CliffsNotes Chemistry Quick Review: 2nd Edition (Cliffsnotes Quick Review) - Softcover

Nathan, Harold D.; Henrickson, Charles H.

Softcover

ISBN 10: 0470905433 ISBN 13: 9780470905432

Publisher: Cliffs Notes Inc.,U.S., 2011

Synopsis

CliffsNotes Quick Review gives you a clear, concise, easy-to-use review of the basics. Introducing each topic, defining key terms, and carefully walking you through sample problems, this guide helps you grasp and understand the important concepts needed to succeed. Access 500 additional practice questions at www.cliffsnotes.com/go/quiz/chemistry

About this title

Excerpt. © Reprinted by permission. All rights reserved.

CliffsNotes Chemistry Quick Review

By Harold D. Nathan Charles Henrickson

John Wiley & Sons

Copyright © 2011 John Wiley & Sons, Ltd
All right reserved.
ISBN: 978-0-470-90543-2

Chapter One

ELEMENTS

Chapter Check-In

Discovering the building blocks of all substances

Understanding the arrangement of the periodic table

Learning about atomic number and atomic weight

With all the different substances that exist, you may be surprised to learn that they are formed from a relatively small number of elements. For example, carbon is one of 112 known elements. Yet carbon can combine with other elements (like hydrogen and oxygen) to form thousands of substances (like sugar, alcohol, and plastics). Although some elements have been known from the earliest times, most were discovered during the last 300 years.

Each element is unique with its own characteristics. Each element is represented conveniently by a symbol. For example: H is for hydrogen, O for oxygen, and Cl for chlorine. In addition, each element has its own atomic number and atomic weight. The symbols of the elements, along with their respective atomic numbers and atomic weights, are displayed in a special arrangement called the periodic table. Although the elements are arranged in order of increasing atomic number, the elements with similar chemical characteristics appear in columns. This organization is useful for studying chemistry.

Discovery and Similarity

The modern science of chemistry began during the eighteenth century, when several brilliant natural philosophers classified the products of decomposition into a small number of fundamental substances. For example, in 1774, the Englishman Joseph Priestley discovered that when the red powder mercuric oxide was heated, it decomposed to liquid metal mercury and to a colorless gas capable of supporting combustion. (This gas was later named oxygen.) Most substances similarly can be decomposed into several simpler substances by heat or by an electrical current; however, the most fundamental substances cannot be broken down further, even with extraordinary temperature or electric voltage. These basic building blocks of all other substances are known as the chemical elements.

When the French chemist Antoine Lavoisier published his famous list of elements in 1789, there were only 33 elements, several of which were erroneous. By 1930, the diligent labors of thousands of chemists had increased the tally of naturally occurring chemical elements to 90. More recently, physicists in high-energy laboratories have been able to create about 20 highly radioactive, unstable elements that do not exist naturally on Earth, although they are probably produced in the hot cores of some stars.

The number of chemical elements has now reached 112, and the list is growing. Fortunately for students, only about 40 are relevant to basic chemistry. Please take a glance at the periodic table of chemical elements (found at the end of this chapter) and find calcium, element number 20. You need to be acquainted with the symbol and general properties of the 20 simplest elements up to calcium, plus another 20 of chemical significance that you will encounter in this book.

Notice that the key concepts of chemistry are set in boldface on their first appearance in the text to alert you to their importance. These terms are used repeatedly in this book, and you cannot master chemistry without understanding them. You can find the definitions of these key concepts in the glossary at the end of this book.

Groups of elements in nature have similar chemical properties. Helium, neon, argon, krypton, and xenon are all colorless gases, only two of which combine with other elements under very special conditions; their lack of reactivity leads to the name inert gases (or noble gases) for this group of similar elements. By contrast, fluorine and chlorine are corrosive greenish gases that form salts when they readily combine with metals, hence the name halogens (salt formers) for fluorine, chlorine, bromine, and iodine.

As a final example of a group of elements with similar properties, the metallic elements lithium, sodium, and potassium have such low densities that they float on water and are so highly reactive that they spontaneously burn by extracting oxygen from the water itself. These light metals form strong alkalis and appropriately are called the alkali metals. You should locate each of these columns of similar elements, as shown in Figure 1-1, on the periodic table.

Similar elements also occur in the same natural environment. For instance, the halogens are markedly concentrated in seawater. (The major salt in ocean brines is sodium chloride.) The other halogens are extracted from seawater that has been further concentrated—bromine from salt beds formed by evaporation and iodine from kelp, which grows in oceans.

The first indispensable key to making sense out of the extensive system of facts and principles called chemistry is the rule that the behavior of an element or compound can be predicted from similar substances.

Atomic Masses

By the early nineteenth century, chemists were striving to organize their rudimentary knowledge of the chemical elements. It was known that differing masses of elements reacted to form compounds. For example, they found that 3 grams of magnesium metal reacted with precisely 2 grams of oxygen to form magnesium oxide with no residual magnesium or oxygen. The same mass of oxygen, however, required 5 grams of calcium metal to react completely to form calcium oxide. Table 1-1 summarizes these relative combining masses.

Chemists gradually discovered that such relative masses in chemical reactions were fundamental characteristics of the elements. The English chemist John Dalton realized that all the known combining masses were nearly whole-number multiples of the combining mass of the lightest element—hydrogen. In 1803, he proposed an atomic theory in which all other elements would be built from multiple hydrogen atoms. Consequently, he based his scale of atomic masses on hydrogen being equal to 1.

Although Dalton's theory was found to be unrealistically simple, he did compel chemists to adopt a standard scale of atomic weights. Because the combining mass of oxygen is approximately 16 times that of hydrogen, the preceding chart can be revised, as shown in Table 1-2.

The modern masses for calcium, magnesium, and oxygen are still nearly in the 5:3:2 ratios of the original masses. Notice especially that the atomic mass of hydrogen is not precisely equal to 1, because the atomic mass scale is now based on the most common variety of carbon being exactly 12 atomic mass units. Dalton's bold conjecture that all the heavier elements have masses that are integral multiples of hydrogen is not strictly valid, but his theory was a good approximation that eventually led to the discovery of the particles composing the atoms.

The Periodic Table

In 1869, the Russian chemist Dmitri Mendeleyev published his great systematization called the periodic table. He arranged all known chemical elements in order of their atomic masses and found that similar physical and chemical properties recurred every 7 elements for the lighter elements and every 17 elements for the heavier ones. (The inert gases had not been discovered at that time; the correct values for similar properties are 8 and 18.) The periodic table is based on atomic masses and similar properties. In each row, the atomic masses increase toward the right. Each column contains a group of elements with similar chemical behavior.

In the modern periodic table, each box contains four data, as shown in Figure 1-2. Besides the element name and symbol, the atomic mass is at the bottom, and the atomic number is at the top. The elements are arranged in order of increasing atomic number in horizontal rows called periods.

In the preceding section, you reviewed the historical determination of atomic masses. Notice that the elements also seem to be arranged in order of increasing atomic mass. But several exceptions exist. Compare the atomic mass of tellurium (Te) to iodine (I). (See Figure 1-3.)

Consider the proper placement of tellurium and iodine in the periodic table, as shown in Figure 1-3. Te has the heavier atomic mass. The chemical properties of tellurium are like those of selenium because both are semimetallic elements that form compounds like those of sulfur. Iodine resembles bromine because these elements are nonmetallic halogens that form compounds like those of chlorine. Therefore, the order in the table cannot be based solely on atomic mass.

The atomic number, which appears above each element symbol, represents the meaningful order in the periodic table. When an element is referred to by an integer, this number means the atomic number, not the atomic mass. Thus, element 27 is cobalt (whose atomic number is 27), not aluminum (whose atomic mass is 27). In Chapters 2 and 3, these two concepts are more carefully defined; for now, simply bear in mind the distinction between atomic number and atomic mass.

The periodic table displays the pattern of properties of the elements. The lightest are at the top of the chart; the atomic masses increase toward the bottom of the chart. The elements to the upper right, above a diagonal line from aluminum (13) to polonium (84), are nonmetals, about half of which exist as gases under normal laboratory conditions. All the elements in the middle and left of the table are metals, except gaseous hydrogen (1). Most of the metals are shiny, deformable solids, but mercury has such a low melting point that it is a liquid at room temperature. All the metals have high conductivities for heat and electricity. Many simple chemical compounds are formed from a metal reacting with a nonmetal.

In the periodic table, shown on the following pages, elements in columns have similar properties, and elements so related (like sulfur, selenium, and tellurium) are members of the same group or family and are congeners of one another.

This information was obtained from the International Union of Pure and Applied Chemistry (IUPAC) Web site: www.IUPAC.org.

Chapter Check-Out

1. The number of known elements is approximately ______.

a. 44 b. 112 c. 520

2. Elements aligned in columns on the periodic table ______.

a. have similar atomic weights b. have similar atomic numbers c. have similar chemical properties

3. The atomic number of iron (Fe) is ______.

a. 26 b. 53 c. 55 d. 85

4. Element number 4 is ______.

a. helium (He) b. beryllium (Be) c. carbon (C)

5. Ultimately, the periodic table arranges elements from left to right in order of ______.

a. the first letter of the element's name b. increasing atomic numbers c. increasing atomic weights

Answers: 1. b 2. c 3. a 4. b 5. b

(Continues...)

Excerpted from CliffsNotes Chemistry Quick Reviewby Harold D. Nathan Charles Henrickson Copyright © 2011 by John Wiley & Sons, Ltd. Excerpted by permission of John Wiley & Sons. All rights reserved. No part of this excerpt may be reproduced or reprinted without permission in writing from the publisher.
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Edition details

• Publisher: Cliffs Notes Inc.,U.S.
• Publication date: 2011
• ISBN 10: 0470905433
• ISBN 13: 9780470905432
• Binding: Paperback
• Language: English
• Edition number: 2
• Number of pages: 192