Halogens and Noble Gases, Second Edition

Brian Nordstrom, Monica Halka

Description

In spite of their adjacency in the periodic table, halogens and nonmetals have very different properties. Halogens are among the most chemically reactive elements in the periodic table, exhibiting a diverse chemistry in terms of the large numbers of compounds they can form. On the other hand, noble gases are the least chemically reactive elements. In fact, before the 1960s, chemists referred to these elements as inert gases, because it was believed that they exhibited no chemistry whatsoever.

Providing the basics of these elements, including their role in history and some of the important scientists involved in their discovery, this newly updated, full-color resource features up-to-date scientific understanding in a clear and accessible format. Halogens and Noble Gases, Second Edition examines the ways humans use halogens and noble gases and the resulting benefits and challenges to society, health, and the environment. Fluorine, chlorine, bromine, iodine, helium, and krypton are covered in this eBook, along with the fundamentals of chemistry and physics as well as possible future developments in halogen and noble gas science and its applications.

Sujets

• Chimie
• Science
• Discovery
• Research
• Elements
• Chemistry
• Element
• Physics
• Chemical substance
• Sciences et techniques
• Chemical

Informations

• Publié par: Infobase Publishing
• Date de parution: 01 décembre 2019
• Nombre de lectures: 0
• EAN13: 9781438182094
• Langue: English
• Poids de l'ouvrage: 1 Mo

Extrait

Halogens and Noble Gases, Second Edition
Copyright © 2019 by Monica Halka, Ph.D., and Brian Nordstrom, Ed.D.
All rights reserved. No part of this publication may be reproduced or utilized in any form or by any means, electronic or mechanical, including photocopying, recording, or by any information storage or retrieval systems, without permission in writing from the publisher. For more information, contact:
Facts On File An imprint of Infobase 132 West 31st Street New York NY 10001
ISBN 978-1-4381-8209-4
You can find Facts On File on the World Wide Web at http://www.infobase.com
Contents Chapters Chemistry and Physics Background Halogens Fluorine Chlorine Bromine Iodine and Astatine Noble Gases Helium Neon Argon Krypton and Xenon Radon Support Materials Chronology Further Resources General Resources Index
Chapters
Chemistry and Physics Background

What is an element? To the ancient Greeks, everything on Earth was made from only four elements—earth, air, fire, and water. Celestial bodies—the Sun, moon, planets, and stars—were made of a fifth element: ether. Only gradually did the concept of an element become more specific.
An important observation about nature was that substances can change into other substances. For example, wood burns, producing heat, light, and smoke and leaving ash. Pure metals like gold, copper, silver, iron, and lead can be smelted from their ores. Grape juice can be fermented to make wine and barley fermented to make beer. Food can be cooked; food can also putrefy. The baking of clay converts it into bricks and pottery. These changes are all examples of chemical reactions. Alchemists' careful observations of many chemical reactions greatly helped them to clarify the differences between the most elementary substances ("elements") and combinations of elementary substances ("compounds" or "mixtures").
Elements came to be recognized as simple substances that cannot be decomposed into other even simpler substances by chemical reactions. Some of the elements that had been identified by the Middle Ages are easily recognized in the periodic table because they still have chemical symbols that come from their Latin names. These elements are listed in the following table.

Elements Known to Ancient People Iron: Fe ("ferrum") Copper: Cu ("cuprum") Silver: Ag ("argentum") Gold: Au ("aurum") Lead: Pb ("plumbum") Tin: Sn ("stannum") Antimony: Sb ("stibium") Mercury: Hg ("hydrargyrum") *Sodium: Na ("natrium") *Potassium: K ("kalium") Sulfur: S ("sulfur") *Sodium and potassium were not isolated as pure elements until the early 1800s, but some of their salts were known to ancient people.

Modern atomic theory began with the work of the English chemist John Dalton in the first decade of the 19th century. As the concept of the atomic composition of matter developed, chemists began to define elements as simple substances that contain only one kind of atom. Because scientists in the 19th century lacked any experimental apparatus capable of probing the structure of atoms, the 19th-century model of the atom was rather simple. Atoms were thought of as small spheres of uniform density; atoms of different elements differed only in their masses. Despite the simplicity of this model of the atom, it was a great step forward in our understanding of the nature of matter. Elements could be defined as simple substances containing only one kind of atom. Compounds are simple substances that contain more than one kind of atom. Because atoms have definite masses, and only whole numbers of atoms can combine to make molecules, the different elements that make up compounds are found in definite proportions by mass. (For example, a molecule of water contains one oxygen atom and two hydrogen atoms, or a mass ratio of oxygen-to-hydrogen of about 8:1.) Since atoms are neither created nor destroyed during ordinary chemical reactions ("ordinary" meaning in contrast to "nuclear" reactions), what happens in chemical reactions is that atoms are rearranged into combinations that differ from the original reactants, but in doing so, the total mass is conserved. Mixtures are combinations of elements that are not in definite proportions. (In salt water, for example, the salt could be 3 percent by mass, or 5 percent by mass, or many other possibilities; regardless of the percentage of salt, it would still be called "salt water.") Chemical reactions are not required to separate the components of mixtures; the components of mixtures can be separated by physical processes such as distillation, evaporation, or precipitation. Examples of elements, compounds, and mixtures are listed in the following table.
The definition of an element became more precise at the dawn of the 20th century with the discovery of the proton. We now know that an atom has a small center called the "nucleus." In the nucleus are one or more protons, positively charged particles, the number of which determine an atom's identity. The number of protons an atom has is referred to as its "atomic number." Hydrogen, the lightest element, has an atomic number of 1, which means each of its atoms contains a single proton. The next element, helium, has an atomic number of 2, which means each of its atoms contain two protons. Lithium has an atomic number of 3, so its atoms have three protons, and so forth, all the way through the periodic table. Atomic nuclei also contain neutrons, but atoms of the same element can have different numbers of neutrons; we call atoms of the same element with different number of neutrons "isotopes."

Examples of Elements, Compounds, and Mixtures Elements Compounds Mixtures Hydrogen Water Salt water Oxygen Carbon dioxide Air Carbon Propane Natural gas Sodium Table salt Salt and pepper Iron Hemoglobin Blood Silicon Silicon dioxide Sand

There are roughly 92 naturally occurring elements—hydrogen through uranium. Of those 92, two elements, technetium (element 43) and promethium (element 61), may once have occurred naturally on Earth, but the atoms that originally occurred on Earth have decayed away, and those two elements are now produced artificially in nuclear reactors. In fact, technetium is produced in significant quantities because of its daily use by hospitals in nuclear medicine. Some of the other first 92 elements—polonium, astatine, and francium, for exam-ple—are so radioactive that they exist in only tiny amounts.
All of the elements with atomic numbers greater than 92—the so-called transuranium elements—are produced artificially in nuclear reactors or particle accelerators. Between 2000 and 2009, the discoveries of elements 113, 115, 117, and 118 were reported. On November 28, 2016, the International Union of Pure and Applied Chemistry (IUPAC) approved names for these elements, thereby completing period seven of the periodic table. Element 113 was named nihonium (Nh), element 115 was named moscovium (Mc), element 117 was named tennessine (Ts), and element 118 was named oganesson (Og).
Scientists are now exploring the possibility of creating still larger elements, although technological breakthroughs will be needed to synthesize elements beyond 120. Even if they are produced, large atomic nuclei tend to be extremely unstable and short-lived because the protons in the nucleus repel one another. However, scientists predict that extra super heavy elements within a region called the island of stability, which have the “magic number” of protons and neutrons that maximize binding energy holding the nucleus together, will be relatively stable. Element 126 is in the island of stability and, if it is ever produced, is predicted to have an isotope stable enough to be studied and used.
When the Russian chemist Dmitri Mendeleev (1834–1907) developed his version of the periodic table in 1869, he arranged the elements known at that time in order of atomic mass or atomic weight so that they fell into columns called groups or families consisting of elements with similar chemical and physical properties. By doing so, the rows exhibit periodic trends in properties going from left to right across the table, hence the reference to rows as periods and name "periodic table."
Mendeleev's table was not the first periodic table, nor was Mendeleev the first person to notice triads or other groupings of elements with similar properties. What made Mendeleev's table successful and the one we use today are two innovative features. In the 1860s, the concept of atomic number had not yet been developed, only the concept of atomic mass. Elements were always listed in order of their atomic masses, beginning with the lightest element, hydrogen, and ending with the heaviest element known at that time, uranium. Gallium and germanium, however, had not yet been discovered. Therefore, if one were listing the known elements in order of atomic mass, arsenic would follow zinc, but that would place arsenic between aluminum and indium. That does not make sense because arsenic's properties are much more like those of phosphorus and antimony, not like those of aluminum and indium.
To place arsenic in its "proper" position, Mendeleev's first innovation was to leave two blank spaces in the table after zinc. He called the first element eka-aluminum and the second element eka-silicon, which he said corresponded to elements that had not yet been discovered but whose properties would resemble the properties of aluminum and silicon, respectively. Not only did Mendeleev predict the elements' existence, he also estimated what their physical and chemical properties should be in analogy to the elements near them. Shortly afterward, these two elements were discovered and their properties were found to be very close to what Mendeleev had predicted. Eka-aluminum