26 | The History and Use of Our Earth’s Chemical Elements (1830–1895) and John Alexander Reina Newlands (1837–1898) both developed charts based

on the atomic weights of the elements as known in their time. They recognized the “octal” (eight) nature of repeating characteristics of several elements, which preceded the concept of the horizontal periods for the current periodic table. Newland’s table grouped elements in vertical columns of seven elements according to their atomic mass. (The noble elements that would have been the eighth column had not yet been discovered.) Julius Meyer and Dimitri Mendeleev (1834–1907) independently produced periodic charts of the elements. However, Meyer’s chart was based on the physical characteristics of elements (e.g., volume versus mass, melting points, and so on), whereas Mendeleev placed data for each element on separate cards and then arranged them in various orders until he noticed gaps in some of the sequences of atomic masses in his card arrangement. He not only corrected the known atomic masses of some elements, but he also left blank spaces in his chart to represent unknown, yet-to-be discovered elements. When the gap was in the middle of a triad of two known elements, he would guess at the atomic mass number (the average mass of the other two elements), atomic proton number, and other characteristics for the missing element. Then he named these with the prefix “eka,” meaning “first” in Sanskrit. For instance, eka-aluminum was later identified as gallium, and eka-silicon proved to be germanium. Within 15 years the “missing” elements were discovered with the basic chemical and physical characteristics based on Mendeleev’s predictions. Meyer’s chart was published in 1870, and Mendeleev’s table appeared the following year.

Several discoveries about the structure of the atom provided information that led to improvements of the older tables and the development of the modern periodic table. For example, in 1897 J. J. Thomson discovered that the tiny negatively charge particle in the atom had a mass of less than 1/2000 of that of the nucleus of hydrogen. In 1920, Ernest Rutherford discovered that during the spontaneous disintegration of radioactive elements, such as uranium, a small but heavy positively charged particle was produced. This was later identified as the alpha particle that was actually the nucleus of helium. When Rutherford shot alpha particles toward a thin gold foil, most particles went through as if the foil did not exist. Just a few of the alpha particles bounced back, and a few were deflected sideways of the foil, giving evidence of the small, heavy positively charged nucleus surrounded by vast amounts of space. In 1932, James Chadwick (1891–1974) discovered the neutron, which had no electrical charge but was of approximately the same mass as the proton. In the late 1800s, Henry Gwyn Moseley (1887–1915) used an X-ray spectrometer to examine the electromagnetic wavelengths of different atoms that he exposed to the X-rays. He observed that each element produced its own specific wavelength, which he then considered as a separate integer that was proportional to the square root of the frequency of its wavelength. This integer is now referred to as the element’s atomic number.

However, Mendeleev received credit for devising the modern periodic table of the elements, even though his table was based on atomic mass numbers rather than the atomic proton numbers of the elements. In 1871 he arranged the elements not only by their atomic mass in horizontal rows (periods), but also in vertical columns (groups, also called families) by their valences as well as other chemical and physical characteristics.

The conﬁguration of electrons around the nuclei of atoms is related to the structure of the periodic table. Chemical properties of elements are mainly determined by the arrangement of electrons in the outermost valence shells of atoms. (Other factors also inﬂuence chemical

The Periodic Table of Chemical Elements | 27 properties, such as the atoms’ size and mass.) An important factor is the “ground state” of an

atom, which is the basic energy level in which atoms are usually found. The ﬁrst electron out from the atom’s nucleus is in the lowest energy state. The second shell of electrons represents the next higher level, and so on. Each atom’s shell and orbital can hold only so many electrons before additional electrons form new orbitals.

Remember, valence is the combining power, in whole numbers, of one element with another. The most common combining ratio of atoms of diﬀerent elements to form molecules is 1:1, but other ratios are possible. For instance, hydrogen (H) and chlorine (Cl) both have a valence of one, so they combine to form a molecule of HCl, hydrogen chloride or hydrochloric acid. Nitrogen has a combining power (valence) of three. When hydrogen combines with nitrogen,

it takes three H atoms for each N atom to form the resulting compound, ammonia (NH 3 ). Mendeleev’s chart with the vacant spaces provided him and other scientists a plan to pre- dict where new elements would fit as new discoveries were made. As the chart began to fill in, some problems became evident. In 1913, the British physicist Henry Gwynn Mosely, by using X-rays, identified that the positive charges of nuclei increased as the atomic mass became greater. Thus, the sizes of atoms increased. This discovery led to a correction of the periodic table. Instead of arranging the elements according to their atomic weight, Mosely arranged them by their atomic number—that is, the number of positive protons in the nucleus. This is how the periodic table is constructed today.

Let’s take a more detailed look at the periodic table and discover its great symmetry and usefulness. (Refer to the periodic table of chemical elements reproduced on the inside front and back covers of this book.)

Rules for Cataloging the Elements

1. The order in which the elements appear in the periodic table follows speciﬁc rules. 2. Changes in properties of the elements repeat in orderly ways.

3. The arrangement of the elements in the table is according to the following factors:

a. A conﬁguration of electrons surrounds the nuclei of all atoms. The electrons are negatively charged particles (units) located at diﬀerent energy levels, referred to as shells, surrounding the nucleus. For most elements, the number of the electrons in the outermost shells determines the element’s “combining power” or valence, or oxidation number.

b. The number of positively charged particles units (protons) in the nucleus of an atom dictates that atom’s atomic number. In other words, the number of protons determines the atomic number, which is also referred to as the proton number.

c. All of the positive protons plus all of the neutrons, which have no electrical charge, are found in the nucleus. Together, the total number of both the protons and neutrons determines the atomic weight (mass) of each atom as well as the isotopes for that element. For example: atomic number (protons) + number of neutrons = atomic weight. (Note: The atomic weight is actually the average of that element’s isotopes’ weights as arrived at by their respective proportions found in a particular element.)

4. Periods are the rows that run left to right and follow the octet rule.

5. Both the atomic number (protons) and electrons increase in number from left to right in each period.

28 | The History and Use of Our Earth’s Chemical Elements

6. When a principal energy level (shell) receives its full complement of electrons (e.g., inert noble gases in group 18 (VIII)), a new row begins, which is the start of new period.

7. The vertical columns represent “families” of elements that exhibit similar characteristics. These families of elements are called groups. Elements in the same group, in general, exhibit similar combining powers (valence), but do not exhibit the same degree of reac- tivity to other elements. Some of these family (group) characteristics are atomic radius, ionization energy, and electron aﬃnity.

8. Most elements in the same group (family) have the same number of valence electrons in their outermost shell. The outer electrons of an atom involved in chemical reactions are the valence electrons. These electron conﬁgurations assist in explaining the recurrence of chemical and physical properties of elements in the same group (family).

9. In general, for groups, when the table is read from left to right, the number of valence electrons in the outer shell of elements increases. For instance, lithium is located at the start of period 2 and has one electron in its outer shell and is found in group 1 (IA), and neon, located at the end of period 2 in group 18 (VIIIA), has eight electrons in its completed outer shell.

There are several methods for numbering the groups from left to right. The International Union of Pure and Applied Chemistry (IUPAC) uses a notation system of Arabic numerals, 1 to 18, for the groups. Roman numerals are used in the old form. Hydrogen is located on the upper left of the table in the ﬁrst period in group 1. Because hydrogen has properties of both a metal (electropositive) and a nonmetal (electronegative ), some periodic tables place hydrogen at the top of group VIIA, next to helium, as well as in the ﬁrst period of group 1. Hydrogen is the only element to exhibit both characteristics.

10. Metals (on the left side and central area of the table), in general, have fewer electrons in their outer valence shell than do nonmetals on the right side of the table. 11. Metals give up valence electrons (thus are electropositive) or share electrons with nonmetals. The most active metals are the ones on the left side of the table that have the least number of valence electrons.

12. Most nonmetals located in the groups on the right side of the table have more electrons in their outer valence shells than do metals. 13. The most reactive nonmetals (electronegative) are those in group 17 (VIIA) on the right side of the table. (See exception noted in rule 19c.) They tend to accept valence electrons from the metals to complete their outer valence shells from seven electrons to form full outer shells of eight electrons.

14. Atomic and ionic sizes of elements, in general, increase from the top of the table to the bottom as the atomic mass and number of electron shells increase. 15. Thus, for both metals and nonmetals, the atomic size increases as the atomic mass (weight) and atomic number increase. 16. The neutron is a fundamental particle of matter found in the nucleus. The neutron has about the same mass as the proton, but, unlike the proton, the neutron has no electrical charge. 17. The total atomic mass (weight) of an atom consists almost entirely of the total mass of both the protons and the neutrons. The negatively charged electron has a mass of less than 1/2000 that of the proton. (The precise ﬁgure is 1/1837.) Therefore, the electron’s mass is negligible when considering the total atomic mass of an atom.

The Periodic Table of Chemical Elements | 29

18. The atomic number is the number of protons (positive charge units) in the nucleus. 19. Exceptions:

a. Elements in the far right column, group 18 (VIIIA), all have completed outer shells of eight valence electrons. Thus, they do not easily react with other elements. They are known as the noble elements or inert gases. However, under certain circumstances, they can form some compounds.

b. By exchanging and sharing outer electrons in chemical reactions, atoms tend to gain a full complement of two or eight electrons in their outer shells and become neutral. If the outer shell has less than four electrons, the element normally gives up electrons in chemical reactions and, thus, is electropositive. On the other hand, if the outer shell has more than four electrons, the atom tends to accept or gain enough electrons to complete the outer shell with a complement of eight electrons. Those atoms that gain electrons are electronegative. The result is that atoms that have combined to form a new molecule have now become more stable.

c. Two elements that do not ﬁt the primary pattern of the table are hydrogen and helium. In some forms of the periodic table, hydrogen and helium are placed together in a separate category because they do not seem to ﬁt any other position in the table.

Hydrogen, with one valence electron tends to be a diatomic gas (H 2 ). Although not a metal, hydrogen is sometimes placed in the table to head the alkali metals in the ﬁrst group. Hydrogen has one electron in the K shell, which is its ﬁrst and only shell. Because as an atom it has only one electron, it can also collect an electron from another element,

including another metal, to form a hydride—for example, 2Na + H 2 → 2 NaH (sodium hydride). So hydrogen might be thought of as both electropositive and electronegative considering that it can give up its electron, receive an electron, or share its electron. (See the entry for hydrogen in the guide section of this book for more on hydrides.)

Helium is placed at the top of the far right column of group 18 (VIIIA) because it has two electrons in the K shell (the ﬁrst shell), which completes its outer valence shell. Helium is inactive. Therefore, it is included in the group with the noble inert gases, even though it has a completed outer “K” shell of two instead of eight electrons in the outer most shell of the other inert gases. 20. The transition elements are found in three series in the center of the periodic table, starting

in group 3. The ﬁrst series is in period 4, from scandium ( 21 Sc) to zinc ( 30 Zn). The second series is in period 5, yttrium ( 39 Y) to cadmium ( 48 Cd). The third series is a bit diﬀerent. It is in period 6, following the lanthanide series, and starts at group 4 at the element hafnium ( 72 Hf) and continues to include mercury ( 80 Hg). The transition elements represent a change in the chemical characteristics of the elements in these series from metals to nonmetals. They show a gradual shift from being strongly electropositive (giving up electrons), as do elements in groups 1 (IA) and 2 (IIA). The shift (transition) continues to the electronegative elements, those gaining electrons, as in groups 15, 16, and 17. Thus, they progress from having some properties and characteristics similar to those of metals to having several properties more like those of metalloids, semiconductors, and nonmetals.

21. There are other series of elements that appear separately in the periodic table, namely the lanthanide series and the actinide series. The superactinide series and the super heavy elements (SHE) are additional series of newly discovered elements. These series of elements are extensions to the normal periodic order of the periodic table.

30 | The History and Use of Our Earth’s Chemical Elements

a. The lanthanide series is also considered metal-like because these elements have two electrons in their outer shells. Because they are diﬃcult to ﬁnd and identify, they were initially considered scarce. Although it has been determined that they are not scarce, they are still called rare-earth elements. There are 15 elements in this series, starting with

group 3 in period 6. They include the elements lanthanide ( 57 La) to lutetium ( 71 Lu). b. The actinide series is both metal-like and radioactive. This series also starts at group 3, but in period 7. It includes the element actinium ( 89 Ac) and ends with lawrencium ( 103 Lr). They are unstable and radioactive. 22. The transuranic elements are a subseries within the actinide series with atomic numbers higher than uranium ( 92 U). They include the actinides neptunium ( 93 Np) up to lawrencium ( 103 Lr). They are man-made, very unstable, and radioactive with very short half-lives. The actinides with higher atomic masses are so unstable that they exist only for microseconds or minutes, and only a few atoms of some have been artiﬁcially created.

The transactinide series includes elements that range from rutherfordium ( 104 Rf) up to the element with atomic number 113 or 114. The superactinide series begins where the transactinide series ends, usually with the element with atomic number 114, and continues to the element with atomic number 118.

Super heavy elements (SHE) and possible future elements may someday be discovered beyond element 118.

26 | The History and Use of Our Earth’s Chemical Elements (1830–1895) and John Alexander Reina Newlands (1837–1898) both developed charts based