Lanthanides and Actinides, Second Edition

By Monica Halka and Brian Nordstrom

The general public may not be familiar with lanthanides, actinides, and transactinides, but these elements comprise approximately 35 percent of the total number of known elements. Attempts to produce new elements—or new isotopes of known elements—constitute an active area of scientific research. 

Providing high school and college students with an up-to-date understanding of these elements, Lanthanides and Actinides, Second Edition explains how they were discovered, as well as the practical applications that these elements have in today's scientific, technological, medical, and military communities. Actinium, thorium, protactinium, uranium, and the transuranium elements are just some of the elements covered in this comprehensive resource. Coverage also includes past, present, and future uses of lanthanides and actinides in science and technology.

• Language: English
• Publisher: Facts On File
• Release date: Dec 1, 2019
• ISBN: 9781438182100

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Lanthanides and Actinides, 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:

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Contents

Chapters

Chemistry and Physics Background

Lanthanides

Actinium, Thorium, and Protactinium

Uranium

Transuranium Elements

Transactinide Elements

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

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

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 was called gallium and eka-silicon was called germanium. These discoveries validated the predictive power of Mendeleev's arrangement of the elements and demonstrated that Mendeleev's periodic table could be a predictive tool, not just a compendium of information that people already knew.

This illustration shows Dmitri Mendeleev's 1871 periodic table. The elements listed are the ones that were known at that time, arranged in order of increasing relative atomic mass. Mendeleev predicted the existence of elements with masses of 44, 68, and 72. His predictions were later shown to have been correct.

Source: Infobase Learning.

The second innovation Mendeleev made involved the relative placement of tellurium and iodine. If the elements are listed in strict order of their atomic masses, then iodine should be placed before tellurium, since iodine is lighter. That would place iodine in a group with sulfur and selenium and tellurium in a group with chlorine and bromine, an arrangement that does not work for either iodine or tellurium. Therefore, Mendeleev rather boldly reversed the order of tellurium and iodine so that tellurium falls below selenium and iodine falls below bromine. More than 40 years later, after Mendeleev's death, the concept of atomic number was introduced, and it was recognized that elements should be listed in order of atomic number, not atomic mass. Mendeleev's ordering was thus vindicated, since tellurium's atomic number is one less than iodine's atomic number. Before he died, Mendeleev was considered for the Nobel Prize, but did not receive sufficient votes to receive the award despite the importance of his insights.

The Periodic Table Today

All of the elements in the first 12 groups of the periodic table are referred to as metals. The first two groups of elements on the left-hand side of the table are the alkali metals and the alkaline earth metals. All of the alkali metals are extremely similar to each other in their chemical and physical properties, as, in turn, are all of the alkaline earths to each other. The 10 groups of elements in the