence as proof.
  2. Write the balanced equation for the decomposition of copper(II) sulfate pentahydrate.

Part B

  1. What is the name and formula of the salt in your unknown hydrate?
  2. Calculate the mass of your original sample before heating. 5 Calculate the total mass lost by your hydrate upon heating 6 Calculate the percentage of water in your unknown hydrate.
  3. (a) Determine the empirical formula for your assigned hydrate.

(b) How would you name this compound?

  1. A student heated a hydrated salt sample with an initial mass of 4.8702 g. After the first heating the mass had decreased to 3.0662 g.
    1. If the sample was heated to constant weight after reheating, what is the minimum mass that the sample can have after the second weighing? Show how you determined your answer.
    2. The student found that the mass lost by the sample was 1.8053 g. What was the percent water in the original hydrate? Show all work.
    3. What experimental error would explain results that showed a percent water that was higher than it should be? Lower than it should be?

64 Chemistry 120

Laurie LeBlanc


Analysis of Iron Nails


figure 13.1

An oxidation-reduction (redox) reaction will be performed in this lab. It will take place between iron found in common household nails and a solution of copper(II) sulfate. In a single displacement reaction, copper(II) ion will be reduced to form solid copper. You will isolate the copper using the technique of vacuum filtration. During the redox reaction the solid iron in nails will be oxidized to

form either the Fe2 or Fe3 ion. You will determine the oxidation number of the iron from the results of your experiment.

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CuSO4 dissociates (or ionizes) in water to form the ions Cu2 and SO -2. The copper reacts with the iron nails in solution to form iron ions in solution (Fe2 or Fe3; you will determine which one) and solid copper. By remembering the definition of a spectator ion, you should see why SO -2 qualifies as a spectator in this reaction. In this reaction electrons are exchanged between the elements copper and iron. Each iron atom will lose electrons while each copper atom gains electrons. The overall reaction can be divided into two half-reactions: reduction and oxidation.



The reduction half-reaction: The chemical species that undergoes reduction gains electrons. Cu2 gains two electrons to produce the neutral solid, copper, as follows:

Cu2  2 electrons Ò Cu

(aq) (s)

The oxidation half-reaction: The chemical species that undergoes oxidation loses electrons. Iron solid will lose either two electrons to form Fe2 or lose 3 electrons to form Fe3:

Fe(s) Ò Fe2  2 electrons or Fe Ò Fe3  3 electrons

(aq) (s) (aq)

In order to obtain an equivalent number of exchanged electrons in an overall balanced reaction, a half-reaction must often be multiplied by a coefficient. The table on the next page shows how to deter- mine the overall balanced reaction for each possible iron product. First, look at the left hand column. The number of electrons used is already the same as the number of electrons produced, so the reaction equations do not need to be multiplied by coefficients. Combine the two half-reactions to give your overall balanced reaction. In the second column (if Fe3 is produced), the two half-reactions must be multiplied by the appropriate number to make numbers of electrons equal. Then the two can be added together to give the overall balanced equation.


A. If Fe+2 is the product:B. If Fe+3 is the product:
Reduction: Cu2  2 e Ò Cu(s)


Reduction: 3 x [Cu2  2 e Ò Cu ]

(aq) (s)

Oxidation: Fe(s) Ò Fe2  2 e


Oxidation: 2 x [Fe(s) Ò Fe3  3 e]


Overall balanced equation:

Cu (aq)  Fe(s) Ò Fe (aq)  Cu(s)

2 2

Overall balanced equation:

3 Cu2  2 Fe Ò 2 Fe3  3 Cu

(aq) (s) (aq) (s)

**Note that the number of electrons in the overall balanced equation cancel, as equal numbers are present on both sides.

Your experimental data will tell you which iron ion is the final product because you will calculate the molar ratio of Cu(s) formed to Fe(s) consumed.

word image 409

Experimental Procedure

  1. Using a pipet, add 50.00 mL 0.5 M CuSO4 solution to a 250 mL beaker (the concentration of the solution will be approximately 0.5 M. Be sure to record the EXACT concentration of the solution into your notebook).
  2. Obtain two clean, dry iron nails. Record the mass of the nails to the nearest 0.0001 g on the analytical balance.
  3. Place the nails into the beaker with the CuSO4. Record the start time and let the reaction proceed for 2 hours, gently swirling every few minutes to displace any copper that has plated out on the iron. As the nails become coated with copper, use the rubber policeman on your stirring rod (and maybe a second stirring rod) to scrape them clean so that the iron on the nails is in continual contact with the solution. During the two hours, record any observations. Note the reaction’s ending time.
  4. Using gloves and tweezers, pick up the nails, one at a time. With the rubber policeman, scrape any excess copper from the nails back into the beaker. Without disturbing the copper already in the beaker, swirl the policeman and scraped nails in the liquid remaining in the beaker to rinse off any remaining copper and finally rinse both into the beaker with dI water to avoid the loss of any copper. Place the scraped nails onto a paper towel to dry.
  5. After the nails are completely dry, measure the mass of the nails and record.
  6. Obtain a piece of filter paper and place it into a Buchner funnel. Record the combined mass of the funnel top and filter paper.
  7. Set up the vacuum filtration as demonstrated by your instructor. Place the filter assembly into a filter flask that will serve to collect the liquid. Wet the filter paper with deionized water and start the vacuum. Slowly pour the contents of the beaker (including the solid copper) onto the filter paper, letting it follow your stirring rod; rinse the beaker with dI water as necessary to remove all solid from the beaker.
  8. Next, using a Pasteur pipet, rinse the solid in the filter assembly with about 25 mL of 1 M hydrochloric acid followed by about 25 mL deionized water and finally about 25 mL acetone.
  9. Leave the copper in the filter apparatus with the vacuum on for at least 15 minutes to dry it. Remove and cover with a Kimwipe and rubber band until the next class so that the sample dries completely.
  10. At the start of the next lab period, determine the combined mass of the copper and funnel top. You can then calculate the mass of the copper solid formed using the difference in recorded masses.

Waste Disposal

Put ALL liquid chemical waste in the waste container provided. This includes the liquid in your filter flask. Put solid waste (i.e., copper and iron and filter paper) in the jar for solid waste. Disposable pipettes go in the glass disposal boxes. Bulbs for disposable pipettes are NOT disposable and should be returned.

66 Chemistry 120


word image 410


  1. For each set of two half-reactions, indicate which is the oxidation and which is the reduction.
    1. Cu2  2e Ò Cu and Zn Ò Zn2  2e
    2. Au3  3e Ò Au and Na Ò Na  e
  2. For each set of half-reactions in question 1, combine them into an overall balanced equation as shown in the table on the second page of this lab. Show all work.
  3. Calculate the number of moles of Cu2+ you would be using in this experiment if the solution concentration of copper(II) sulfate is 0.5012 M. [Hint: All copper comes from the copper(II) sulfate solution.]
  4. What is the difference between a galvanized nail and a plain iron nail? Be specific as to composition. (Cite your source for this answer.)
  5. (a) Why is neglecting to do a final rinse of the stirring rod (with policeman) into the beaker containing the accumulated copper a source of error in this experiment? (b) What would the effect of this error be on your mass of copper?

(c) What would the effect of this error be on the ratio of mol Cu/mol Fe?

Cuyamaca College Experiment 13 n Analysis of Iron Nails 67

  1. Calculate the total number moles of Cu2 available at the beginning of the reaction.
  2. What mass of copper was produced in the reaction?
  3. Calculate the moles of copper produced in the reaction. 4 What was the mass of iron consumed in the reaction? 5 Calculate the moles of iron consumed.


6 You could possibly have produced either Fe+2 or Fe+3 ions since both are possible oxidation numbers for iron ion.

word image 411

Below are the two possible reactions:

2 2

Cu  Fe Ò Cu  Fe or

(aq) (s) (s) (aq)

3 Cu2  2 Fe Ò 3 Cu  2 Fe3

(aq) (s) (s) (aq)

  1. Reexamine your calculations in questions 3 and 5 and use those answers to determine the ratio of moles of copper solid produced to moles of iron solid consumed (moles Cu/ moles Fe).
  2. Does your data agree with the stoichiometry shown in one of the overall chemical reactions above?

Which iron ion has been formed?

10 (a) Calculate the % error in your molar ratio:

% error = [(experimental ratio) – (expected ratio)] / (expected ratio) x 100

(b) Comment on the magnitude of your % error.

68 Chemistry 120


Group A Ionic Compounds/Polyatomic Ions

  1. Give the name or formula, as required:



word image 263

Name: Formula:

  1. sodium sulfide
  2. magnesium nitride
  3. calcium oxide
  4. lead(II) sulfate
  5. ammonium chloride
  6. potassium iodide
  7. aluminum bromide
  8. lithium phosphide
  9. beryllium nitrate
  10. strontium chlorate
  11. Na2O
  12. CsF
  13. SnF4
  14. Al2O3
  15. BaSO4
  16. NH4Br
  17. Mg3P2
  18. Mg3(PO4)2
  19. K2CO3
  20. Be(OH)2

Cuyamaca College Appendix n Worksheet 1 71

Group A Ionic Compounds/Oxoanion Series

  1. Give the name or formula, as required:



word image 264

Name: Formula:

  1. sodium nitrate
  2. potassium chlorite
  3. barium perchlorate
  4. bismuth sulfate
  5. cesium nitrite
  6. lead (IV) hypochlorite
  7. rubidium phosphate
  8. calcium nitride
  9. aluminum hyponitrite
  10. beryllium sulfite
  11. CaSO4
  12. K3PO3
  13. Sn(NO2)2
  14. KClO4
  15. Ca(ClO3)2
  16. BeCO3
  17. Mg(ClO)2
  18. Pb(OH)2
  19. Li2SO3
  20. SrS

Cuyamaca College Appendix n Worksheet 2 73

Group B Ionic Compounds

  1. Give the name or formula, as required:



word image 265

Name: Formula:

  1. nickel(II) fluoride
  2. manganese(II) carbonate
  3. zinc nitrate
  4. silver sulfate
  5. cobalt(II) chlorate
  6. chromium(VI) chloride
  7. titanium(IV) sulfide
  8. mercury(II) oxide
  9. cadmium fluoride
  10. iron(III) chloride
  11. AuCl3
  12. CuS
  13. AgNO3
  14. VO
  15. TiCl4
  16. PtO2
  17. Mn3P2
  18. PdO
  19. ScCl2
  20. PbS

Cuyamaca College Appendix n Worksheet 3 75


  1. Give the name or formula, as required:



word image 266

k. Fe(NO3)3 • 9 H2O

j. MnSO4 • 4 H2O

i. Na2CO3 • H2O

h. CuSO4 • 5 H2O

g. KF • 2 H2O

f. Nickel(II) nitrate hexahydrate

e. Calcium sulfate trihydrate

d. Copper(II) chloride dihydrate

c. Sodium carbonate decahydrate

b. Magnesium sulfate heptahydrate

a. Cobalt(II) chloride dihydrate



l. CoCl2 • 6 H2O

Cuyamaca College Cuyamaca College Appendix n Worksheet 4 77

Covalent Compounds

  1. Give the name or formula, as required:



word image 267

Name: Formula:

    1. dinitrogen pentoxide
    2. boron triiodide
    3. phosphorus pentachloride
    4. diiodine nonoxide
    5. dinitrogen tetroxide
    6. carbon tetrachloride
    7. carbon disulfide
    8. dichlorine heptoxide
    9. oxygen monofluoride
    10. carbon monoxide
    11. P2S3
    12. P4O10
    13. SCl4
    14. BCl3
    15. OF2
    16. ICl
    17. NCl3
    18. SiF4
    19. N2O5
    20. P2O5

Cuyamaca College Appendix n Worksheet 5 79


  1. Give the name or formula, as required:



Name: Formula:

  1. sulfuric acid
  2. nitrous acid
  3. phosphoric acid
  4. phosphorous acid
  5. hydrochloric acid
  6. chloric acid
  7. perchloric acid
  8. carbonic acid
  9. hydroiodic acid
  10. chlorous acid
  11. H2SO3
  12. HBr(aq)
  13. HNO3
  14. H3PO3
  15. HNO2(aq)
  16. H2CO3
  17. HClO

Bromic Acid

  1. HBrO3
  2. HBrO2
  3. HBrO4

word image 268

Cuyamaca College Appendix n Worksheet 6 81


  1. Give the name or formula, as required:



word image 269

Name: Formula:

  1. diphosphorus trisulfide
  2. perchloric acid
  3. phosphoric acid
  4. diiodine nonoxide
  5. hydroiodic acid
  6. nitrogen dioxide
  7. hypochlorous acid
  8. hydrogen iodide
  9. carbon tetrachloride
  10. diiodine nonoxide
  11. CS2
  12. HCl(aq)
  13. H2CO3
  14. OF
  15. HBr
  16. SiCl4
  17. P2O5
  18. ICl
  19. Cl2
  20. HNO2

Cuyamaca College Appendix n Worksheet 7 83

Mixed Ionic Compounds

  1. Give the name or formula, as required:



word image 270



k. Mg(NO)2

j. copper(I) chloride

i. ammonium nitride

h. beryllium bicarbonate

g. silver nitrate

f. iron(III) chlorate

e. potassium phosphide

d. zinc hydroxide

c. sodium sulfate

b. sodium sulfite

a. sodium sulfide

l. MnO

q. AuCl3

p. Ni(ClO4)2

o. BaF2

n. CrPO3

m. Al(NO2)3

r. PtO2

s. SrI2

Cuyamaca College Appendix n Worksheet 8 85


  1. Give the name or formula, as required:



word image 271

Name: Formula:

  1. sodium hypochlorite
  2. manganese(II) chlorate
  3. cesium hydroxide
  4. lead(IV) sulfide
  5. ammonium iodide
  6. vanadium(II) oxide
  7. aluminum sulfite
  8. lithium phosphide
  9. cobalt(II) chloride dihydrate
  10. beryllium bicarbonate
  11. SnF4
  12. Sr3(PO3)2
  13. AuCl3
  14. Ag2O
  15. CuSO4•5H2O
  16. NH4Br
  17. Mg3P2
  18. PtO
  19. FeCO3
  20. Be(OH)2

Cuyamaca College Appendix n Worksheet 9 87

Various Compounds (Ionic/Covalent/Hydrate/Acid)

word image 272

  1. Give the name or formula, as required:



Name: Formula:

  1. sodium hypochlorite
  2. diiodine nonoxide
  3. chromium(III) oxide
  4. hydroiodic acid
  5. ammonium phosphite
  6. copper(II) chloride dihydrate
  7. nitrogen
  8. carbon monoxide
  9. tetraphosphorus decoxide
  10. nitric acid
  11. HNO3
  12. BaCl2
  13. B2F4
  14. CoF2 • 6 H2O
  15. P2O5
  16. N2O4
  17. SnCO3
  18. Sr(NO3)2
  19. I2O9
  20. HNO2

Cuyamaca College Appendix n Worksheet 10 89

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