The periodic table

The total number of different states with principal quantum number n is obtained from the sum 

In Section 8.6 we will see that symmetry considerations alone, without any appeals to special functions or series solutions, will allow us to predict these results from the model, up to a factor of two. 

Tilf itottiic jAf/L CftloveA Kit A iAr jAabrb orAitiil jfD cl- kAicA tAfy irr 

Why should the spectral data for the alkali atoms resemble the spectral data for hydrogen Our model of the hydrogen atom, along with the Pauli exclusion principle (Section 1.2) and some other assumptions, provides an answer. For example, consider lithium, the third element in the periodic table. Its nucleus has a positive charge of three and it tends to attract three electrons. The Schrodinger operator for the behavior of a single electron in the presence of a lithium nucleus is 

The same argument can be made for each alkali atom because there is only one outer electron, one can model an alkah atom as a hydrogen-like atom with one electron and a nucleus made up of the true nucleus and the inner electrons. As above, this argument hinges on the fact that the inner electrons tend to be in the lowest possible states, while the Pauli exclusion principle forbids any two electrons from occupying the same state. And indeed, spectral data for alkali atoms resembles spectral data for hydrogen. Moreover, the chemical properties of the alkali atom is similar. For example, each combines easily with chlorine to form a salt such as potassium chloride, lithium chloride 

These views were certainly worthy of respect in 1922 Mellor, who was already well known as an author, probably knew as much factual inorganic chemistry as anyone alive. Should we take heed of Mellor today  

There are a number of anomalies in the long form, whose exposition has caused the spilling of much ink. One of Mendeleev s greatest triumphs was his prediction of a new element ( eka-silicon ) between Si and Sn in his Table. Mendeleev had the audacity to predict some chemical properties of this new element, and his prophecies were substantially fulfilled a few years later by the isolation of germanium and a preliminary exploration of its chemistry. These predictions were made simply by interpolation between Si and Sn. Chemists, and chemistry students, have come to expect that the chemical properties within a Group follow monotonic trends properties can be predicted by interpolation and extrapolation. Experimental observations which do not fit such simple trends lead to the identification of anomalies . At one extreme, there may be a tendency to sweep such anomalies discreetly under the carpet, or even to question the validity of the data at the other extreme, strenuous efforts are made to account for anomalies by means of elaborate and sometimes fanciful theorising. 

Paradoxically, germanium is a case in point. In recent years, there has been much discussion of the middle element anomaly in the p block Ga, Ge, As, Se and Br do not quite fit in with the other members of their respective Groups, although the anomaly does fade away towards the right-hand side of the Table. Anomalies have also been noted in the heavier p block elements Tl, Pb etc. 

It must be remembered that close chemical relationships are to be 

Nowadays, we need not be surprised by these observations. Compare the ground state electron configurations of Mg and Zn  

The electron configurations of the rest of the elements are derived in the same way. The M shell n = 3), with a maximum capacity of 18 electrons, consists of one 3s orbital, three 3p orbitals and five 3d orbitals. The N shell n = 4) can hold 32 electrons in one 4s, three 4p, five 4d and seven 4f orbitals. The maximum number of electrons in each shell is 2n, where n is the principal quantum number. 

To summarise, the building up procedure we have used is called the Aufbau principle. Each electron occupies one electron state, represented by four quantum numbers, one of which represents the spin of the electron. Each orbital can contain two 

Group 2/11 Group 13/III Group 14/IV Group 15/V Group 16/VI Group 17/VII Group 18/VIII 

No Electron-electron spin-orbit Moonetlc fiel Interaction repulsion (Russell- inWflon 

Spectra are a record of transitions between electron energy levels. Each spectral line can be related to the transition from one energy level to another. It is found that an ion in a magnetic field has a more complex spectrum, with more lines present, than has the same ion in the absence of the magnetic field. The presence of additional lines in the spectrum of an atom or ion when in a magnetic field is called the Zeeman effect. A similar, but different, complexity, called the Stark effect, arises in the presence of a strong electric field. The electron configurations described are not able to account for all of the observed transitions, and to derive the possible 

The majority of elements are naturally occurring. How are these elements distributed on Earth, and which are essential to hving systems  

Earth s crust extends from the surface to a depth of about 40 km (about 25 mi). Because of technical difficulties, scientists have not been able to study the iimer portions of Earth as easily as the crust. Nevertheless, it is believed that there is a solid core consisting mostly of iron at the center of Earth. Surrounding the core is a layer called the mantle, which consists of hot fluid containing iron, carbon, silicon, and sulfur. 

Of the 83 elements that are found in nature, 12 make up 99.7 percent of Earth s crust by mass. They are, in decreasing order of natural abundance, oxygen (O), silicon (Si), aluminum (Al), iron (Fe), calcium (Ca), magnesium (Mg), sodium (Na), potassium (K), titanium (Ti), hydrogen (H), phosphorus (P), and manganese (Mn). In discussing the natural abundance of the 

The accompanying table lists the essential elements in the human body. Of special interest are the trace elements, such as iron (Fe), copper (Cu), zinc (Zn), iodine (I), and cobalt (Co), which together make up about 0.1 percent of the body s mass. These elements are necessary for biological functions such as growth, transport of oxygen for metabolism, and defense against disease. There is a delicate balance in the amounts of these elements in our bodies. Too much or too little over an extended period of time can lead to serious illness, retardation, or even death. 

Element Percent by Mass Element Percent by Mass  

The chemical properties of an element are determined primarily by the protons and electrons in its atoms neutrons do not take part in chemical changes imder normal conditions. Therefore, isotopes of the same element exhibit similar chemical properties, forming the same types of compounds and displaying similar reactivities. 

Sample Problem 2.2 shows how to calculate the number of protons, neutrons, and electrons using atomic numbers and mass numbers. 

Strategy Recall that the superscript denotes the mass number (A), and the subscript denotes the atomic number (Z). In the case where no subscript is shown, as in parts (c) and (d), the atomic number can be deduced from the elemental symbol or name. For the purpose of determining the number of electrons, remember that atoms are neutral, so the number of electrons equals the number of protons. 

Setup Number of protons = Z, number of neutrons = A - Z, and number of electrons = number of protons. Recall that the 14 in carbon-14 is the mass number. 

For the hydrogen atom, and for the hydrogen-like ions such as He, Li, . with a single electron in the field of a nucleus with charge +Ze, the hamiltonian (the quantum mechanical form of the energy) is given by 

Because of the electron-electron repulsion term in Equation (7.2) the hamiltonian cannot be broken down into a sum of contributions from each electron and the Schrddinger 

each orbital characterized by particular values of n and f can accommodate 2(2f - - 1) electrons. It follows that an ns orbital with 1 = 0 can accommodate two electrons, an np orbital with = 1 can take six, an nd orbital with = 2 can take ten, and so on. 

Filling up the 4/ orbital is a feature of the lanthanides. The 4/ and 5d orbitals are of similar energy so that occasionally, as in La, Ce and Gd, one electron goes into 5d rather than 4f. Similarly, in the actinides, Ac to No, the 5/ subshell is filled in competition with 6d. 

During the next twenty-five years, eleven more elements were discovered chlorine (Cl), manganese (Mn), molybdenum (Mo), tellurium (Te), tungsten (W), zirconium (Zr), uranium (U), titanium (Ti), yttrium (Y), chromium (Cr), and beryllium (Be). 

the electric battery had been discovered by the Italian physicist Alessandro Volta. 

Early in the nineteenth century, the English chemist Humphry Davy was experimenting with a great big battery and a compound which we know as potassium hydroxide. This compound was well known, but no one fully understood what it contained. 

Davy discovered he could make a new metallic element by melting it and passing an electric current through it. 

To repeat this experiment today we would simply melt potassium hydroxide in a metal crucible attached to one terminal of a power supply. When a platinum wire from the other terminal is dipped into the molten compound, a small quantity of potassium metal forms around the end of the wire. 

We now know of the existence of over one hundred elements. A century ago, more than sixty of these were already known, and naturally attempts were made to relate the properties of all these elements in some way. One obvious method was to classify them as metals and non-metals but this clearly did not go far enough. 

Among the metals, for example, sodium and potassium are similar to each other and form similar compounds. Copper and iron are also metals having similar chemical properties but these metals are clearly different from sodium and potassium—the latter being soft metals forming mainly colourless compounds, whilst copper and iron are hard metals and form mainly coloured compounds. 

Among the non-metals, nitrogen and chlorine, for example, are gases, but phosphorus, which resembles nitrogen chemically, is a solid, as is iodine which chemically resembles chlorine. Clearly we have to consider the physical and chemical properties of the elements and their compounds if we are to establish a meaningful classification. 

By 1850. values of atomic weights (now called relative atomic masses) had been ascertained for many elements, and a knowledge of these enabled Newlands in 1864 to postulate a law of octaves. When the elements were arranged in order of increasing atomic weight, each 

IB as shown. The element at the top of each group was called the head element. Group VIII contained no head element, but was made up of a group of three elements of closely similar properties, called transitional triads . Many of these terms, for example group, period and head element, are still used, although in a slightly different way from that of Mendeleef. 

The periodic table lists the elements from left to right in the order of their atomic numbers. Each horizontal row is called a period- Tire vertical columns are called groups or families. There are at least two methods used to number the groups. The newer method is to number them 1 through 18 from left to right. An older method, which is still used, is to separate the groups into sections A and B. These sections are then numbered with Roman numerals as shown below. 

Tire periodic table below divides the elements into three sections 1) nonmetals on the right (dark orange) 2) metals on the left (light orange) and 3) metalloids along the yellow-shaded diagonal separating the metals from the nonmetals. 

Metals are large atoms that tend to lose electrons to form positive ions or form positive oxidation states. To emphasize their loose hold on their electrons and the fluid-like nature of their valence electrons, metals are often described as atoms in a sea of electrons. The easy movement of electrons within metals gives them their metallic character. Metallic character includes ductility (easily stretched), malleability (easily hammered into thin strips), thermal and electrical conductivity, and a characteristic luster. Metal atoms easily slide past each other allowing metals to be hammered into thin sheets or drawn into wires. Electrons move easily from one metal atom to the next transferring energy or charge in the form of heat or electricity. All metals but mercury exist as solids at room temperature. Metals typically form ionic oxides such as BaO, (BeO is one exception that is not ionic.) 

Know the era met eristics of metals lustrous, ductile, malleable, themcliy and eiectrirally conductive. 

Nonmetafs have diverse appearances and chemical behaviors. Generally speaking, nonmetals have lower melting points than metals. They form negative ions. Molecular substances are typically made from only nonmetals. Nonmetals form covalent oxides such as Si02 or C02. 

The periodic table lists all known elements. Each element has a special symbol that describes it. Some symbols are the first letter of the element. The first element has the letter H for hydrogen. O is for oxygen. C is for carbon. Most of the 

Note The subgroup numbers 1-18 were adopted in 1984 by the International Union of Pure and Applied Chemistry The names of elements 112-118 are the Latin equivalents of those numbers. 

The periodic table lists over 100 elements. The atomic number determines the arrangement of each element. Rows and columns help to organize the elements according to specific properties. A row going across is called a period. The atomic number in each period increases by one with each element 

MB Aluminum Silicon Phosphorus Sulfur Chlonne Argon  

Zinc Gallium Germanium Arsenio Selenium Bromine Krypton  

The periodic table is a systematic representation of elements in a particular order. From the periodic table, we can gather a tremendous amount of information about the characteristics of an element. In this chapter, we will look at the trends and other important aspects of the periodic table. 

The properties of elements are periodic functions of their atomic numbers. 

The vertical columns of elements represented in the periodic table are called groups, and the horizontal rows are called periods. There are seven periods in the periodic table. The groups are usually designated by roman numerals followed by the letter A or B as shown in the periodic table. 

The periodic table as was noted above is the playing field for the materials technologist. A systematic understanding of the chemical and physical behavior of the elements and their compounds would be invaluable but an adequate text would consist of many volumes. Thus, in place of full descriptions of the elements, some periodicities in the table are included to acquaint readers with the elements and indicate their peculiarities.  

32 electrons). These numbers are accounted for by the orbital model. As the valence electrons are strong determinants of most properties, the periodicity in electron occupation parallels the properties of the elements. The periodic table is based on those periodicities of the elements or their compounds. 

The first division of the elements is into the groups of metals and nonmetals. The diagonal going from beryllum to polonium forms the boundary of the metals groups, but that boundary is not sharp. 

To the left of the Be-Po diagonal the metallic elements form intermetallic solid compounds or alloys with each other. Metallic atoms are characterized by a small number of valence electrons and low electronegativity. These compounds and alloys are themselves usually metals but not always, e.g., CsAu is a semiconductor. The stoichiometries in intermetallics (these are the numbers n and m in the compound A BJ do not usually correspond to atomic valencies, as they would in ionic and covalent compounds. The existence regions of intermetallic compounds may be wide, which means that a compound such as A B may have continuous ranges of n and m in the same phase. In other words, intermetallics can be berthollides. 

Ionization energy, electron affinity, and electronegativity which are more fully discussed in Section 2.7, are used to describe the chemical bond in certain models and form the basis for estimating the properties of compounds. 

We urge you to keep Rutherford s planetary model of the atom (Section 1.4) in mind while reading this chapter. That model, with its consideration of the electrical forces within atoms and molecules, provides the foundation of the entire subject of chemical bonding and molecular structure. 

Consequently, the elements listed in order of increasing Z can be arranged in a chart called the periodic table, which displays, at a glance, the patterns of chemical similarity. The periodic table then permits systematic classification, interpretation, and prediction of all chemical information. 

The modern periodic table (Fig. 3.f and the inside front cover of this book) places elements in groups (arranged vertically) and periods (arranged horizontally). 

There are eight groups of representative elements, or main-group elements. In addition to the representative elements, there are ten groups (and three periods) of transition-metal elements, a period of elements with atomic numbers 57 through 71 called the rare-earth or lanthanide elements, and a period of elements from atomic numbers 89 through 103 called the actinides, all of which are unstable and 

All matter is composed of the same building blocks called atoms. There are two main components of an atom. 

The charge on a proton is equal in magnitude but opposite in sign to the charge on an electron. In a neutral atom, the number of protons in the nucleus equals the number of electrons. 

This quantity, called the atomic number, is unique to a particular element. For example, every neutral carbon atom has an atomic number of six, meaning it has six protons in its nucleus and six electrons surrounding the nucleus. 

In addition to neutral atoms, we will also encounter charged ions. 

The number of neutrons in the nucleus of a particular element can vary. Isotopes are two atoms of the same element having a different number of neutrons. The mass number of an atom is the total number of protons and neutrons in the nucleus. Isotopes have different mass numbers. 

In the modem periodic table, a horizontal row of elements is called a period, and a vertical column is a group or family. The traditional designations of groups in the United States differ from those used in Europe. The International Union of Pure and Applied Chemistry (lUPAC) has recommended that the groups be numbered 1 through 18, a recommendation that has generated considerable controversy. In this text, we will 

Groups (American tradition) lA IIA IIIB IVB VB VIB VIIB VIIIB 

Groups (European tradition) lA IIA IIIA IVA VA VIA VIIA VIII 

This was put forward by Mendeleev in 1869, predating the elucidation of the structure of the atom. 

Two groups of elements are shown below. For each element, the electronic configuration of its atoms is given. 

The three elements in each group have something in common  

Group I The atoms of these elements each have 1 electron in the outer shell. 

Using this idea, scientists have divided all the elements into groups. Here is how the arrangement works. First, the elements are listed in order of increasing atomic number. Hydrogen comes at the top, since its atomic number is 1. Next, the list is sub-divided in the following way  

FIGURE 4.1 On the periodic table, groups are the elements arranged as vertical columns, and periods are the elements in each horizontal row. 

FIGURE 4.2 On the periodic tabie, each vertical column represents a group of elements, and each horizontal row of elements represents a period. 

FIGURE 4.4 Lithium (Li), sodium (Na), and potassium (K) are some alkali metals from Group lA (1). 

Cl What physical properties do these aikaii metais have in common  

FiGURE 4.5 Chiorine (Ciz), bromine (Br2), and iodine (i2) are haiogens from Group 7A (17). 

Marie (NLP 1903, NLC 1911 ) and Pierre (NLP 1903 ) Curie took up further study of Becquerel s discovery. In their studies, they made use of instrumental apparatus, designed by Pierre Curie and his brother, to measure the uranium emanations based on the fact that these emanations turn air into a conductor of electricity. In 1898, they tested an ore named pitchblende from which the element uranium was extracted and found that the electric current produced by the pitchblende in their measuring instrument was much stronger than that produced by pure uranium. They then undertook the herculean task of isolating demonstrable amounts of two new radioactive elements, polonium and radium, from the pitchblende. In their publications, they first introduced the term radio-activity to describe the phenomenon originally discovered by Becquerel. After P. Curie s early death, M. Curie did recognize that radioactive decay (radioactivity) is an atomic property. Further understanding of radioactivity awaited the contributions of E. Rutherford. 

The early years of the twentieth century saw giant advances in man s understanding of nature which must be mentioned in any synopsis of the scientific history of this era. Thus, in 1901, M. Planck (NLP 1918 ) published his first paper on the black-body radiation law which ushered in the era of quantum mechanics. In 1905, A. Einstein (NLP 1918 ) published his Anna Mirabilis Papers on the photo effect, on Brownian motion, and on the theory of special relativity and the equivalence of matter and energy. 

In the meantime, E. Rutherford (NLC 1908 ) studied the radioactivity discovered by Becquerel and the Curies. He determined that the emanations of radioactive materials include alpha particles (or rays) which are positively charged helium atoms, beta particles (or rays) which are negatively charged electrons, and gamma rays which are similar to x-rays. He also studied the radioactive decay process and deduced the first order rate law for the disappearance of a radioactive atom, characterized by the half-life, the time in which 50% of a given radioactive species disappears, and which is independent of the concentration of that species. 

1 The Historic Papers ofSoddy and Fajans in 1913 and the Work Leading up to These Papers 

As already noted, while it is the intention for the rest of this chapter to give references to original literature, this will not be done for the (extensive) literature on studies of radioactive decay series which preceded the publication of the historic papers of Soddy and Fajans in 1913. This material was well reviewed by F. Soddy (NLC 1921 ) in his Nobel Lecture which is used as the sole reference here (Soddy 1922). 

During the 50 years after the periodic tables of Mendeleev and Meyer were proposed, experimental advances came rapidly. Some of these discoveries are listed in Table 2.1. 

Parallel discoveries in atomic spectra showed that each element emits light of specific energies when excited by an electric discharge or heat. In 1885, Balmer showed that the energies of visible light emitted by the hydrogen atom are given by the equation 

1897 J. J. Thomson Showed that electrons have a negative charge, with charge/mass = 1.76 X 10 C/kg 

1911 E. Rutherford Established the nuclear model of the atom a very small, heavy nucleus surrounded by mostly empty space 

1913 H. G. J. Moseley Determined nuclear charges by X-ray emission, establishing atomic numbers as more fundamental than atomic masses 

In 1869 and 1870 respectively, Dmitri Mendeleev and Lothar Meyer stated that the properties of the elements 

The lUPAC (Intemational Union of Pure and AppUed Chemistry) has produced guidelines for naming blocks and groups of elements in the periodic table. In summary, 

In mid-2011, the number of elements in the periodic table stood at 117 (see Section 21.5)  

Effective nuclear charges, Z jf, experienced by electrons in different atomic orbitals may be estimated using Slater s rules. These rules are based on experimental data for electron promotion and ionization energies, and Z jj is determined from the equation  

Mercury is a silvery, shiny element that is a liquid at room temperature. Mercury can enter the body through inhaled mercury vapor, contact with the skin, or ingestion of foods or water contaminated with mercury. In the body, mercury destroys proteins and disrupts cell function. Long-term exposure to mercury can damage the brain and kidneys, cause mental retardation, and decrease physical development. Blood, urine, and hair samples are used to test for mercury. 

In both freshwater and seawater, bacteria convert mercury into toxic methyhnercury, which attacks the central nervous system. Because fish absorb methyhnercury, we are exposed to mercury by consuming mercury-contaminated fish. As levels of mercury ingested from fish became a concern, the Food and Drug Administration set a maximum level of one part mercury per million parts seafood (1 ppm), which is the same as 1 mg of mercury in every kilogram of seafood. Fish higher in the food chain, such as swordfish and shark, can have such high levels of mercury that the Environmental Protection Agency recommends they be consumed no more than once a week. 

One of the worst incidents of mercury poisoning occurred in Minamata and Niigata, Japan, in 1950. At that time, the ocean was polluted with high levels of mercury from industrial wastes. Because fish were a major food in the diet, more than 2000 people were affected with mercury poisoning and died or developed neural damage. In the United States between 1988 and 1997, the use of mercury 

This mercuty fountain, housed in glass, was designed by Alexander Calder for the 1937 World s Fair in Paris. 

Use the periodic table to identify the group and the period of an element identify the element as a metal, a nonmetal, or a metalloid. 

Let us first recall from Chapters 3 and 4 that atoms are built of particles of three kinds protons, neutrons, and electrons. The nucleus of each atom is made of protons and neutrons. The number of protons (the atomic number) determines the electric charge of the nucleus, and the total number of protons and neutrons (the mass number) determines its mass. In a neutral atom the number of electrons surrounding the nucleus is equal to the atomic number. 

The periodic law states that the properties of the chemical elements are not arbitrary, hut depend upon the structure of the atom and vary with the atomic number in a systematic way. The important point is that this dependence involves a crude periodicity that shows itself in the recurrence of characteristic properties. 

For example, the elements with atomic numbers 2, 10, 18, 36, 54, and 86 are all chemically inert gases. Similarly, the elements with atomic numbers one greater —namely, 3, 11, 19, 37, 55, and 87 —are all light metals that are very reactive chemically. These six metals, lithium (3), sodium (11), potassium (19), rubidium (37), cesium (55), and francium (87), all react with chlorine to form colorless compounds that crystallize in cubes and show a cubic cleavage. The chemical formulas of these salts are similar LiCl, NaCl, KCl, RbCl, CsCl, and FrCI. The composition and properties of other compounds of these six metals are correspondingly similar, and different from those of other elements. 

The periodic recurrence of properties of the elements with increasing atomic number may be effectively emphasized by arranging the elements in a table, called the periodic table or periodic system of the elements. 

Several alternative forms of the periodic table have been proposed and used. We shall base the discussion of the elements and their properties in this book on the table shown as Table 5-1 (it is also reproduced inside the front cover of the book). 

One of the difficulties that persisted in the development of the periodic table was the controversy over assigning atomic masses. Without a unique scale for atomic mass, scientists drew different conclusions about how to arrange the elements in order.  

2 Historical Development of the Lewis/Kossel Model 2.1 The Periodic Table 

He established that the elements, if arranged according to their atomic weight, exhibit an apparent periodicity of properties and his conclusions were summarised as follows  

Elements which are similar regarding their chemical properties have atomic weights which are either of nearly the same value (e.g. Pt, Ir, Os) or which increase regularly (e.g. K, Rb, Cs). 

The elements which are the most widely diffused have small atomic weights. 

The atomic weight of an element may sometimes be amended by a knowledge of those of its contiguous elements. Thus, the atomic weight of tellurium must lie between 123 and 126 and cannot be 128. Tellurium s atomic mass is 127.6, and Mendeleev was incorrect in his assumption that atomic mass must increase with position within a period.) 

Elements are often referred to collectively by their periodic table group number (Group 1 A, Group 2A, and so on). For convenience, however, some element groups have been given special 

57 La Lamhaniun 58 Ce Cerium 59 Pr Praseo(fymium 60 Nd Neodymium 61 Pm Promethium 62 Sm Samarium 63 Eu Europium 64 Gd Gadolium 65 Tb Terbium 66 Dy Dy sprosum 67 Ho Holmium 68 Er Erbium 69 Tm Thulium 70 Yb Ytierimm 

Aebnium Thorium Prcxactinium Uranium Neptunium Plutoiium Americium Curium Beiteliufn Galifc um Einsteinium Fermium Mendetevium Nobelium 

FiQUre 2.12 The modern periodic table. The elements are arranged according to atomic number, which is shown above each element s symbol. 

The periodic table is a handy tool that correlates the properties of the elements in a systematic way and helps us to predict chemical behavior. At the turn of the twentieth centuiy. the periodic table was deemed the most predictive tool in all of science. We will take a more detailed look at this keystone of chemistry in Chapter 7. 

When elements are considered in order of increasing atomic number, it is observed that their properties are repeated in a periodic manner. For example, elements with atomic numbers 2,10, and 18 are gases that do not undergo chemical 

The periodic table gets its name from the fact that the properties of elements are repeated periodically in going from left to right across a horizontal row of elements. The table is arranged such that an element has properties similar to those of other elements above or below it in the table. Elements with similar chemical properties are called groups of elements and are contained in vertical columns in the periodic table. 

It should be noted that there is not always this consistency in property variations withm the periodic table. Physical properties change in a more or less regular manner however, there are some rather abrupt changes when one moves across a period or down a group. 

This chapter describes the connections between the one-electron configurations of the elements, the structure of the Long Form of the Periodic Table and the physical properties of the elements, namely their size, ionisation energies and electron affinities or attachment enthalpies. 

The main variations of properties of the elements that are summarised in the Periodic Table can be divided into physical and chemical properties. These will be briefly described with the elements restricted to the first four rows of the Periodic Table in order to conserve space. The three most important physical properties of the elements are their size, ionisation potential and electron affinity or attachment enthalpy each of these will be discussed briefly. 

A major step forward in answering questions hke these was made by the Russian chemist Dmitri Mendeleev (1834-1907), who, in 1869, pubhshed a periodic table of the elements, arranging the elements in a systematic fashion that, with modifications, is still followed today. With a couple of exceptions, Mendeleev hsted the elements in order of increasing atomic weight and then arranged them into rows, or periods, such that elements that fell into the same column, called a family or group, possessed similar chemical and physical properties. 

By leaving gaps in the periodic table where no known elements appeared to fit, and then estimating what the properties of those elements should be, Mendeleev s table had the great power of prediction. 

Elements are pure substances that contain only one kind of atom. Compounds are pure substances that contain more than one kind of atom. 

The term periodic means having repeated cycles. The properties of elements in the periodic table are repeated at regular intervals. 

Within a matter of a few years, other chemists used Mendeleev s predictions to discover several elements that had been unknown in 1869—galUum, germanium, and scandium. These elements fit neatly into the spaces that Mendeleev had left blank, and they had chemical and physical properties that were very close to the ones that Mendeleev had predicted. At the time Mendeleev developed his periodic table, the noble gases had not yet been discovered. Mendeleev lived to witness their discovery and their placement into their own family of elements. 

As the list of known elements expanded during the early 1800s, attempts were made to find patterns in chemical behavior. These efforts culminated in the development of the periodic table in 1869. We will have much to say about the periodic table in later chapters, but it is so important and useful that you should become acquainted with it now. You will quickly learn that the periodic table is the most significant tool that chemists use for organizing and remembering chemical facts. 

If F is a reactive nonmetal, which other element or elements shown here do you expect to also be a reactive nonmetal  

The arrangement of elements in order of increasing atomic number, with elements having similar properties placed in vertical columns, is known as the periodic table (T figure 2.15). The table shows the atomic number and atomic symbol for each element, and the atomic weight is often given as well, as in this typical entry for potassium  

You may notice slight variations in periodic tables from one book to another or between those in the lecture haU and in the text. These are simply matters of style, or they might concern the particular information included. There are no fundamental differences. 

Groups - vertical columns containing elements with similar properties 

The horizontal rows of the periodic table are called periods. The first period consists of only two elements, hydrogen (H) and helium (He). The second and 

Elements in a group often exhibit similarities in physical and chemical properties. For example, the coinage metals —copper (Cu), silver (Ag), and gold (Au)—belong to group IB. These elements are less reactive than most metals, which is why they have been traditionally used throughout the world to make coins. Many other groups in the periodic table also have names, Hsted in Table 2.3. 

Chemists like to put things together into groups based on similar properties. This process, called classification, makes studying a particulcir system much Ccisier. Scientists grouped the elements in the periodic table so they don t have to memorize the properties of individucd elements. With the periodic table, they can just remember the properties of the various groups. 

The periodic table is the most important tool a chemist possesses. So in this chapter, I show you how the elements eire arranged in the table, and I show you some important groups. I also explain how chemists and other scientists use the periodic table. 

In nature, as well as in things that humankind invents, you may notice some repeating patterns. The seasons repeat their pattern of fall, winter, spring, and summer. The tides repeat their pattern of rising and falling. Tuesday follows Monday, December follows November, and so on. A pattern of repeating order is called periodicity. 

In the mid-1800s, Dmitri Mendeleev, a Russian chemist, noticed a repeating pattern of chemical properties in the elements that were known at the time. Mendeleev arranged the elements in 

Thorium Protactinium Uranium Neptunium Plutonium Americium 

---

CH2CI-CO-CH3. Colourless lachrymatory liquid b.p. 119°C. Manufactured by treating propanone with bleaching powder or chlorine. It is used as a tear gas and is usually mixed with the more potent bromoacetone. chloro acids Complex chloroanions are formed by most elements of the periodic table by solution of oxides or chlorides in concentrated hydrochloric acid. Potassium salts are precipitated from solution when potassium chloride is added to a solution of the chloro acid, the free acids are generally unstable. 

This missing synuuetry provided a great puzzle to theorists in the early part days of quantum mechanics. Taken together, ionization potentials of the first four elements in the periodic table indicate that wavefiinctions which assign two electrons to the same single-particle fiinctions such as... 

Nuclei with spin about tliree-quarters of the periodic table, have a quadnipolar moment and as a... 

The usual acceptor and donor dopants for Al Ga As compounds are elements from groups II, IV and VI of the periodic table. Group II elements are acceptors and group VI elements are donors. Depending on the growth conditions. Si and Ge can be either donors or acceptor, i.e. amphoteric. This is of special interest in LEDs. 

Note. The electronic configuratioa of any element can easily be obtained from the periodic table by adding up the numbers of electrons in the various quantum levels. We can express these in several ways, for example electronic configuration of nickel can be written as ls 2s 2p 3s 3d 4s. or more briefly ( neon core ) 3d 4s, or even more simply as 2. 8. 14. 2... 

The detailed electronic configurations for the elements atomic numbers 5 5-86 can be obtained from the periodic table and are shown below in Table 1.5. 

The periodic table also contains horizontal periods of elements, each period beginning with an element with an outermost electron in a previously empty quantum level and ending with a noble gas. Periods 1, 2 and 3 are called short periods, the remaining are long periods Periods 4 and 5 containing a series of transition elements whilst 6 and 7 contain both a transition and a rare earth senes,... 

In any group of the periodic table we have already noted that the number of electrons in the outermost shell is the same for each element and the ionisation energy falls as the group is descended. This immediately predicts two likely properties of the elements in a group (a) their general similarity and (b) the trend towards metallic behaviour as the group is descended. We shall see that these predicted properties are borne out when we study the individual groups. 

Discuss, as far as possible, how far the valencies chosen are in agreement with expectations in the light of the position of these elements in the Periodic Table. (L. S)... 

---