Student's Guide to Fundamentals of Chemistry. Brescia, Arents, Meislich, Turk

420 Pages·1980·8.49 MB·English

by  Jo A. Beran (Auth.)

Student's Guide to
FUNDAMENTALS
OF CHEMISTRY
Fourth Edition
Brescia, Arents, Meislich, Turk
Dr. Jo A. Beran
Texas A & I University
Academic Press
A Subsidiary of Harcourt Brace Jovanovich Publishers
9
New York London Toronto Sydney San Francisco
Copyright © 1980, by Academic Press, Inc.
All rights reserved.
No part of this publication may be reproduced or
transmitted in any form or by any means, electronic
or mechanical, including photocopy, recording, or any
information storage and retrieval system, without
permission in writing from the publisher.
Academic Press, Inc.
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United Kingdom Edition published by
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ISBN: 0-12-132397-8
Printed in the United States of America
To the Student
You are about to embark upon one of the more rewarding disciplines in your
college career. Basic chemical principles are used in nearly all the sciences,
from agriculture to engineering and from biology to medicine. Although you
may not plan to become a research chemist, the principles you will study in
this course could be the very foundation of your expertise in your chosen
profession.
This book is written to parallel Fundamentals of Chemistry, Fourth Edition by
Brescia, Arents, Meislich, and Turk. The text introduces you to the basic
chemical principles covered in most general chemistry courses. You will find
that by excelling in this course, you will have a knowledge of chemistry that
parallels any in the country.
This Study Guide is written to assist you in understanding the material in
each section of every chapter in the accompanying text. Although the Study
Guide is nearly self-contained, it is not designed to replace the text as a
primary resource book. In the Study Guide each section is divided into an
Objective and a Focus. Examples with worked-out solutions appear wherever
appropriate. Representative study questions are located at the end of the
chapter with answers.
To effectively use the Study Guide, you should take the following steps:
1. Read the Objective(s) in the Study Guide before reading the correspond-
ing section in the text. This will identify the purpose and goal for that
section in the text and prepare you for what is to come.
2. Read the material in the text and work the examples. Study the explana-
tion for each principle.
3. Read the Focus and work through the examples in the Study Guide. The
focus is written to help stress the important principle(s) covered in that
section of the text. Rework the text examples, and then solve the corre-
sponding study questions at the end of the Study Guide chapter. If you
have difficulty with these questions, it indicates that you need to study
the principles and work the Study Guide examples and questions again.
4. Now you should be ready to answer the problems at the end of the
chapter in the text.
Where does the instructor enter the learning scheme? The chemistry in-
structor is another available source of information. The lecturer will explain
the principles in the most effective manner for students at your school. How-
ever, you must learn the material. You should prepare for the lecture the
same way you use this Study Guide: read the course syllabus and the corre-
sponding sections of the chapter before class; check the portions of each
vii
viii
chapter you do not understand; attend class regularly and listen especially
carefully when those difficult areas are presented. Take notes in lecture but
do not attempt to write everything the lecturer says; spend most of your time
listening. Jot down notes when those difficult areas are being discussed. If
those difficult areas are still unclear, see your instructor or the teaching
assistant as soon as possible. Also review the Study Guide again (and again, if
necessary). Do not slide over difficult areas. Chemistry builds, and what
you learn today is tomorrow's building block.
In the Study Guide, the problems are solved using a basic problem-solving
technique, that of dimensional analysis. In chemistry, every measurement has
a unit. The solution to a problem follows a format such that units are multi-
plied, divided, added, and subtracted just like the numbers. When the
units of the answer are known, the problem is solved by starting with the
given quantity and units, and by applying the conversion factors and units that
are necessary to obtain the desired quantity and units. Thus the units of the
given and desired quantities serve as the guide in the selection of conversion
factors. You will find the following format for a large number of problems in
the Study Guide:
Given quantity and units:
Desired quantity and units:
Conversion factor(s): (These are the relationships that permit you to con-
vert the given units into the desired units.)
Calculation:
desired quantity and units = given quantity and units x conversion factor
x conversion factor
units = units
Above all, become comfortable with working problems. Proficiency in sports
is similar to proficiency in chemistry; it requires hours of practice, mental
discipline, and endurance.
The answers to the questions at the end of each Study Guide chapter may
vary by one or two units in the last significant digit. This is due to the se-
quence in which numbers are rounded off when calculating the answer. Do
not become alarmed because of this difference.
The approaches to chemical principles and problem solving that appear in
this Study Guide have been developed through my experience in the instruc-
tion of over 1500 general chemistry students. Final suggestions were made by
Drs. Jerry Mills of Texas Tech University, Wayne Boring of Stephen F. Austin
State University, Emil Mucchetti of Texas A&I University, and by the authors
of the text: Frank Brescia, John Arents, Herbert Meislich, and Amos Turk.
Most importantly, I express my heartfelt gratitude to Judi, Kyle and
Greg for giving up familiy activities such as evenings out, the movies,
weekends at the beach, and our summer vacation, thus allowing me time to
complete this project.
Chapter
1
SOME FUNDAMENTAL TOOLS
OF CHEMISTRY
SECTION 1.1 SCIENTIFIC METHOD FOR CHEMISTS
Objectives
• To explain how a chemist, engineer, or an auto mechanic seeks to find a
logical explanation to a phenomenon.
• To define an element and to recognize the more common names and
symbols.
Focus
Chemistry is a logical science. A logical thought process is used to account for
the many observations made in the chemistry laboratory or in nature. From a
qualitative point of view, a chemist is like an automobile mechanic. If an
automobile engine is not operating smoothly, mechanics do not correct the
problem by immediately changing spark plugs, points, condenser, oil, and so
on. Instead they define the problem, perform a series of experiments (check
the points, plugs, compression, and so on), and finally correct it. When
mechanics work on other engines and their tested theory corrects a similar
problem, they soon become "good" mechanics.
A chemist or, for that matter, anyone who uses a rational approach in
explaining a phenomenon uses quantitative experimental methods to test a
theory.
Historically, chemists have been able to explain and test many observations
and theories by breaking down matter into two categories of pure
substances —elements and chemical combinations of two or more elements
called compounds. You should learn the names and symbols of the more
common elements.
No statement is absolutely true to the scientist. Theories valid for years have
been discarded because of further experimentation and interpretation.
Throughout this course, remember that the principles you are learning are
supported only by current experimental data. Your children or grandchil-
dren, however, may consider today's understanding of chemistry as old-
1
2 1. SOME FUNDAMENTAL TOOUS OF CHEMISTRY
fashioned and incorrect as a result of their interpretation of new and more
sophisticated experimentation. This does not mean that we should stop testing
today. The test of time only makes good theories better.
SECTION 1.2 MEASUREMENT AND THE INTERNATIONAL SYSTEM OF UNITS
Objectives
• To name the common units of length, mass, time, and temperature in the
International System of Units (SI System).
• To identify the prefixes used to express multiples or fractions of the
common units.
• To calculate the density of a substance.
• To interconvert temperature readings from the Celsius, Kelvin, and
Fahrenheit temperature scales.
Focus
The SI System is recognized as the set of standard weights and measures
relative to which all scientific measurements are made throughout the world.
For this reason, become familiar with the SI units of length, volume, mass,
time, and temperature.
In addition, smaller or larger multiples of length, volume, and mass should
be memorized, especially the prefixes from mega- through nano- listed in the
text, Table 1.2. Remember that the prefix is equivalent to some multiple of a
standard measurable unit. Also be familiar with expressing numbers in the
exponential notation discussed in the text, Appendix B-l. To illustrate:
1. centi —(c) is equivalent to 10~2. cm = 10~2 m produces two equivalent
conversion factors:
cm _ 10~2 m _
IO"2 m cm ~~
2. micro — (μ) is equivalent to IO-6, μg = 10~6 g, resulting in two equivalent
conversion factors:
Mg 10"6g j
= =
10"6g Mg
3. mega—(M) is equivalent to 106. MW (megawatt) = IO6 W and
MW _ 106W
106W " MW ~ l
The use of these "equivalent" relationships are very useful in converting to
new SI or English units.
Section 1.4 treats the mathematical method for these conversions. For now,
study the meaning and the symbols of the SI prefixes and the following
equivalent relationships between the English and the SI units:
1.2 MEASUREMENT AND THE INTERNATIONAL SYSTEM OF UNITS 3
1 liter = 1.057 quarts
2.54 cm = 1 inch
454 g = 1 pound
Also know the additional conversions in the text, Table 1.3.
Length (and Volume)
Chemists often use measurements of length when referring to the wavelength
of light, to the dimensions of an atom, or to the distance between atoms in a
molecule (bond length). Therefore, the meter (= 39.37 inches) is inconve-
niently large; instead chemists discuss length in units of nanometers (nm =
10~9 m), micrometers (μ,ιη = 10~6), or angstroms (an old unit where 1
angstrom (Â) = 10~10m).
The common SI unit for volume is the liter (L), equal to 1 dm3. Chemists
generally work with smaller volumes in the laboratory; the milliliter (mL =
10~3 L = cm3) and the microliter (μ-L = IO"6 L) are common units of volume.
Note that 1 mL = 1 cm3.
Mass
The terms mass and weight are often interchanged by chemists; technically,
however, the difference is important. Weight is a force, dependent upon
gravitational pull,
weight = mass X acceleration due to gravity
whereas mass is independent of gravitational pull. Balances are used in chemical
laboratories to measure the mass of a sample by comparing it with standard
masses. Spring-loaded scales on the other hand measure the weight of an object.
An example of a scale is the ordinary bathroom scale.
The SI unit of mass is the kilogram. Because chemists use much smaller
quantities in conducting research, the gram (g), milligram (mg), and occasion-
ally the microgram (/ug) are common units of mass.
The density of a sample, defined as its mass per unit volume, is generally
expressed in units of g/mL, g/cm3, or g/L, depending on its physical state
(liquid, solid, or gas).
Example 1.1 A 27.8 g sample of magnesium metal displaces 16.0 cm3 of water in a
graduated cylinder. Calculate the density of magnesium metal.
Answer 1.74 g/cm3
Solution Because the magnesium metal displaces 16.0 cm3 of water, then its volume
is also 16.0 cm3. The density of the sample, defined as mass/volume, is
H n itv = mass = 2789 9
'^ volume 16.0 cm 3 = 1 74 cm3
Time
Chemists frequently use time when discussing the rate of a chemical reaction.
The reaction may occur within hours, seconds, milliseconds, and even mi-
croseconds.
4 1. SOME FUNDAMENTAL TOOLS OF CHEMISTRY
Temperature
The Celsius and Kelvin temperature scales are commonly used in the chemis-
try laboratory. The Celsius scale, standardized to read 0° at the freezing point
of water and 100° at its boiling point, is by far the most common. A one-
degree interval on the Celsius scale is equal to a one-degree interval on the
Kelvin scale, except that 0°C is equal to 273.15 K.
T(K) = i(°C) + 273.15
Absolute zero is defined as 0 K, or —273.15°C. The Fahrenheit scale, no
longer used in the scientific fields, is still used by meteorologists.
A comparison of the two temperature scales shows how conversions can be
made.
212°F 100°C
180 divisions 100 divisions
32°F 0°C
Example 1.2 Convert 65.0°F to the Celsius scale.
Answer 18.3°C
Solution The mercury level of the thermometer has passed through 33 (65-32)
Fahrenheit divisions above the freezing point of water. Since there are 180 Fahrenheit
divisions compared to only 100 Celsius divisions between the same two reference
points on the two temperature scales, the mercury level passes through fewer Celsius
divisions by a ratio of 100/180, which reduces to 5/9. Therefore the mercury level
passes through 5/9 (33) or 18.3 Celsius divisions above the freezing point of water.
The formula for converting from °F to °C is:
°C = Ç^- (°F - 32°F)
Example 1.3 Convert 65.0°F to Kelvin.
Answer 291.4K
Solution From the previous example, t = 18.3°C. Then, since 7"(K) = f(°C) + 273.15,
7"(K) = 18.3°C + 273.15 = 291.4K.
The equation for converting °F to °C, °C = 5°C/9°F (°F - 32°F), can also
be used for converting from Celsius to Fahrenheit. Thus for the conversion
only one formula needs to be known.
SECTION 1.3 SIGNIFICANT FIGURES IN MEASUREMENT
Objective
• To express measurements and data calculations using the correct number of
significant figures.
1.3 SIGNIFICANT FIGURES IN MEASUREMENT 5
Focus
If the mass of a copper penny is recorded as 3.74 grams, it is assumed that
the first two digits are known accurately, but there is some uncertainty in the
last recorded digit. This implies that the accuracy of the measurement is 3.74
± 0.01 grams.
The number of significant figures in any recorded data for a physical
measurement is equal to the number of digits known accurately plus one
more that has some uncertainty. Therefore, 3.74 grams has three significant
figures.
For very large or very small numbers in which zeros are used for placement
of the decimal point, the numbers should be expressed in the proper scientific
notation. For example, 0.0031 cm is expressed as 3.1 x 10~3 cm. The number
3.1 indicates two significant figures. Scientific notation is presented in
Appendix B.l.
Scientific Significant
Measurement Notation Uncertainty Figures
4.223 g 4.223 x 10° g 4th digit 4
21.0m 2.10 x IO1 m 3rd digit 3
0.00067 mL 6.7 x 10"4mL 2nd digit 2
600.0 mg 6.000 x IO2 mg 4th digit 4
1001 Mg 1.001 x IO3 Mg 4th digit 4
Learn rules 1, 2, and 3 in the text. These rules are extremely helpful in
deciding how to transfer an answer from your calculator to an answer that
indicates the precision of the data.
Example 1.4 Determine the total mass of the following mixture. Express the answer to
the correct number of significant figures.
4.227 g gold
11.27 g silver
24.774 g water
163.2 g beaker
Answer 203.5 g
Solution The sum given by the calculator is 203.471 g. However, since the mass of
the beaker is uncertain at the "tenths" digit (±0.1 g), then the total mass should be
rounded off to ±0.1 g or to 203.5 g. (See rule 1.)
Example 1.5 Determine the density of an object that has a mass of 17.012 g and
occupies a volume of 12.6 cm3. Express your answer with the correct number of
significant figures.
Answer 1.33 g/cm3
Solution The calculator answer for the division
17.012 g (5 significant figures)
12.6 cm3 (3 significant figures)
6 1. SOME FUNDAMENTAL TOOLS OF CHEMISTRY
is 1.3501587. Since the answer can have no more significant figures than the mea-
surement which has the fewest number of significant figures, the density can be
expressed to only three significant figures. Rounding off 1.3501587 to three significant
figures produces an answer of 1.35 g/cm3. See rule 2 for multiplication and division
and rule 3 for rounding off.
Finally, when recording significant figures, remember that some conversions
are expressed by exact numbers which do not impose any limit on the
number of significant figures in the answer. For example, in the conversion
factors
12in 100 cm 103g 3 ft
lft lm 1kg 1yd
all numbers are exact.
SECTION 1.4 CONVERSION FACTORS
Objectives
• To use the mathematical procedure of dimensional analysis in solving
problems.
• To become more familiar with conversion factors in the SI and English
System of measurement
Focus
Dimensional analysis is a very convenient method for solving problems in all
the physical sciences for two reasons: (1) the units retain the physical meaning
of the measured quantity in the calculation and (2) the units serve as a guide
when doing algebraic calculations.
In nearly all chemistry problems there is a given or limiting quantity (and
unit) from which the desired quantity (and unit) can be calculated. Starting
with the given quantity and unit, appropriate conversion factor(s) are used so
that like units can be canceled.
Example 1.6 Determine the width, in inches, of an airbase 6.0 miles wide.
Answer 3.8 x 105 inches
Solution The given quantity and unit is 6.0 miles, and the desired quantity and unit is
the number of inches. In converting from miles, we know that 1 mile = 5280 feet and
1 foot = 12 inches. As conversion factors, these equivalent relationships become
1mi 5280 ft ^ 1ft 12 in
5280ft or —1— m—i and 1„2„. in or 1ft
The correct conversion factors are dependent on the logical approach to the solution.
Let's set up our problem using the dimensional analysis approach:
desired quantity = given quantity x conversion x conversion
and unit factor factor