LAPC Chemistry Lab Report

NO RETURNSChemistry 60
LABORATORY EXPERIMENTS FOR
BEGINNING
CHEMISTRY
Proceeds of this manual are used for Pierce student scholarships
Department of Chemistry
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Contents
Introduction ………………………………………………………………………..……6
Laboratory Manners ………………………………………………………………..……8
Tips on Technique ….………………………………………………………………..……9
Measurements and Error …..……………………………………………………….……10
Use of Laboratory Equipment to Measure Mass and Volume………………………..…..14
Your Laboratory Notebook …………………………………………………………..…16
How To Write Your Lab Report ………………………..………………………………18
EXPERIMENT 1. Measurement and Volume ……………………………………..……24
EXPERIMENT 2. Preparation and Measurement of the Density of Solutions …….…….36
EXPERIMENT 3. Distinctive Properties of Chemicals ………………………….………42
EXPERIMENT 4. Separation of Components of Mixtures …..………………….……….48
EXPERIMENT 5. Solubility and Chemical Change ……………………………………..54
EXPERIMENT 6. Chemical Reactions and Equations …………………………….……68
Guide To Chemical Reactions Supplement- Six Simple Types….…..76
EXPERIMENT 7. Determination of a Chemical Formula………………….……………84
EXPERIMENT 8. Qualitative Analysis..…..…………………………………….………92
EXPERIMENT 9. Mass Relationships in a Chemical Reaction ………….…………….109
EXPERIMENT 10. Titration and Solution Stoichiometry ……………………………..113
APPENDICES
Laboratory Safety Rules
Student Academic Integrity Policy
Student Discipline Procedures
Student Grievance Procedures
Answers to Sample Quizzes
Laboratory Safety Quiz
Locker Item List
Some Common Laboratory Equipment
List of Common Cations
List of Common Anions
Common Monoatomic Ions in the Periodic Table
Significant Figures Workshop
Measurement and Calculation Workshop
Conversion Factor Workshop
Percentage Workshop
Nomenclature/Formula Writing Workshops
Chemical Calculations Workshop
Stoichiometry Workshop
Chemical Reactions Workshops
How To Create A Graph Workshop
Ion Flashcards
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INTRODUCTION
A laboratory is a place where an investigator can examine some small part of the world of nature.
The scientific investigator is looking for answers to specific questions about the composition and
behavior of matter under various conditions. To find the answers to the questions the investigator
creates situations in which certain substances and processes can be observed and measured. The
purposeful creation of situations to observe is called experimentation.
Observations that can be made with your own senses, without using a measuring instrument, are
qualitative observations. The fact that an insoluble substance forms when two solutions are
mixed is a qualitative observation. The observations that hydrogen sulfide has a foul odor and
that copper sulfate forms a blue solution in water are qualitative in nature.
How much insoluble material forms when measured amounts of two solutions are mixed is a
quantitative question. To answer it, a measurement must be made with an appropriate measuring instrument. In general, qualitative observations lead to answers about what is happening;
quantitative measurements help answer how much.
EXAMPLE 1:
Does an insoluble precipitate form when a silver nitrate solution and a hydrochloric acid solution
are mixed?
Answer (by direct observation): Yes.
In many cases the measurement or direct observation you make is not the answer to the question
posed. It is often necessary to deduce (figure out) from other known facts and relationships just
what the observation means with respect to what you want to find.
EXAMPLE 2:
What is the insoluble substance formed when silver nitrate and hydrochloric acid solutions are
mixed?
You cannot answer this question by direct observation because one white powder looks pretty
much like any other white powder. You must deduce the answer from known facts about the solubility of the possible products of the reaction. This means you must know what species are present and which ones can react with each other. By examining the properties of the possible products, you can arrive at a conclusion about the probable identity of the white powder.
EXAMPLE 3:
Is aluminum denser than water?
From the direct observation that a solid piece of aluminum sinks in water, one could deduce that
aluminum is denser than water.
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EXAMPLE 4:
What is the density of aluminum?
This is a quantitative question, so we must make some measurements. However, there is no
instrument that measures the density of solids directly. We must measure something that is related to density and calculate what the density must be. To know what measurements to make,
we must know what the relationship is.
Each experiment in this course will be proposed to you as a problem, either qualitative or
quantitative, for you to solve by making the appropriate observations or measurements and deductions. The background facts and relationships will be reviewed. Untested assumptions about
the experiment will be examined. Then we will decide what measurements we can make with the
available apparatus to find the answer to the problem.
Knowledge of the characteristics of the materials, substances, and apparatus to be used in a
particular experiment can be obtained by reading the laboratory textbook and by prior planning
of the experiment and your laboratory notebook. You will be amazed at how this will speed and
smooth your journey through this laboratory course.
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LABORATORY MANNERS
Your laboratory experience will be pleasant if everyone in the laboratory observes courteous
behavior and common sense. Common sense is easier to follow if you know what you are doing
and why you are doing it. The following points are offered as reminders of good laboratory
manners.
1. Protect your eyes with approved safety eyewear at all times. You, of course, will be careful
and never have an accident, but you never know about your neighbor!
2. Dress for lab: always wear closed toed shoes that protect the tops and soles of your feet while
you are in the laboratory. People do spill things, and concentrated acid between the toes is a thrill
you will never forget. Tie long hair back and do not wear loose flowing garments.
3. Let nothing touch your lips during a laboratory session. The laboratory is not a kitchen.
Food, drinks, and gum can too easily become contaminated by toxic (poisonous) substances.
Avoid touching your face with your hands- if you must, wash them first. Wash your hands
before you leave lab to avoid contaminating your lunch.
4. Find out where fire extinguishers, safety showers, and eye-wash fountains are and know when
and how to use them.
5. In case of chemical spills on skin or clothing, wash the affected area immediately with plenty
of water; then tell the instructor.
6. Follow the directions in your lab manual. If you think you have a better way of doing an
experiment, discuss it with the instructor. Please do not experiment on your own yet.
7. Keep your work area clean and dry. You will discover that spills and breakage occur less
often if you keep your glassware at the back of the bench. Look before you reach.
8. Take only as much chemical as you need from the stock bottles. Never return unused
chemicals to a stock bottle; you might contaminate the rest of the contents of the bottle.
9. Use the fume hood or lab snorkels for reactions which release toxic or noxious gases- your lab
manual will direct you when these are to be used.
10. For your health and that of the environment, dispose of chemicals properly- your lab
manual will indicate where to dispose of used chemicals.
11. Focus and awareness is important in the laboratory- so do not use any electronic devicescell phones, texters, ipods, etc.- while in lab.
12. When using common equipment, like balances or shared reagent bottles, always leave them
clean when you are done and return them to their original location.
13. Electronic balances are expensive and delicate instruments, read the posted instructions
before using them. Never add chemicals to a container on a balance, remove it first, add the
chemical and then return it for reading.
14. Return all of your glassware and lock your lockers at the end of lab. Return all common
equipment (rings, test tube clamps, wire gauze, burets, etc.) to where they belong.
15. Report all accidents to your instructor.
16. Place all backpacks, jackets and books not being used in a cubby on the back wall.
17. Read the experiment carefully and prepare your lab notebook before each lab
experiment. Use common sense and ask your instructor whenever you are unsure.
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TIPS ON TECHNIQUE
1. Never heat a beaker or flask directly with a flame. Use a wire gauze support to spread the
heat. A direct flame may cause the glass to crack.
2. Never heat soft (non-Pyrex) glass or plastic apparatus. Graduated cylinders, funnels, vials,
etc., will break or melt every time. Glassware labeled Kimex or Pyrex can be heated.
3. Never heat volumetric glassware (those used to accurately measure volumes) such as
volumetric flasks, pipets, or burets. Even if they do not break the heat will distort them and
destroy their accuracy.
4. When dispensing a liquid from a reagent drop bottle, do not touch the tip to your test tube,
spot plate or flask. This can contaminate the drop bottle. Always drop the reagent holding the
dropper bottle above your container.
5. Keep books, papers, and personal gear off the lab bench- store them in the cubbies in the lab.
The only exception is your laboratory notebook. Soggy textbooks are hard to read and
impossible to sell.
6. ALWAYS follow the instructions in your lab book and if you are not sure about anything ask
your instructor.
7. Do not leave a burner or an apparatus with running water unattended. Beakers boil over or boil
dry when you are not looking, and hoses invariably disconnect when you turn your back.
8. Always heat the contents of a test tube by placing the test tube in a hot water bath. Heating a
test tube directly in a flame can be dangerous to those around you. The heated liquid can be
ejected from the tube and burn someone else in the lab.
9. Do not waste water- it is a precious commodity in California.
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MEASUREMENTS AND ERROR
SIGNIFICANT FIGURES IN MEASUREMENTS AND CALCULATIONS
Measurement is the operation of comparing a property of a substance to the same property of a
standard. We find the length of an object by counting the number of standard lengths laid end to
end that correspond to that dimension. We find the mass of an object by counting the units of
standard mass which experience the same force (weight) in the earth’s gravity. Since every
measurement amounts to a physical comparison, there is always a certain amount of inexactness
or uncertainty in the value of that measurement. This error or uncertainty may come about in
several ways:
1. The measuring instrument used may not reproduce the standard accurately. One does not
use a ruler to construct an airplane, for example.
2. The measuring instrument may not be sensitive to small differences in the value of a
measurement. A truck scale is not used to measure the weight of the truck driver.
3. The measurement taken with a certain instrument may not be reproducible in repeated
trials (precision). A precise instrument reproduces a measurement within a very small
interval.
4. Different methods of measuring may result in different values.
Notice that carelessness and mistakes are not included in the list. The errors and uncertainties
that we are discussing are those that a conscientious worker cannot avoid with the available
measuring devices. You should be aware of these limitations so that you do not place more reliance on your answers than the data warrants.
ACCURACY
The accuracy of a measuring device can be determined by measuring a primary standard. If a
mass balance reports an average value of 1 kilogram for a standard kilogram, it is accurate, at
least at that one value. If a thermometer reads zero degrees Celsius in a mixture of pure water
and ice, it is accurate at that temperature, because the freezing point of pure water is defined as
0°C.
If the instrument is accurate at one value it probably gives true values at other measurements not
too far removed from the test value. In general, however, a number of different independent
measurements must be made to test for accuracy. If the measurements all agree, the value
probably is true.
For example, we will determine the volume of a metal object by three independent methods. We
will determine the volume (1) from its dimensions, (2) from the indicated liquid level change on
immersion in a graduate, and (3) by application of the buoyancy law. If the values all agree
within a certain interval, we can be confident about the accuracy of the value.
SENSITIVITY
The sensitivity of a measuring device is the smallest unit that can be read or estimated of the
property it is measuring. An analytical balance (used in Chemistry 101) will give a measurable
response to the weight of the fingerprint of the operator (± 0.0001 gram). The milligram top
loading balance we will use in this class will not be affected by fingerprints but can be affected
by air currents in the lab. It has a sensitivity of ± 0.001 gram and is ten 10 times less sensitive
than the analytical balance. The capacity and sensitivity of an instrument are usually inversely
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proportional to each other. If a measuring device has a large capacity or range, it is generally not
very sensitive. This makes it difficult to detect small differences in large measurements.
PRECISION
The precision of a measurement is related to the reproducibility of the value. If an archer puts
five arrows within a two-inch circle he is shooting precisely. If the circle happens to be twelve
inches from the bull’s-eye, he is precise but not very accurate.
The precision of an instrument can be determined only by repeated measurements of the same
property of an object. The precision is related to the sensitivity of the measuring device and is
usually less than the last decimal unit that can be read (or estimated) on the scale of the
instrument.
The average weight value is 50.88 / 5 = 10.18 grams.
The deviation listed to the right is the amount each value differs from the average. The average
deviation is the sum of the absolute value of each deviation, divided by the number of trials. In
this case, the average deviation is 0.20 / 5 = 0.04 g. That is, on the average, there is an
uncertainty of ±0.04 in our answer. We can say that the probable weight of the object is 10.18
±.04 grams. The value has four significant figures, even though the last one is somewhat
uncertain. The digits 1, 0, 1, and 8 are all significant.
If we divide the average deviation by the average value, we obtain a measure of the relative
uncertainty of the measurement. In this measurement the relative uncertainty is
0.04g/10.18g = 1/254 or one part in 254.
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Balance 2: The scale on this balance can be read without estimate to 0.01g, and the thousandths
unit estimated. On this balance the values for the mass of the same object might be:
The value can be reported to five significant figures: 10.182 ±0.008 grams. The relative uncertainty of the balance is 0.008/10.182, or one part in 1270. This balance is five times as precise as
Balance 1.
SIGNIFICANT FIGURES (or DIGITS)
The number of significant figures in a measurement is an indication of the precision of the
measurement. The last digit in a value is presumed to be uncertain. If you report that you obtained 1.0 gram of product, the reader should presume that you mean something between 0.5 and
1.5 grams. If you report 1.00 gram, the reader knows that you measured it to the nearest tenth of
a gram and there is some uncertainty about the third significant figure. If not specified, the
uncertainty in the last significant figure is presumed to lie within 5 units of the reported digit. For
example, 1.283 means 1.283 ±.005, or “at least 1.278 and not more than 1.288.”
The rules for counting significant figures are as follows;
All non-zero digits are significant
Zeros between non-zero digits are significant
Leading zeros are not significant
Trailing zeros are significant if a decimal point is shown; otherwise they are assumed to
be insignificant.
Exact numbers have infinite significant figures.
Defined numbers have infinite significant figures.
If the measurement has no significant figures to the right of the decimal point, it is best to write
the number in exponential form to indicate the number of significant figures. For example, you
cannot be sure how many significant figures there are in 2500. If there are four, you should write
2.500 X 103 If only two significant figures are justified by the measurement, record it as 2.5 X
103.
If we have had no prior experience with an instrument, a single measurement leaves us in doubt
about the precision of the recorded value. We should take enough measurements to treat the
values statistically. However, time is also important, and after sufficient experience with one
instrument, you will know what its inherent precision is. Then you will know how much confidence to place on a single measurement.
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USE OF LABORATORY EQUIPMENT TO MEASURE MASS
AND VOLUME
Measuring the Mass of Solids
Using the electronic balance
In this lab, we will use the electronic balance to measure the mass of an object. Electronic
balances have many advantages over other types of balances mainly the ease with which
measurements are obtained. To measure the mass of an object follow the rules below.
Make sure the electronic balance is placed on a flat, stable surface. Wind, shaky surfaces,
or any kind of force can affect the reading accuracy of the balance.
Press the “ON” button and wait for the balance to show zeroes on the digital screen. If the
balance was already on and not reading 0.000 g, press the “zero/tare” button and wait for
the display to read zero. The symbol “ ” will appear on the left side of the display when
the balance has stabilized.
Place the solid object on the balance and wait for the mass to stabilize. When the symbol
“ ” appears, record the mass in your notebook. Always report all the digits in the mass
that the balance indicates. For our electronic balances, it’ll be to 3 places after decimal.
If you are measuring the mass of a liquid or a solid chemical, you must weigh the
substance in a beaker or a flask. To automatically deduct the weight of the container,
press the “Tare” or “zero” button while the empty container is on the balance and wait for
“ ” to appear. Carefully add the substance to the container. Ideally this is done while the
container is on the balance but if it’s easier, you may remove the container from the
balance, add the appropriate amount of chemical, then return the container to the balance
to determine the mass of the chemical. Record the mass as indicated on the display.
While determining the mass of an object, use the same balance for the entire experiment.
Measuring the Volume of Liquids
The pieces of laboratory glassware we will use in this semester for measuring volume include the
beaker, graduated cylinder, volumetric flask, Erlenmeyer flask, pipet, and buret.
To Deliver and To Contain
The glassware used for measuring liquid can be divided into two categories: those that contain
and those that deliver a specified amount of liquid. Glassware that contain include graduated
cylinders and volumetric flasks (labeled with “TC” meaning “to contain”); for these, the
inscribed markings increase in value from bottom up. Glassware that deliver include pipets and
burets (labeled with “TD” meaning “to deliver”); for these, the inscribed markings increase in
value from top to bottom, thus indicating how much liquid has been delivered not how much
remains in the glassware.
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Glassware Used for Measuring Volume
Some of the glassware in the chemistry lab is inscribed with markings for measuring the volume
of liquids. This includes beakers, Erlenmeyer flasks, graduated cylinders, and burets. Even
though beakers and Erlenmeyer flasks have markings on them they are generally not accurate
enough for measuring volume. Thus, these pieces of glassware are primarily used to hold liquids
during experiments. When measuring the volume of a liquid using the graduated cylinder or
buret, the rule is to estimate one more digit past the smallest division on the measuring device.
Other glassware, such as the volumetric flask and the pipet only have a single line that indicates
volume since they are used to measure only one specific volume and thus are much more
accurate than the other measuring devices.
The Meniscus
When liquid, such as water, is placed in a glass or plastic container, surface tension causes the
surface of the liquid to take on a curved shape, called a meniscus. Lab glassware used for
measuring liquid is calibrated so that an accurate result is obtained by reading the bottom of the
meniscus, when it is viewed at eye level. Viewing the meniscus at any other angle will give
incorrect results. The accuracy of the markings on the different glassware vary significantly. The
50 mL graduated cylinder we will use in this lab have a precision of ±0.5 mL, the 10 mL
graduated cylinder has a precision of 0.1 mL, the buret with a precision of 0.05 mL, the pipet
0.01 mL.
Note that the buret and pipet are more precise than the graduated cylinder.
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YOUR LABORATORY NOTEBOOK
A laboratory notebook is a permanent record of experiments carried out in the laboratory.
Everyone who may be occupied in scientific or technological research must learn to keep a
proper record of their experimental conditions and observations. There are a number of reasons
for this practice. Nobody’s memory for details is infallible, and it is often necessary to repeat an
experiment or report the findings in a communication. The notebook serves as an original source
for the data. For people doing original research, the laboratory notebook is a legal record that
establishes their claim of discovery.
You will benefit in several ways by keeping a careful record of your laboratory experience. The
experiments in this text are illustrative of most of the major lecture topics, so your notebook will
help you review for examinations. Careful observing and recording habits will give you an
advantage in subsequent courses where notebooks are required. Communication skills are
valuable in employment applications, too.
The laboratory notebook is your diary of your laboratory experience. As such, it may contain
anything that you feel is pertinent. The minimum content of the notebook must show:
1. What was done
2. When it was done
3. What was observed
4. Who performed the experiment
To be of the greatest value, all notebook entries of observations and data should be made at the
time of the experiment. Notes made on stray scraps of paper are worthless. Entries should be
made in ink in a permanently bound notebook with sequentially numbered pages, dated and
signed at the bottom of each page. In research laboratories, it is also required that each page be
witnessed by a competent person who is familiar with the work.
A laboratory worker sets up or contrives a situation to observe in such a way that the observations or measurements answer a question. Each experiment therefore starts with a question to
be answered.
Before you start any laboratory work, you should write a general outline of what you propose to
do and how the work relates to the problem you seek to solve. In your outline, note the known
facts and assumptions which bear on the problem.
As you complete each phase of the experiment, describe your actions in sufficient detail so that
someone else could repeat the experiment if necessary. Name or describe the apparatus used, and
give the amounts of all substances employed. Record the measurements and observations in ink
at the time that you make them. List repetitive measurements in parallel columns when repeated
trials are made.
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USE OF BOUND LABORATORY NOTEBOOK
Information obtained from scientific work is of interest to scientists from all over the world,
especially if they are investigating similar things. When a set of successful experiments is
completed and analyzed the results are typically published in a scientific journal allowing anyone
to see what was done. These results can now be carefully scrutinized and added to, or corrected,
by other scientists to increase our body of knowledge. Eventually this information may also be
of interest to nonscientists as well, as it will help us to better understand the natural world around
us and make more informed decisions about how to live our lives.
The information obtained from experiments in its simplest form is called data and it serves as a
record of what occurred during an experiment. Without the data there is nothing to support the
results or conclusions that are made. It is necessary that an accurate and permanent record of the
data be made at the time it is collected. Without this record the experimental work cannot be
published. This record is also needed should the experimental work be challenged by other
scientists or in a court of law. All experimenters, whether working in academics or industry,
must keep a notebook with these records. It is a legal requirement. To help you to follow this
regulation you will be keeping a notebook of all of your experimental data during this semester.
The notebook should be clearly organized and labeled using neat columns or tables whenever
possible. It takes some experience to learn how to construct these tables properly so to help you
initially sample data tables will be given in this lab book for the first few experiments. Copy
these tables neatly into your notebook using ink before coming to lab. When in lab, write all of
your data directly into these tables.
Please follow these guidelines when using your notebook for lab reports.
Use black or blue ball point pen only and do not at any point use whiteout; if you make a
mistake, simply put a line through it and write the corrected information next to it.
NEVER erase any work which is incorrectly recorded. Draw a neat line through the
incorrect number and make the new record nearby:
21.08g 21.48 g
1. Write your name, your instructor’s name, course number, and lab time on the front label.
Write your last name across the cut bottom edges of the pages. Write your locker number
and lock combination on the inside of the front cover.
2. Your notebook has duplicate numbered pages. Insert the cardboard backing behind each
set of pages before beginning to write. Ensure that the writing appears clearly on the
duplicate copy. These are the copies you will turn in to your instructor. If these are not
legible, turn in the original.
3. Use the first two pages to make a Table of Contents:
Date Experiment #
Title of Experiment
Page #
Grade
4. Start every experiment on a fresh page, using only right-hand pages. Put the date, experiment # and title of experiment at the top. Use clear sub-headings where several experiments are done under one main heading. Write going across the page rather than
down the column.
5. In writing up an experiment, organize it in the following way, unless directed otherwise:
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HOW TO WRITE YOUR LAB REPORT
Prelab (to be completed before experiment and class)
a. Objective: State concisely what the purpose or objective of the experiment is.
b. Background: Write the definition of the terms asked for in the manual.
c. Procedure: Write as briefly as possible what is to be done and the correct order
of doing it. It is acceptable to use an outline and number the steps.
d. Prelab Assignment: Answer all prelab questions when applicable.
In-lab (to be completed while doing the experiment before leaving lab)
e. Data: Every measurement or reading must be recorded in ink when it is made. A
clear label or statement should show exactly what has been measured. Digits
should be arranged in columns vertically under each other. The value of a neat,
permanent record made at the time of observation cannot be overemphasized.
NEVER record laboratory data on any piece of paper other than your notebook.
f. Calculations: Arrange each calculation neatly and in detail on the page (it is best
to first do calculations on scratch paper and then copy them into your notebook).
Always include units and show a prerounded and then final rounded answer. This
page may also be used for any omissions from the write-up, for drawing diagrams
or for any required graphs.
Post-lab (to be completed after doing the experiment before next class)
g. Results: List the results of the calculations. When appropriate, give a table of
summarized results. Whenever possible express the results in tabular form.
Clearly label all results and insure that all numbers are accompanied by
appropriate units.
h. Discussion: Write a statement summarizing your results and observations.
Explain what the results mean! (What is the “take-home lesson”? What significant
conclusions can be reached on the basis of the results?) Comment on the probable
sources of error and limitations of any analytical methods employed. Briefly
comment on the precision and accuracy of your own work and any unusual
observations made.
IMPORTANT: A well thought out and well written discussion shows your understanding of the experiment and your grade will depend on this aspect of the writeup.
i. Answers to all Post lab Exercises and/or Questions when applicable.
Come to lab PREPARED to do the assigned experiment. This requires that you will have read
and studied the experiment at home and prepared your notebook by writing the problem,
theory, outlined procedure, answers prelab questions, and set up tables for recording your
data into. A sample lab report appears below. Examples of how to properly prepare your
notebook will be given for the first few experiments in this book.
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The best time to complete the write-up of an experiment is as soon as possible. When there is
time left, start working on your calculations right in lab- this way you can obtain help. Your
instructor will tell you when the completed report will be due.
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SAMPLE NOTEBOOK ENTRY
OBJECTIVE:
To determine the percent water in copper sulfate pentahydrate by heating.
BACKGROUND:
Empirical formula- the smallest whole number mole ratio of a compound/molecule.
PROCEDURE:
A quantity of copper sulfate hydrate was measured by weighing a test tube with and without the
sample. The sample was then heated with a flame of a Bunsen burner in the manner pictured
until no more weight loss could be detected. A second sample was treated in the same way.
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RESULTS:
% water in salt hydrate
Trial 1
Trial 2
Average
36.0%
36.0%
36.0%
DISCUSSION: The weight percent water in copper sulfate pentahydrate averages 36.0% to
three significant figures. The two trial results were very precise, in fact exactly the same,
suggesting that the % mass of water in a particular salt hydrate is constant and independent of
how much salt hydrate is used. Sources of error include not evaporating all the water, which
would increase the calculated percent of water in the sample.
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What you need to know for Quiz Before Beginning Lab
Significant figures
Common laboratory equipment (page A13)
Safety in the lab (safety quiz) (page A10)
Basic parts of the lab report
Sample Quiz 1
1. Name the lab glassware shown to the right is:
__________________________________
2. How many significant figures are in each of the following numbers?
47 students ____________
0.004506800 _______________
3. Calculate the following to the correct number of significant figures:
2.34 + 1.235 x 3.7 = _________________
4. Answer True/false
Long hair need not be tied up if you are not using a burner____________
Small amounts of acids or base can be poured down the sink _______________
5. Where should the following information be recorded in a students’ notebook? Choose
from the following;
objective, background, procedure, data, calculations, results, discussion
Mass of sodium chloride in test tube: 2.342 g _________________________
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EXPERIMENT 1
MEASUREMENT, VOLUME, AND
DENSITY
PRE-LAB
Objective: Write short statements describing the purpose of each part of the experiment. Since
there are four parts to this experiment, write a separate statement for each part.
Background: Define the terms density, buoyancy force, extensive property, intensive
property, precision, and accuracy.
Procedure: Briefly write/ outline what you will be doing in each part of the experiment. Write
a separate procedure for each part.
Data: Write separate data tables for parts a, b, c, and d. Use the lab manual as your guide for
parts a, b, and c. Put all data tables following each other without leaving space for calculations.
Create a data table for part d by reading through the procedures and noting all measurements you
are asked to record. Create clear labels for each measurement.
OBJECTIVE OF EXPERIMENT 1
To calculate the density of an aluminum cylinder from its mass and volume. The mass will be
measured using an electronic balance (Part A), and the volume will be found using three
different methods: by measuring dimensions (part B), by displacement (part C), and by buoyancy
force (part D).
BACKGROUND OF EXPERIMENT 1
DENSITY
Matter has mass and it occupies space (has a volume). Mass is the amount of matter present in
an object, while volume is the amount of space a substance occupies. These two properties, mass
and volume, are called extensive properties. An extensive property is one that depends upon
the amount of substance present. The more we have of any sample of a substance the larger will
be its mass and volume.
A relationship exists between these two extensive properties and it is called the density. Density
is the ratio of the mass to the volume of a given amount of matter.
Density = mass/volume
Density is an intensive property. An intensive property is a property independent of the amount
of substance present. It is a fundamental characteristic of the substance itself. Any object made
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of a particular pure substance, no matter what its size or shape, has the same density. For
example any aluminum object, a piece of foil or a cylinder, has a density of 2.702 g/cm3 (or
g/mL) and any amount of pure water has a density of 0.997 g/mL, at room temperature. It is only
because the volume of an object changes slightly as temperature changes that the density values
will also change. The volume of matter almost always increases with temperature, so density
decreases as the temperature rises and increases as temperature falls. It is for this reason that any
density value must be reported with a temperature. (Since the change of density with temperature
is small for solids, it may be ignored in this experiment.)
PART A: MEASURING MASS
OBJECTIVE A
To measure the mass of the aluminum cylinder.
BACKGROUND A
Mass is defined as the amount of matter present in an object. The mass of an object is always
constant. Mass is determined using an instrument known as a balance. There are many different
types of balances. In this class you will be using an electronic top loading balance. Like any
mechanical device, the balance may be out of adjustment and so each semester all of the
balances are calibrated (adjusted) with samples of known masses to ensure that they are reading
properly. This calibration process need only be done periodically. An important adjustment that
must be made before every use of the balance is called taring or zeroing. This is done by
simply touching the zero button when the balance is clean and there is nothing positioned on it,
and then waiting until it reads 0.000 g.
Error in Experiment
All measurements have some error associated with them. The extent of this error depends
upon both the measuring devices being used and the measurer herself. To minimize error the
device should be working properly and be used properly. Calculating deviation and average
deviation are one way to determine error when related to a repeating series of a measurement.
They quantitate the precision (see page 11) or the closeness (reproducibility) of a series of
measurements of the same dimension. Another way is to calculate error and percent error.
These measurements can only be calculated when the quantity being determined
experimentally is known, such as the density of aluminum in this experiment. When a value
such as this is known it is called the true value. In this case it is 2.702 g/cm3. Error and
percent error quantitate the accuracy (see page 10) or closeness of a measured (experimental)
value to its true value. Formulas for error and percent error appear on page 27.
PROCEDURE A
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Make three determinations of the mass of your cylinder on the electronic balance and record
them in ink into the data table you have constructed in your notebook (see below). Important:
always wait until the stabilization symbol “ ” appears in the lower left of the display before
recording the mass or adding or removing anything from the balance (read page 14, “Measuring
the mass of solids” before using the balance).
IN-LAB
Data A: Complete the data table for part A.
Calculations A: perform the calculations asked for in part A on the next page.
Repeat this format (data followed by calculations) for the remaining parts of this experiment.
Sample Data Table for part 1A:
Copy this format into your notebook
Data: Mass Determination
Measurement
Deviation from
average
Trial 1
__________
__________
Trial 2
__________
__________
Trial 3
__________
__________
Sum
__________
__________
Average (in g)
__________
__________
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CALCULATIONS A
Calculations: (Do these on scratch paper first and then transfer into your notebook under the
calculations section.
1) Calculate the average mass and the average deviation for the set of measurements and
transfer them into your notebook in the data section (as shown above).
The average is the result obtained when the sum of all measured quantities is divided by the
total number of quantities.
Sum of quantities
Average =
Number of quantities
The deviation is the absolute value of the difference between each value and the average
value.
Deviation= │trial value — average value│
PART B:
CALCULATING VOUME BY MEASURING EXTERNAL
DIMENSIONS
OBJECTIVE B
To measure the external dimensions of an aluminum cylinder and calculate its volume.
BACKGROUND B
A vernier caliper is a device used for measuring the distance between two calibration marks. As
opposed to a ruler where the measurer positions the object and reads the measurements, a caliper
fixes the position of an object between its “jaws”. Therefore the caliper is a more precise
measuring device than the centimeter ruler. After calibrating the caliper to read zero when the
jaws are completely closed, the jaws of the caliper, once tightened around an object, will
accurately define the distance to be measured. The digital read out makes reading the
measurement simple. The digital caliper you will use provides millimeter readings to the nearest
hundredth place (all non-zero numbers on a digital instrument are considered significant).
VOLUME
The volume of a regularly shaped solid object may be calculated from its dimensions if the geometry is not too complicated. It is known that the volume of a cube is the cube of the length of
an edge, and the volume of a rectangular parallelepiped (a box) is the length x width x height.
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These two volumes are illustrated in Figure 1.2, along with two other common shapes, the right
cylinder and the sphere. The abbreviations used in the formulas are s, side; 1, length; w, width; h,
height; d, diameter. π (the Greek letter pi) is the mathematical constant equal to approximately
3.1416.
PROCEDURE B
Remove the aluminum cylinder in your locker and ensure that it is clean and free of tape. What
dimensions must you measure in order to calculate its volume?
In this part of the experiment, you will measure the dimensions of the cylinder with a vernier
caliper. After calibrating your caliper to read zero with the jaws of the device completely closed,
insert your cylinder between the set of jaws on the bottom of the caliper and tighten. Make three
measurements of each dimension of your object with the vernier caliper and record the data in
your data table. Be sure to label the dimensions properly.
Sample Data Table for part 1B:
Copy this format into your notebook
Dimensions to be measured: _height (h)_
Diameter_(d)_
Measurement
Measurement
Trial 1
__________
__________
Trial 2
__________
__________
Trial 3
__________
__________
Sum
__________
__________
__________
__________
Average (in mm)
Average (in cm)
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__________
Page 28
CALCULATIONS B
1) Calculate the average value, in units of mm, for each dimension measured for the caliper
(height and diameter)
2) Convert the average values for height and diameter from units of mm to cm and copy them
into your data table.
3) Calculate the volume of your cylinder using the average of the values (in cm) you determined
from the caliper and highlight your final answer. Use the formula:
π
Volume =
d 2h
4
and show all of your work. Transfer all of these calculations neatly into your notebook.
Here is a sample calculation.
π
V=
d2 h =
4
3.1416
(1.23 cm)2(2.65 cm) = 0.7854 (1.5129cm2)(2.65cm) =
4
3.1488 cm3 = 3.15 cm3
(Show any formulas you are using and all of your steps here including units next to each
measurement. Always round off just your final answer using appropriate significant figures
that rely on your original data- in this case 3 digits!)
4) Calculate the density of your object using the average volume you calculated above and the
average mass you found in part a.
5) The true density for aluminum is 2.702 g/cm3. Calculate the error and percent error for your
experimentally determined densities using the formulas below.
Error = │Experimental Value — True Value │
Error
% Error =
x 100
True Value
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PART C:
MEASURING VOLUME BY DISPLACEMENT
OBJECTIVE C
To measure the volume of a cylinder by direct displacement and by Archimedes’ principle
BACKGROUND C
A solid object submerged in a liquid will displace a volume of liquid equal to the volume of the
object. Displacement is a very common sense concept and explains why you wouldn’t fill your
bathtub up to the very top- it would obviously spill over when you got in it. In fact if you
measured the amount of water that spilled out of your full tub when you carefully submerged
yourself in it, that volume would be equal to the volume of your body! So, this is a very easy
way to determine the volume of an irregular shaped object- like your body or a regular shaped
object like a cylinder. If some liquid is in a graduated cylinder and you drop an object into the
cylinder, the increase in the height of the liquid surface is a direct measure of the volume of the
object.
Volume of cylinder = Volume after submerging cylinder – Volume before submerging cylinder
The method is only as precise as the precision of the graduated cylinder, and a 50 mL cylinder is
usually not precise to more than ±0.5 milliliter. This method, therefore, chiefly serves to get an
estimate of the volume of irregular objects, or to check the arithmetic on calculations of volume
based on the dimensional measurements.
PROCEDURE C
Add some deionized water to a clean 50 mL graduated cylinder and record the initial volume to
the tenth place. Read the volume with your eye level with the meniscus, or curved surface of the
water and record it in ink in your notebook. It is typical to read the bottom surface of the
meniscus (see Figure 1C). Note that the error in the graduated cylinder is ±0.5 mL (which means
that you would record a 0 or a 5 in the tenths place for it to be a proper measurement). Take your
cylinder and drop it gently into the graduated cylinder. Shake any bubbles out and record the
volume of the water with the object submerged. The difference between the final and initial
volumes is the volume of liquid displaced. This is a measure of the volume of the object.
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Figure 1C
Sample Data Table for part 1C:
Copy this format into your notebook
Initial volume in the graduated cylinder: _____________
Final volume in the graduated cylinder: _____________
CALCULATIONS C
1) Calculate the volume of the object to the correct number of significant figures by finding
the amount of water displaced.
2) Calculate the density of your object using the volume you found above and the average
mass of the object from part A.
3) Calculate the error and percent error for your experimentally determined density just like
you did in part B.
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PART D:
MEASURING VOLUME BY ARCHIMEDES’ PRINCIPLE
OBJECTIVE D
To measure the volume of a cylinder by Archimedes’ (buoyancy) principle.
BACKGROUND D
A more precise determination of volume can be made using a balance and Archimedes’
principle. Archimedes’ principle states that when an object is immersed in a liquid, the liquid
exerts an upward force known as the buoyancy force. This buoyancy force is equal to the weight
of the liquid displaced (Figure 1D).
Figure 1D
Imagine that the anchor in the figure above is your cylinder. The buoyancy force can be
calculated by taking the difference between the weight of the cylinder in air and its weight when
suspended fully submerged in the liquid (weightin air – weightin water). If the buoyancy force on the
cylinder is measured, that will tell us the mass of the amount of water displaced by the cylinder.
If the liquid is pure water, its density can be looked up (at the relevant temperature). This allows
us to solve for the cylinder’s volume using the formula d = m/V, rearranged to V=m/d.
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Mass of water displaced = buoyancy force = weightin air – weightin water
Volume of water displaced = Volumecylinder (The water is displaced by the cylinder, therefore
the volume of the water displaced is the
same as the volume of the cylidner.)
Mass of water displaced (buoyancy force)
Density of water displaced (densitywater) =
Volume of water displaced (volumecylinder)
buoyancy force
densitywater =
volumecylinder
buoyancy force
volumecylinder =
densitywater
Equation to find
volume Part D
PROCEDURE D
To measure the buoyancy force on your cylinder you will need to use the coiled wire hanging
from the bottom of the balance. (Do not remove it or pull on it. See your instructor if you are
having difficulties.) Tare (zero) the balance with the wire coil hanging freely. Prepare your 150
mL beaker so that it is about 2/3 filled with deionized water and place it on the jack directly
below the coil on the balance. Place your cylinder in the coil and completely submerge the coilcylinder pair in the beaker of water. Be careful to insure that the coil/cylinder is not touching the
bottom or sides of the beaker. You can use the jack to adjust the height of your beaker. (Is it
necessary to know the weight of the beaker of water for this problem?) Measure the apparent
mass of the submerged object (How should this compare to the same object’s mass in air?)
Record the measurement in your notebook. Return with your beaker and the cylinder to your
bench. Obtain a thermometer and determine the temperature of the water in your beaker to the
nearest tenth oC (one place after decimal) and record it in your data table. Using table 1 below,
look up the density of water at this temperature and record it.
Data Table for part 1D
For this part of the experiment and most future ones you will be creating your own data table.
Use the examples given in the previous parts to help guide you.
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CALCULATIONS D
1) Calculate the volume of the cylinder by using the mass of the water displaced (buoyancy
effect) and the density of the water from Table 1. Hint: Use the “equation to find volume” on
the previous page.
2) Using the volume determined above and the average mass in air from part A, calculate the
density of the cylinder.
3) Calculate error and percent error for your experimentally determined density of your
cylinder.
Table 1: Density of Water (in g/cm3) at Temperatures 18 to 24oC in 0.1oC increments.
0
C
0.0
0.1
0.2
0.3
0.4
0.5
0.6
0.7
0.8
0.9
18
0.99860 0.99858 0.99856 0.99854 0.99852
0.99850 0.99848 0.99846 0.99844 0.99842
19
0.99840 0.99838 0.99836 0.99834 0.99833
0.99831 0.99829 0.99827 0.99824 0.99822
20
0.99820 0.99818 0.99816 0.99814 0.99812
0.99810 0.99808 0.99806 0.99804 0.99801
21
0.99800 0.99797 0.99795 0.99793 0.99790
0.99788 0.99786 0.99784 0.99782 0.99779
22
0.99777 0.99775 0.99772 0.99770 0.99768
0.99766 0.99763 0.99761 0.99758 0.99756
23
0.99754 0.99751 0.99749 0.99747 0.99744
0.99742 0.99739 0.99737 0.99734 0.99732
24
0.99730 0.99727 0.99725 0.99722 0.99720
0.99717 0.99715 0.99712 0.99710 0.99707
POST-LAB
Results: Prepare a summary results table (see below) containing your results from all four parts
of the experiment.
Mass
Part A
Part B
——–
Part C
——–
Part D
——–
Volume
Density
Error
% Error
——–
——–
——–
——–
Discussion: Write a thoughtful discussion that analyzes the results and addresses these
questions.
a. How do the different values for the volume of your cylinder compare? Which do you
think is the most accurate? Why?
b. How do the different values for density compare? Which do you think is the most
accurate and why?
c. What were some experimental sources of error that could have led to an increase or a
decrease in the calculated density? Note that carelessness and mistakes are not sources
of error. Errors are those that a conscientious worker cannot avoid with the available
measuring devices.
(This section should show whether or not you understood the experiment.)
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What you need to know for Experiment 1 Quiz
Reading the volume of liquid in a graduated cylinder
Extensive/intensive properties
Calculating density/ error/% error/ average
How different sources of error would affect calculated density
Sample Quiz 2
1. The following is a graduated cylinder which has an error of ±0.5 mL:
What is the volume of liquid shown? __________________
mL
30
20
10
2. State whether each of the following is an extensive or intensive property:
Temperature ____________________
mass _________________________
3. Find the average of the three trials: 5.683 g, 5.681 g, 5.688 g
_________________________
4. What is the error in density for a sample if the experimental density was 2.13 g/cm3 and
the true value is 2.437 g/cm3? _________________________
5. A student sets out to determine the density of a metal sample just like you did in
experiment 1. To measure the volume, he immersed his sample in a graduated cylinder
containing water and measured the volume of the displaced liquid. If some water
splashed out when he placed his sample in the cylinder, how would his calculated density
compare to the true density of his sample. Choose from higher, lower, no change, can’t
tell.
___________________________
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Experiment 2
Preparation and Measurement of
the Density of Solutions
PRE-LAB
Objective: Write a short statement describing the purpose of the experiment.
Background: Define the terms solution, solvent, solute, and aqueous solution.
Procedure: Briefly write/outline what you will be doing in the experiment.
Prelab Calculations: Perform the two calculations asked for on page 40.
Data: Set up the data table for all the information you will be collecting in this experiment.
OBJECITVE
To prepare and measure the density of a 10.0 percent composition (by mass) salt solution
and a 40.0 percent composition (by volume) alcohol solution.
BACKGROUND
A solution is a homogeneous mixture made up of two or more pure substances that has all of
these components thoroughly mixed at the molecular level. The component present in the largest
amount is called the solvent and the component(s) present in the smaller amount(s) is known as
the solute(s). When solids are dispersed in a liquid the solids are usually the solutes and the
liquid is the solvent. Many of the solutions we will work with use water as the solvent and are
called aqueous solutions. They can be represented by writing the formula of the solute followed
by (aq). For example, sodium chloride dissolved in water would be written NaCl(aq). When the
solvent of a solution is not specified, the solvent is presumed to be water.
Another property of a solution is that its components can be mixed in varying amounts. This
creates solutions of different concentration. Addition of a lot of solute to a fixed amount of
solvent gives a more concentrated solution while a small amount of solute gives a more dilute
solution. The composition of the mixture may be quantitated by specifying the ratio of the mass
of one component to the total mass of the solution (% composition by mass or % m/m). The
composition may also be quantitated by the ratio of the volume of one of the components to the
sum of the volumes (% composition by volume or % v/v).
As we learned in experiment 1, the density of a pure substance is a physical property that is
independent of the amount of the substance we have. Each pure substance has a characteristic
density value. For example, we saw that the density of water is approximately 1.0 g/mL
(remember that density varies slightly with temperature). When substances with a density higher
than 1.0 g/mL are added to water the density of the resultant solution is greater than 1.0 g/mL.
Conversely, adding substances with a density lower than 1.0 g/mL to water decreases the density
of the resultant solution to less than 1.0 g/mL. In this experiment we will prepare aqueous
solutions containing either sodium chloride or ethanol as the solute and then measure their
density using the buoyancy concept of Archimedes’ principle.
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To prepare a given solution we must know the pure substances to be used and how much of the
solution to prepare. For example let’s say we want to make 100.0 g of a 10.0% (by mass) salt
solution [NaCl(aq)]. Masses are always additive, so the sum of the masses of the components is
the mass of the solution. Since % represents parts out of 100, in the case of a 10.0% solution,
10.0 mass parts out of every 100.0 mass parts of solution is solute (salt) and the rest, 90.0 mass
parts, is water. The following statement can be made for a 10.0% NaCl(aq):
10.0 g salt ≡ 90.0 g water ≡ 100.0 g solution,
where ≡ means equivalent to. Conversion factors can be constructed from any combination of
two of these terms (see example 1 on the next page).
In general we cannot say that volumes of different liquids are additive. Therefore volumes of
unlike substances do not give a volume of solution which equals the sum of the volumes of the
components. A 40.0 % (by volume) alcohol solution means 40.0 volume parts alcohol and the
rest, 60.0 volume parts, is water. The volume of the solution is not necessarily 100 volume parts.
The following statement can be made for a 40.0% alcohol solution:
40.0 mL alcohol ≡ 60.0 mL water
Conversion factors can be constructed from any combination of these terms (see example 2 on
the next page).
Percent composition, like density, is an intensive property since the amount of the solution
doesn’t affect its % composition.
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In order to prepare a solution one must first decide what quantity of solution is needed, and then
compute the quantities of components to be measured. Below are examples of how calculations
could be made for preparing solutions of each type of % composition- mass and volume.
EXAMPLE 1: % composition by mass
Prepare 80.0 g of a 10.0% (by mass) salt solution (NaCl(aq)). How much salt and how much
water are needed?
Given:
g of salt solution
Want:
g of salt
Road map:
g of salt solution
g of salt
Use conversion factor:
10.0 g salt/100 g solution
10.0 g salt
Calculation:
80.0 g of solution X
100 g solution
= 8.00 g salt
The remainder is water: 80.0 g solution – 8.00 g salt = 72.0 g water
There is no way to prepare a specified volume of a solution by adding volumes of unlike
substances, like we did above with masses, unless the densities of the solution and the
components are known. The volume of one of the components is specified and the volume of the
other is computed from its relationship to the one specified.
EXAMPLE 2: % composition by volume
Prepare a 40.0% (by volume) alcohol solution (C2H6O(aq)) containing 25.0 mL of alcohol. How
much alcohol and how much water are needed?
Given:
mL of alcohol solution
Want:
mL of water
Road map:
mL of alcohol solution
mL of water
Use conversion factor:
60.0 mL water/40.0 mL alcohol
60.0 mL water
Calculation: 25.0 mL of alcohol solution X
= 37.5 mL water
40.0 mL alcohol solution
After measuring out 37.5 mL of water in your cylinder, you would add 25.0 mL of alcohol to it.
The densities of the solutions could be found by measuring the mass and the volume of each
solution. If by chance an accurate measuring device for volume is not available as is the case
here, we can use the buoyancy law to calculate the density. Remember that the buoyancy law
says that the difference between an object’s weight in air and its weight when it is submerged in a
liquid is the weight of the liquid displaced.
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We have an aluminum cylinder whose accurate volume can be calculated from its actual density
value of 2.702 g/cm3 and its mass in air. If we know the mass of solution that the cylinder
displaces (which is the buoyancy force) we can calculate the solution’s density.
weight of the solution displaced by cylinder (buoyancy force)
d soln =
volume of solution displaced by cylinder (volume of cylinder)
Note from experiment 1: buoyancy force = mass in air – mass in solution
PROCEDURE
Salt Solution: Prepare 100.0 grams of a 10.0 percent, by mass, salt solution. Put on your
protective eyewear. (You and your lab partner may prepare one solution to share between you.)
Tare a 150 mL beaker on the balance- this means putting your dry beaker on the balance and
pressing the tare knob. Wait for the balance to read 0.000 g. Remove the beaker and add sodium
chloride to the beaker until you reach the calculated amount (remember to remove the beaker
from the balance when adding any chemical- the balance will retain its tared memory). Once
you get close to 10.0 g on the balance, (its not essential to get 10.000 g since you only need 3
significant figures), return with your beaker and salt to your bench. So that you don’t have to
weigh the mass of the water on the balance also, assume that water has a density of 1.0 g/mL,
and measure the calculated mass of deionized water using your graduated cylinder . Pour it into
the same beaker. Stir the mixture until the salt is completely dissolved and this will give you a
10.0% by mass salt solution Label the beaker as NaCl solution and go on to prepare the solution
in Part B.
Alcohol Solution: Prepare a 40.0 percent, by volume, alcohol solution containing 30.0 milliliters
of alcohol. (You and your lab partner may prepare one solution to share between you).
The volumes of alcohol and deionized water may be measured separately in a graduated
cylinder. (The alcohol solution can be found in a hood.) Mix the liquids in your 150 mL beaker
and label it as alcohol solution.
Once you and your partner have made the solutions, take both solutions, each of your aluminum
cylinders, each of your lab notebooks, and pens to a free balance. Tare the balance and place
your aluminum cylinder in the coiled wire below the balance (gently dry the wire with a paper
towel if it is wet). Wait for the balance to stabilize and record its mass in your notebook. This is
the mass of your cylinder in air. Now place your beaker with NaCl solution under the balance
and completely submerge the coil-cylinder pair in the beaker of NaCl solution. Be careful to
insure that the coil is not touching the bottom or sides of the beaker. After the balance reading
stabilizes, record it in your notebook. This is the mass of your cylinder in NaCl solution.
Replace the beaker containing NaCl solution with the beaker containing the alcohol solution and
determine the submerged mass of your cylinder again. Record it in your notebook. Allow your
lab partner to repeat the same three measurements using their own cylinder. Using the
determined masses, complete calculations 1-5 below. You may use your notes from the
calculations in experiment 1 part D as a guide to help you calculate the buoyancy effect on your
cylinder when submerged in the solution.
Disposal: The salt solution can be poured down the drain and the alcohol solution should be
poured into the waste container in the hood marked “alcohol solution”.
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PRELAB CALCULATIONS
Carefully read the background section of this experiment and use the examples to guide you
through the following two calculations.
1) What mass of salt and what mass of water are needed to prepare 100.0 grams of 10.0
percent by mass salt solution?
Hint: remember that a 10.0 percent by mass salt solution = 10.0 g salt / 100.0 grams of
solution. Use this mass percent as a conversion factor for the following solution map:
mass solution
mass salt
Remember, the solution contains both water and salt!
2) What volume of water must be added to 30.0 milliliters of alcohol to obtain a 40.0
percent (by volume) solution of alcohol?
To get started, write the percent by volume as a conversion factor between volume of
alcohol and volume of water, as shown in the background section.
IN-LAB
Data: Complete the data table you prepared.
Calculations: Perform the calculations asked for below.
CALCULATIONS
1) Calculate the volume of the aluminum cylinder you used from the actual density of aluminum
(2.702 g/mL) and its mass in air.
2) a) Calculate the buoyancy force on the cylinder due to the salt solution.
b) Determine the experimental density of the salt solution from the buoyancy force and the
volume of the aluminum cylinder (determined in 1) above).
3) According to the CRC Handbook of Chemistry and Physics the actual density of a 10.0% salt
solution at 25°C is 1.0726 g/mL. Calculate the error and % error of your calculated density for
the salt solution.
4) a) Calculate the buoyancy force on the cylinder due to the alcohol solution.
b) Determine the experimental density of the alcohol solution from the buoyancy force and
the volume of the aluminum cylinder (determined in 1) above).
5) According to the CRC Handbook of Chemistry and Physics the actual density of 40.0% (v/v)
alcohol solution at 25°C is 0.926 g/mL. Calculate the error and % error of your calculated
density for the alcohol solution.
Note: You should be able to get errors of less than 10% for both solutions in this experiment. If your
values are higher, repeat the measurement(s) for the mass of the cylinder in the solution(s) and
recalculate.
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POST-LAB
Results: Create a summary results table
Conclusion: Analyze the results and errors in a written discussion. Include any sources of
error that could have led to an increase or a decrease in the calculated density values for the
solutions.
What you need to know for Experiment 2 Quiz
Definition of the terms solution, solvent, solute, and aqueous solution.
Calculating mass (or volume) of solute/solvent needed given percent composition.
Determining density from buoyancy force and volume of cylinder
Sample Quiz 3
1. Fill in the blank with one of the following terms: solution, solvent, solute, aqueous
solution.
The component present in the smaller amount is the _________________________
2. How many mL of alcohol and how many mL of water are needed to prepare a 35.0%
alcohol solution containing 15.0 mL alcohol.
3. Use the following data to answer the questions below: A salt solution was prepared and a
2.45 cm3 aluminum cylinder was submerged in the solution. The mass of the cylinder in
air was 7.325 g and its submerged mass was 3.768 g.
a. What is the buoyancy force on the cylinder? ___________________
b. What is the density of a salt solution? Include units _______________________
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Experiment 3
Solubility and Chemical Change
PRE-LAB
For this experiment, prepare your notebook in the following manner:
On the first page:
Objective: Write a brief summary of the three types of problems that you will address in this
experiment.
Background: Write brief definitions for each of the following terms that we will be using in
this experiment: solution, miscible, immiscible, soluble, insoluble, homogeneous mixture,
heterogeneous mixture, reaction, precipitate, exothermic, and endothermic.
On the next two pages:
Turn your notebook horizontally and divide it into three equally sized columns. Label the
columns with the following headings: procedures, observations and conclusions. Write a brief
summary of the procedure in your lab notebook for each experimental part in the procedure
column before you come to lab. You will enter observations and conclusions in the other
columns when you do the experiment. Spread this table over two pages so that you have plenty
of room to write. The terms you defined in the background section will be used in the conclusion
section.
OBJECTIVE
1. Determine the miscibility of different liquids and their relative densities.
2. Determine whether or not a solid will dissolve in water.
3. Determine whether a chemical reaction occurs when two solutions are mixed.
BACKGROUND
Part 1: Mixing two different liquids.
When two different liquids are mixed they may form a homogeneous mixture or solution. These
liquids are said to be miscible resulting in one visible layer or phase. Some liquids are miscible
in all proportions while others are only partly soluble in each other. When two different liquids
do not mix (or mix very little) with each other they will form two distinct layers or phases and
are said to be immiscible. The phase having the lower density will always float on top of the
other. When the two immiscible liquids are both colorless, it may be difficult to determine
which phase is which liquid. One can add a few drops more of one liquid and see which phase
gets larger in size.
Part 2: Mixing a solid in water.
Some ionic compounds will dissolve in water and others will not. When a small amount of a
solid is added to and dissolves in a particular liquid it forms a homogeneous single phase and is
said to be soluble in the liquid. If the solid is not soluble, or insoluble, it will form a cloudy
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heterogeneous mixture from which solid particles will separate out, forming two phases. All
mixtures should be mixed well before determining whether all of the solid can indeed dissolve.
Solubility tables for ionic compounds in water can be used to confirm your results.
Part 3: Is it a reaction or not?
If two different aqueous solutions are mixed several different observations can be made. A
number of these observations are evidence that a new chemical(s) has been formed. This is
called a chemical change or reaction. One way in which a reaction can be confirmed is when the
mixing of two solutions produces an insoluble solid product. This results in a cloudy,
heterogeneous mixture and the solid formed is called a precipitate. A solubility chart (see Table
3.1 on page 44) lists ionic compounds that are soluble (dissolve in water) and insoluble (do not
dissolve in water).
Other chemical reactions result in a gaseous product being formed that escapes as bubbles from
the solution. There are also some reactions in which dissolved components combine to produce
colored substances. These reactions can be recognized by a color change. Sometimes no
observed change takes place, generally indicating that no reaction has occurred.
If there is a temperature change, you can describe the processes or reactions that occur in any of
the parts above using the terms exothermic or endothermic. The term exothermic describes a
process or reaction that releases energy, usually in the form of heat; while the term endothermic
describes processes or reactions that absorb heat energy.
IN-LAB
Fill in the observations and conclusions columns for each part by answering the questions asked
for under the procedures section. Remember to use the terms you defined in the background
section. The answers to the prompting questions in the procedure should go in the conclusions
section.
Observations vs. Conclusions:
Part of what you will learn in this experiment is how to distinguish between writing
observations and conclusions. Observations are descriptions of what you observe, that is see,
feel, smell, or hear. Your observations should be written in simple, non-scientific language.
Conclusions are statements describing what you are deducing or inferring from your observations.
Here you should use the appropriate scientific language, including the new terms you are
introduced to in this experiment.
For example:
Observation 1: after mixing the chemicals the test tube became warm
Conclusion 1: the process occurring gave off heat and was therefore exothermic
Observation 2: after mixing the 2 liquids well, two separate layers (phases) were visible
Conclusion 2: these two liquids formed a heterogeneous mixture and are therefore
immiscible
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PROCEDURE
Wear your protective eyewear throughout the experiment.
EXPERIMENTAL TIPS
Use of Reagent Droppers Bottles: To avoid contamination when dispensing chemicals using
droppers, never touch the end of the dropper to a container. This may contaminate the end of the dropper.
Quantities: This is a qualitative experiment, meaning that the amounts of the chemicals you use are
approximate and need not be measured carefully. For all parts of this experiment, make sure that you mix
the chemicals well before making your observations
Disposal of Chemicals: You will be using a variety of different chemicals in this experiment.
Years ago chemicals of all sorts were disposed of down the drain or in the trash, without much concern
about their affect on us or the environment. Today this is not the case- there is much more public
consciousness of our safety and the hazards of chemicals in the environment are regulated by many laws
and statutes. Most chemicals must be collected after use and disposed of in the proper means. Learning
about the proper disposal of chemicals is an important part of this laboratory experience. This lab book
will indicate how to dispose of used chemicals so follow them carefully. If you still have questions, ask
your instructor.
Part 1. Mixing Different Liquids:
a. Is hexane (C6H14) miscible (does it mix) with water?
Place about a half inch of deionized water from your wash bottle into a small test tube and
add about a quarter of an inch of hexane (found in the hood). Mix well. What do you
observe? Did the water end up on the top or the bottom? (Since you added more water than
hexane, the layer with the greater volume is the water.) If you can’t tell, add some more
water and see into which layer it goes.
Empty the liquids into the red container in the hood marked “organic waste”.
b. Is methanol (CH3OH) miscible with water?
Place about a half inch of deionized water in a test tube and add about a quarter of an inch
of methanol. Mix as before. What do you observe? Is there any way of telling which has
the greater density? Try dripping some water down the side of the tube and see if it layers
on top or sinks to the bottom to determine relative densities.
Empty the liquids into the red container in the hood marked “organic waste”.
c. Is methanol (CH3OH) miscible with hexane (C6H14)?
Place about a half inch of methanol in a test tube and add about a quarter of an inch of
hexane. Mix as before. Let it stand for a few minutes. Are there two phases? Which phase
is which chemical?
Empty the liquids into the red container in the hood marked “organic waste”.
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Part 2: Dissolving a Solid in Water:
a. Is solid barium chloride (BaCl2) soluble in water?
Place a very small amount (about the size of a match head) of barium chloride dihydrate in
a small, dry test tube and add about an inch of deionized water from your wash bottle. Mix
well. Did the solid dissolve or not? Save this for later use.
b. Is sodium sulfate (Na2SO4) solid soluble in water?
Place a very small amount (about the size of a match head) of sodium sulfate in a small,
dry test tube and add about an inch of deionized water. Mix as before. Did the solid
dissolve or not? Save this for later use.
c. Is barium sulfate (BaSO4) solid soluble in water?
Place a small amount of barium sulfate in a small, dry test tube and add about an inch of
deionized water. Mix as before. Is the mixture cloudy? Did the barium sulfate dissolve?
d. Is there a reaction between barium chloride (BaCl2) and sodium sulfate (Na2SO4)?
Now mix the barium chloride solution from (a) and the sodium sulfate solution from (b)
together. Let it stand for a few minutes before making observations. Does this mixture look
like any of the mixtures of solid and water that you just prepared? What do you conclude?
Did a chemical change occur? If so, can you suggest what one of the products might be?
Empty all solutions from part 2 into the container in the hood marked “aqueous waste.”
Part 3: Mixing Aqueous Solutions: (Carry out these procedures in a clean spot plate.)
a. Is there a reaction between potassium nitrate and sodium chloride solutions?
Add a few drops of potassium nitrate (KNO3) solution into a clean well of your spot plate.
Add an equal amount of sodium chloride (NaCl) solution and mix using your stirring rod.
Does any visible reaction occur? Record your observations.
b. Is there a reaction between iron(III) chloride and potassium thiocyanate solutions?
Into another clean well on your spot plate, add a few drops of iron(III) chloride (FeCl3)
solution and potassium thiocyanate (KSCN) solution and mix. What happens?
c. Is there a reaction between iron(III) chloride and sodium hydroxide solutions?
Place several drops of iron(III) chloride solution into a clean well and then add sodium
hydroxide solution drop by drop. Mix and observe what happens?
d. Is there a reaction between sodium carbonate solution and hydrochloric acid?
Place several drops of sodium carbonate (Na2CO3) solution into a clean well and then add
dilute hydrochloric acid [HCl(aq)] found on the bench top drop by drop, with mixing.
Watch the solution carefully as each drop of acid is added and as you mix. What happens?
Is there a chemical reaction? What type of product (physical state) is formed?
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e. Is there a reaction between copper(II) sulfate and ammonia solutions?
Add a small amount of the dilute ammonia (NH3) solution found on the bench top drop by
drop, while mixing, to a well containing copper(II) sulfate (CuSO4) solution. What
happens? Is there a chemical reaction?
Repeat any reactions that you want to observe again.
Empty all solutions from your spot plate into the container in the hood marked “aqueous
waste.”
POST-LAB
Write a discussion analyzing the results you obtained in each part. Include the following:
For Part 1:
Indicate the relative densities of the 3 liquids: water, hexane and methanol as
determined from the experiment.
Look up the actual densities of these 3 liquids (in g/mL) by “googling” them on the
Internet and write them in your report (write the reference site’s URL in each case).
Use them to confirm or deny the results you obtained and discuss this.
For Part 2:
Compare the solubility results you obtained in this experiment with solubility table 3.1.
Your instructor will review this with you.
Compounds containing these ions are soluble
Exceptions making compound insoluble
Li+, Na+, K+, NH4+, NO3 , C2H3O2
No exceptions
Cl , Br , I
Any of these ions paired with Ag+, Hg22+, Pb2+
SO42
SO42 paired with Sr2+, Ba2+, Pb2+, Ca2+
Compounds containing these ions are insoluble
Exceptions making compound soluble
OH , S2 , CO32 , PO43
S2 or OH- paired with Sr2+, Ba2+, or Ca2+
Any of these ions with Li+, Na+, K+, NH4+
Table 3.1: Solubility Chart
For Part 3:
Summarize your results. What was the specific evidence that suggested that a chemical
reaction had occurred in each case?
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What you need to know for Experiment 3 Quiz
Definition of the terms solution, miscible, immiscible, soluble, insoluble, homogeneous mixture,
heterogeneous mixture, reaction, precipitate, exothermic, and endothermic and applying them to
given scenarios.
Use of the solubility chart
Sample Quiz 4
Choose the correct term(s)
1. Answer the questions below based on the following: A student mixes water and hexane and
observes two phases forming.
a. Based on his results the two liquids are
i. miscible
ii. immiscible
iii. soluble
iv. insoluble
b. the two substances form a
i. homogeneous mixture
j. heterogeneous mixture
2. A student mixed two substances together in a test tube and observed that the test tube
became warm. The process that occurred inside of the test tube was
a. Exothermic
b. Endothermic
3. Based on the solubility chart below, determine whether each of the following substances
will be soluble or insoluble in water:
a. (NH4)2CO3
b. AgCl
Compounds containing these ions are soluble
Exceptions making compound insoluble
Li+, Na+, K+, NH4+, NO3 , C2H3O2
No exceptions
Cl , Br , I
Any of these ions paired with Ag+, Hg22+, Pb2+
SO42
SO42 paired with Sr2+, Ba2+, Pb2+, Ca2+
Compounds containing these ions are insoluble
Exceptions making compound soluble
OH , S2 , CO32 , PO43
S2 or OH- paired with Sr2+, Ba2+, or Ca2+
Any of these ions with Li+, Na+, K+, NH4+
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Experiment 4
Separation of Components
of Mixtures
PRE-LAB
Objective: Write a short statement describing the purpose of the experiment.
Background: Define the terms filtration, extraction, crystallization, and evaporation.
Procedure: Briefly write/outline what you will be doing in the experiment.
Data: Set up a data table for this experiment.
OBJECTIVE
To separate a two component heterogeneous mixture based upon the differing properties of its
components.
BACKGROUND
Most of the materials we encounter in everyday life such as cement, vegetable soup, natural
waters, wood, pizza, and soil are mixtures. A mixture is a blend of two or more different
substances. Sometimes the blending is so uniform and complete that you can’t see any
differences in the composition of the mixture, even with a microscope. These are called
homogeneous mixtures or solutions. They are composed of a solvent, the component found in the
largest quantity, and one or more solutes, those components found in smaller quantities. Some
other mixtures are obviously made of different substances because you can see variations in the
properties from one region to another. These are called heterogeneous mixtures. Each separate
region where the properties are uniform is called a phase. A heterogeneous mixture therefore has
more than one phase present.
A mixture must be heterogeneous at the time a separation of the components is attempted. If the
phases are solid and large enough one could separate them with tweezers. When panning for
gold, we use differences in density to separate gold from the gravel. Iron oxide is separated from
rock by using the magnetic properties of iron oxide. In all of these cases we use differences in
physical properties to separate different phases from each other. In the case of homogeneous
mixtures, we must change the conditions so as to create two phases so they can be separated
from each other.
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A large part of chemical research and industrial production is devoted to finding new ways to
separate minute quantities of substances from each other. Some of these methods are described
as follows:
Filtration: A solid may be separated from a liquid or a gas by straining the fluid through a
porous medium of some kind. If the pores are smaller than the solid phase the liquid passes
through and the solid stays behind. The porous medium is usually paper.
Extraction: Soluble substances may be separated from those that are less soluble by dissolving
them in a solvent. The solution can then be removed from the insoluble components by filtration
or other appropriate means.
Crystallization: A solution of a solid substance may become supersaturated by evaporation of
the solvent or by lowering the temperature. Under the right conditions crystals of the pure solute
may form, which then can be removed by filtration from the impure solution. Some substances
may be purified by crystallization from a solution, because impurities tend to be excluded from a
growing crystal.
Evaporation: Is used as a means of purifying a soluble solid from a liquid by evaporating away
the liquid.
In this experiment we consider the problem of separating sand from a salt. Two phases are
present. (Note: there is only one state, the solid state, but there are two distinct phases). We know
copper(II) sulfate is soluble in water and that sand is insoluble. These facts suggest that we might
extract the salt from the sand with water and filter the sand from the mixture. Since copper(II)
sulfate does not evaporate, we could recover the salt from the solution by evaporation of the
solvent water and by crystallization of the salt.
CuSO4 • 5H2O(s)
+ sand(s)
Add water
(Extraction)
CuSO4 (aq)
+ sand (in water)
(Filtration)
CuSO4 (aq)
(Evaporation/
crystallization)
CuSO4 • 5H2O(s)
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Sand (wet)
Air dry
Sand(s)
Page 49
We may then determine the mass of each component and calculate the mass percent composition
and percent recovery:
mass of pure component A or B
Percent Composition of A or B =
X 100
mass of mixture started with
The above calculation needs to be done once for the salt and once for the sand
sum of mass of both components recovered
Percent recovery =
X
100
mass of mixture started with
PROCEDURE
Wearing your protective eyewear, tare your 150 mL beaker and add a scoop full of the CuSO4 •
5H2O /sand mixture to it. Record the mass in your notebook. Add 10-15 mL of deionized water
to the beaker and swirl the mixture gently. Weigh and record the mass of a piece of clean filter
paper. Assemble the filtering apparatus as follows: fold the piece of weighed filter paper in
quarters. Open the folded paper between the folded edges so as to make a cone (your instructor
will demonstrate). Place the cone-shaped filter in the funnel and wet it with deionized water from
your wash bottle. Press out any air pockets between the filter and the funnel. Support your
funnel with a ring attached to a support rod (rings can be found in the cabinet under the sink).
Record the mass of your clean, dry evaporating dish. Place it beneath your funnel and lower
your funnel so that its stem is touching the inside of the evaporating dish. Slowly pour the
contents of the beaker onto the filter. Take care that the level of liquid is always below the top
rim of the filter paper. Use your wash bottle and a rubber policeman (attached to your stirring
rod) to transfer all of the sand into the filter. Be very careful to not disturb the filter paper with
your stirring rod as this may tear the filter paper. After all the liquid has drained through the
filter, wash the filter with small portions of deionized water until the washings are colorless.
Collect all the filtrate and washings in the evaporating dish.
Remove the filter paper from the funnel, open the filter paper containing the sand on a watch
glass, write your name on a piece of tape and stick it to the watch glass. Place it in the oven or set
it on the shelf assigned to your instructor. Once completely dry, record the mass of the filter
paper and sample when they are dry (this may be the next period depending on what your
instructor decides).
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filtrate
250 mL beaker
Prepare a steam bath by placing a 250 mL beaker filled about 2/3 full of tap water, on a wire
screen on a 3” ring attached to a support rod. Protect the water bath by placing a large ring on
top of the 3” ring. Insure that the rod and ring are tightly clamped and that your Bunsen burner
flame is at the proper height. Carefully place the 250 mL beaker on the wire screen. Place the
evaporating dish carefully on the beaker and heat the water to boiling. Heat the steam bath until
the filtrate has completely evaporated. Do not let the steam bath boil to dryness- add more water
if necessary. Once all the water has evaporated from the evaporating dish, let the set up cool
completely and then carefully remove your evaporating dish. Once it has cooled to room
temperature, dry the bottom of your evaporating dish containing the sample and determine the
mass.
Disposal: After confirming that your calculations are accurate, the CuSO4 • 5H2O solid can be
disposed of in the labeled container and the sand is disposed in the trash.
IN-LAB
Data: Fill out the data table you prepared.
Calculations: Perform the calculations asked for below.
CALCULATIONS
1) Calculate the mass of the:
a) CuSO4 • 5H2O recovered. Remember to subtract the mass of the empty evaporating
dish.
b) Sand recovered. Remember to subtract the mass of the clean filter paper.
2) Calculate the % composition, by mass, of:
a) CuSO4 • 5H2O in the recovered mixture.
b) Sand in the recovered mixture.
3) Calculate the % recovery of the original mixture.
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POST-LAB
Results: Create a summary results table
Conclusion: Analyze the results and errors in a written discussion. What could be the sources
of error and how could these sources of error have affected your calculated percent recovery?
What you need to know for Experiment 4 Quiz
Definition of the techniques filtration, extraction, crystallization, and evaporation.
Calculating percent composition and percent recovery
Sample Quiz 5
1. Fill in the blank with one of the following terms: filtration, extraction, crystallization, and
evaporation
A technique used to separate a soluble substance from a less soluble one by dissolving in
a solvent ____________________
2. Answer the questions below using the following Information: A student attempts to
separate 2.456 g of a sand/salt mixture just like you did in this lab. After carrying out the
experiment, she recovers 1.376 g of sand and 0.658 g of salt.
a. What was the percent composition of sand in the mixture according to the student’s
data?
________________________
b. What was the percent recovery? _______________________
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EXPERIMENT 5
DISTINCTIVE PROPERTIES
OF CHEMICALS
PRE-LAB
Objective: Write a short statement describing the purpose of the experiment.
Background: Define the terms pure substance, mixture, element, compound, ionic compound,
covalent compound, hydrates, homogeneous mixture, heterogeneous mixtures, transparent,
translucent, opaque, malleable, ductile, and brittle. You may need to refer to your textbook or the
internet for the definition of some of these terms.
Procedure: Briefly write /outline what you will be doing in the experiment.
Pre-lab assignment: Complete the first column of the chart on pages 60-65 by filling in
either the name or chemical formula for each sample. Make sure to include the state of matter
when writing a chemical formula. For all pure substances, put a check mark in the correct box to
classify as metal, nonmetal, ionic, or covalent. Each sample only has one classification.
OBJECTIVE
This display of chemicals is designed to familiarize you with the identification of samples of
elements, compounds, and mixtures by observing their properties and different names and
formulas. You will practice making careful observations and learning how to categorize
different types of matter. The chemicals have been organized into nine related groups or series.
The title of each series indicates what those substances have in common. Compare and contrast
the appearance of the substances within each series to see how related chemicals can have
different appearances.
BACKGROUND
All matter is made up of chemicals. Matter can be classified into various groupings (see
Classification of Matter flowchart on page 57). The first categorization is based upon whether
the matter is a pure substance or a mixture. Pure substances are composed of only one type of
chemical and can be elements or compounds. Elements are pure substances made up of atoms
that are the same (they have not truly the same, as isotopes have different numbers of neutrons,
but they are considered the same element). They can be further subdivided into metals or nonmetals. Compounds are pure substances composed of two or more different elements. They
can be categorized in different ways depending on the criteria used. If the criterion is the
presence or absence of ions, then compounds are divided into ionic (those that contain ions) or
covalent (absence of ions and made up only of non-metals). If the presence of C and H is used
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as the criterion, the compounds are divided into organic (contains C, H, and possibly other nonmetals) or inorganic compounds. A mixture is composed of two or more pure substances and it
can be homogeneous, look the same throughout (also called a solution), or heterogeneous, look
visually different. Solutions are comprised of a solvent and one or more solute(s). The solvent is
the component of the solution in the largest amount, while the solute(s) is (are) those present in
smaller amounts.
The basic unit of a sample of matter, which relates only to a pure substance, is the smallest
physical unit in which that substance exists. The basic unit of an element is an atom or, in a few
cases, a molecule. The basic unit of an ionic compound is a formula unit, and that of a covalent
compound is a molecule. Mixtures do not have a basic unit, as they are not pure. Hydrates
(also known as salt hydrates) are crystalline solids that contain chemically bonded water in
definite proportions. They are considered pure and are true ionic compounds. They will have
the symbolism “ •• xH2O” at the end of the formula, where x usually represents a small whole
number. You will find more detailed definitions of these various terms and how to distinguish
between them in your textbook.
PROCEDURE
Using the tables included at the end of this experiment, write the chemical formula and name of
each chemical for a particular series and then complete the remaining columns. Looking
carefully at the names and formulas of the chemicals will help you with the classifications. Skip
a row and go onto another series. You need not go in number order. Use the following
guidelines to fill in each column. Open each container carefully and smell the contents by
wafting a sample with your hand to your nose. Close them tightly when done. Do not open the
other containers of the other series. Use the Classification of Matter flowchart on page 57
Formula and Name of Chemical: Record the chemical symbol or formula and the complete
name of the substance. Be careful to write the formulas correctly.
Physical Properties: State the physical properties of the substance. Indicate the physical state
(solid, liquid, or gas) of the substance. Specify whether the substance is transparent
(transmitting light so images can be clearly seen), translucent (allowing light but not detailed
images to go through), or opaque (not transmitting any light). Do not mistake the container for
the substance. If the substance is solid, report whether it is lustrous, like a metal, or not shiny
(dull). Deduce from the shape and texture whether the substance can be drawn into a wire
(ductile) or a sheet (malleable), or whether it is brittle (a substance that breaks when
hammered), crystalline, or powdery. What is its color, if any? For Series 8, Organic Chemicals,
carefully open each container and, by waving your hand over the top, see if you detect an odor.
Describe it. Close the container firmly. Do some of the properties of the pure substance, such as
its color, also appear in the solution? Do some substances change color in solution?
Pure Substance or Mixture: From the label on the display identify each substance as a pure
substance or a mixture.
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What Type of Pure Substance or Mixture: If it is a pure substance, identify it as an element or
a compound.
What Type of Element: If it is an element, identify it as either a nonmetal or a metal, depending
on what side of the “stairstep” in the periodic table they appear.
What Type of Compound: If it is a compound, identify it as an ionic or a covalent compound.
What type of Mixture: If it is a mixture, indicate whether it is homogeneous or heterogeneous.
Basic Unit: If it is a pure substance, indicate its basic unit. (Note: mixtures do not have a basic
unit.)
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CHEMICAL DISPLAYS
Series 1. IRON
Series 2. COPPER
Fe Iron (powder or wire)
Cu copper (powder, sheet or wire)
Fe2O3 Iron(III) oxide
Cu2O copper(I) oxide
Fe(NO3)3 • 9H2O Iron(III) nitrate
nonahydrate
CuO copper(II) oxide
Fe(NO3)3(aq) Iron(III) nitrate solution
CuSO4 • 5H2O copper(II) sulfate pentahydrate
CuSO4(aq) copper(II) sulfate solution
Series 3. MERCURY
Series 4. SULFUR
Hg mercury (liquid)
S sulfur (powder)
HgCl2 mercury(II) chloride
SO2 sulfur dioxide
HgI2 mercury(II) iodide
FeS iron(II) sulfide
FeSO4 • 7H2O iron(II) sulfate heptahydrate
Series 5. SILICON
and CARBON
Series 6. IODINE
Si silicon (crystal)
I2 iodine (crystal)
C carbon (graphite)
KI potassium iodide
SiC silicon carbide (carborundum)
KI(aq) potassium iodide solution
SiO2 silicon dioxide (quartz)
I2 iodine in ethyl alcohol (C2H6O)
SiCl4 silicon tetrachloride
I2 iodine in carbon tetrachloride (CCl4)
(continued on next page)
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Series 7. GASES
Series 8. ORGANIC CHEMICALS
Cl2 chlorine
C6H14
Br2 bromine (vapor over liquid)
HC2H3O2 Acetic acid
NO2 nitrogen dioxide
C4H8O2 Ethyl acetate
N2 nitrogen
C3H6O Acetone
NH3 ammonia
C2H6O Ethyl alcohol (Ethanol)
H2 Hydrogen
C10H8O 1-naphthol
He helium
C7H6O Benzaldehyde
(Caution: do not remove the stoppers)
(Wear your protective eyewear. Open the containers
smell carefully and then close them tightly.)
Hexane
C8H8O2 Methyl Benzoate
C9H8O Trans-cinnamaldehyde
Series 9. MIXTURES
Brass
Sand + CuSO4 • 5H2O
White wine
Water and BaSO4
Oil and water
Vinegar
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PURE SUBSTANCE
SAMPLE OF MATTER
Element
MIXTURE
Compound
Homo-
Metal
Nonmetal
Ionic
Covalent
Hetero-
(color; physical state;
transparent/translucent/opaque; shiny/dull
malleable/ductile, shape and size of particles
for solids; odor for organic series only)
Series 1. IRON
_________________________
iron (powder, wire)
_________________________
iron(III) oxide
_________________________
iron(III) nitrate nonahydrate
_________________________
iron(III) nitrate solution
Series 2. COPPER
Cu(s)
_________________________
Cu2O(s)
_________________________
CuO(s)
_________________________
CuSO4 • 5H2O(s)
_________________________
CuSO4(aq)
_________________________
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DESCRIPTION
Page 60
BASIC
UNIT
(for Pure
substances
only)
atom,
molecule, or
formula unit
PURE SUBSTANCE
Element
SAMPLE OF MATTER
MIXTURE
Compound
Homo-
Metal
Nonmetal
Ionic
Hetero-
Covalent
Series 3. MERCURY
_______________________
mercury (liquid)
_______________________
mercury(II) chloride
_______________________
mercury(II) iodide
Series 4. SULFUR
S(s)
_______________________
FeS(s)
_______________________
FeSO4 • 7H2O(s)
_______________________
SO2 (g)
_______________________
Chemistry 60 lab
Page 61
DESCRIPTION
(color; physical state;
transparent/translucent/opaque;
shiny/dull malleable/ductile, shape
and size of particles for solids; odor
for organic series only)
BASIC
UNIT
(for Pure
substances
only)
atom,
molecule,
or formula
unit
PURE SUBSTANCE
Element
SAMPLE OF MATTER
MIXTURE
Compound
Homo-
Metal
Nonmetal
Ionic
Hetero-
Covalent
Series 5. SILICON AND CARBON
Si
_______________________
C (graphite)
_______________________
SiC
_______________________
SiO2
_______________________
SiCl4
_______________________
Series 6. IODINE
I2
_______________________
KI
_______________________
KI (aq)
_______________________
I2 in C2H6O
_______________________
I2 in CCl4
_______________________
Chemistry 60 lab
Page 62
DESCRIPTION
(color; physical state;
transparent/translucent/opaque;
shiny/dull malleable/ductile, shape and
size of particles for solids; odor for
organic series only)
BASIC
UNIT
(for Pure
substances
only)
atom,
molecule, or
formula unit
PURE SUBSTANCE
SAMPLE OF
MATTER
Element
MIXTURE
Compound
Homo-
Metal
Nonmetal
Ionic
Covalent
Hetero-
(color; physical state;
transparent/translucent/opaque; shiny/dull
malleable/ductile, shape and size of particles
for solids; odor for organic series only)
Series 7. GASES
____________________
Chlorine
____________________
bromine (vapor over liquid)
____________________
nitrogen dioxide
____________________
Nitrogen
____________________
Ammonia
____________________
Hydrogen
____________________
Helium
Chemistry 60 lab
DESCRIPTION
Page 63
BASIC
UNIT
(for Pure
substances
only)
atom,
molecule, or
formula unit
PURE SUBSTANCE
SAMPLE OF
MATTER
Element
MIXTURE
Compound
Homo-
Metal
Nonmetal
Ionic
Covalent
Hetero-
DESCRIPTION
(color; physical state; transparent/translucent/opaque;
shiny/dull malleable/ductile, shape and size of particles
for solids; odor for organic series only)
Series 8. ORGANIC CHEMICALS
C6H14(l)
Hexane
HC2H3O2(aq)
__________________
C4H8O2(l)
ethyl acetate
C3H60(l)
Acetone
C2H6O(l)
ethyl alcohol (ethanol)
C10H8O(s)
1-naphthol
C7H6O(l)
Benzaldehyde
C8H8O2(l)
methyl benzoate
C9H8O(l)
trans-cinnamaldehyde
Chemistry 60 lab
Page 64
BASIC
UNIT
(for Pure
substances
only)
atom,
molecule, or
formula unit
PURE SUBSTANCE
SAMPLE OF MATTER
Element
MIXTURE
Compound
Homo-
Metal
Nonmetal
DESCRIPTION
Ionic
Covalent
Hetero-
(color; physical state;
transparent/translucent/opaque; shiny/dull
malleable/ductile, shape and size of particles for
solids; odor for organic series only)
Series…