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Chapter 2 the Elements

Chapter 2 the Elements

Chapter 2

The Elements

2.1 Categorizing – a history

Although elements such as , , , , and have been known since antiquity, the first scientific discovery of an element oc- curred in 1649 when Hennig Brand discovered . During the next 200 years, a vast body of knowledge concerning the properties of elements and their compounds was acquired by chemists. By 1869, a total of 63 elements had been discovered. As the number of known elements grew, scientists be- gan to recognize patterns in properties and began to develop classification schemes.

2.1.1 Law of Triads In 1817 Johann Dobereiner noticed that the atomic weight of fell midway between the weights of and , elements possessing sim- ilar chemical properties. In 1829, after discovering the triad com- posed of , , and and the triad of , and he proposed that nature contained triads of elements the middle element had properties that were an average of the other two members when ordered by the atomic weight (the Law of Triads). This new idea of triads became a popular area of study. Between 1829 and 1858 a number of scientists (Jean Baptiste Dumas, Leopold Gmelin, Ernst Lenssen, Max von Pettenkofer, and J.P. Cooke) found that these types of chemical relationships extended beyond the triad. During this time fluorine was added to the halogen ; , , and were

15 Chapter 2 grouped into a family while , phosphorus, , , and were classified as another. Unfortunately, research in this area was hampered by the fact that accurate values were not always available.

2.1.2 First Attempts At Designing a If a periodic table is regarded as an ordering of the chemical elements demon- strating the periodicity of chemical and physical properties, credit for the first periodic table (published in 1862) probably should be given to a French ge- ologist, A.E.Beguyer de Chancourtois. De Chancourtois transcribed a list of the elements positioned on a cylinder in terms of increasing atomic weight. When the cylinder was constructed so that 16 mass units could be written on the cylinder per turn, closely related elements were lined up vertically. This led de Chancourtois to propose that ”the properties of the elements are the properties of numbers.” De Chancourtois was first to recognize that ele- mental properties reoccur every seven elements, and using this chart, he was able to predict the stoichiometry of several metallic oxides. Unfortunately, his chart included some ions and compounds in addition to elements.

2.1.3 Law of Octaves John Newlands, an English chemist, wrote a paper in 1863 which classified the 56 established elements into 11 groups based on similar physical prop- erties, noting that many pairs of similar elements existed which differed by some multiple of eight in atomic weight. In 1864 Newlands published his version of the periodic table and proposed the Law of Octaves (by analogy with the seven intervals of the musical scale). This law stated that any given element will exhibit analogous behavior to the eighth element following it in the table. (see Figure 2.1)

2.1.4 Who Is The Father of the Periodic Table? There has been some disagreement about who deserves credit for being the ”father” of the periodic table, the German Lothar Meyer or the Russian . Both chemists produced remarkably similar results at the same time working independently of one another. Meyer’s 1864 textbook included a rather abbreviated version of a periodic table used to classify the elements. This consisted of about half of the known elements listed in order of

16 The Elements

Figure 2.1: The Law of Octaves arrangement of elements

Figure 2.2: The periodic ordering of elements according to Meyer their atomic weight and demonstrated periodic changes as a function of atomic weight. In 1868, Meyer constructed an extended table which he gave to a colleague for evaluation. Unfortunately for Meyer, Mendeleev’s table became available to the scientific community via publication (1869) before Meyer’s appeared (1870) (see Figure 2.2 and Figure 2.3). While writing a textbook on systematic inorganic , Principles of Chemistry, which appeared in thirteen editions the last being in 1947, Mendeleev organized his material in terms of the families of the known el-

17 Mendeleev's Table

Mendeleev noted patterns in the combining ratios of elements

! The Elements Lithium (Li), Sodium (Na), and Potassium (K) all formed oxides in the ratio of two atoms per oxygen : R2O ! The Elements (Be), (Mg), and Calcium (Ca) all formed oxides in the ratio of one atom per oxygen atom: RO ! (B) and Aluminum (Al) formed R2O3 ! (C) and (Si) formed RO2

Recognizing the patterns of combining ratios or "valency", Mendeleev created a table organized by placing elements with similar combining ratios in the same group. He arranged the elements within a group in order of their atomic mass. Chapter 2

In 1869, the Russian chemist Mendeleev noted that the repeating patterns of behavior could be arranged in a sequence of elements giving rise to the "Periodic Table" of the elements.

Figure 2.3: The periodic ordering of elements according to Mendeleev. The spaces marked with blank lines represent elements that Mendeleev deduced existed but were unknown at the time, so he left places for them in the table. The symbols at the top of the columns (e.g. R2O and RH4) are molecular Specialformulas thanks written to Dr. Paul in Karol's the style "Intro of to the Modern 19th Chemistry" century for providing much of the information on this page.

ements which displayed similar properties. The first part of the text was http://periodic.lanl.gov/mendeleev.htmdevoted to the well known (1 of 2) chemistry [10/24/2001 5:41:07 of the PM] . Next, he chose to cover the chemistry of the metallic elements in order of combining power – alkali metals first (combining power of one), alkaline earths (two), etc. How- ever, it was difficult to classify metals such as copper and mercury which had multiple combining powers, sometimes one and other times two. While trying to sort out this dilemma, Mendeleev noticed patterns in the properties and atomic weights of halogens, alkali metals and alkaline metals. He observed similarities between the series Cl-K-Ca , Br-Rb-Sr and I-Cs-Ba. In an effort to extend this pattern to other elements, he created a card for each of the 63 known elements. Each card contained the element’s symbol, atomic weight and its characteristic chemical and physical properties. When Mendeleev arranged the cards on a table in order of ascending atomic weight grouping elements of similar properties together in a manner not unlike the card ar- rangement in his favorite solitaire card game, patience, the periodic table was

18 The Elements

Figure 2.4: Mendeleev’s Tabelle II in semi-modern Form: To the modern eye, the 1869/71 formulations lacks any Group 18 rare gases and there are few f- elements: formed. From this table, Mendeleev developed his statement of the periodic law and published his work ”On the Relationship of the Properties of the Elements to their Atomic Weights” in 1869. The advantage of Mendeleev’s table over previous attempts was that it exhibited similarities not only in small units such as the triads, but showed similarities in an entire network of vertical, horizontal, and diagonal relationships. In 1906, Mendeleev came within one vote of being awarded the Nobel Prize for his work. At the time that Mendeleev developed his periodic table the experimen- tally determined atomic masses were not always accurate, so he reordered elements despite their accepted masses. For example, he changed the weight of beryllium from 14 to 9. This placed beryllium into Group 2 above mag- nesium whose properties it more closely resembled than where it had been located above nitrogen. In all Mendeleev found that 17 elements had to be moved to new positions from those indicated strictly by atomic weight for their properties to correlate with other elements. These changes indi-

19 Chapter 2 cated that there were errors in the accepted atomic weights of some elements (atomic weights were calculated from combining weights, the weight of an element that combines with a given weight of a standard.) However, even after corrections were made by redetermining atomic weights, some elements still needed to be placed out of order of their atomic weights. From the gaps present in his table, Mendeleev predicted the existence and properties of un- known elements which he called eka-aluminum, eka-boron, and eka-silicon. The elements , and were found later to fit his predictions quite well. He was, however, incorrect in suggesting that the ele- ment pairs of -potassium, - and tellurium-iodine should be interchanged in position due to inaccurate atomic weights. Although these elements did need to be interchanged, it was because of a flaw in the reasoning that periodicity is a function of atomic weight.

2.1.5 Discovery of the Noble Gases In 1895 Lord Rayleigh reported the discovery of a new gaseous element named argon which proved to be . This element did not fit any of the known periodic groups. In 1898, William Ramsey suggested that argon be placed into the periodic table between chlorine and potassium in a family with , despite the fact that argon’s atomic weight was greater than that of potassium. This group was termed the ”zero” group due to the zero valency of the elements. Ramsey accurately predicted the future discovery and properties of .

2.2 Atomic Structure and the Periodic Table

Although Mendeleev’s table demonstrated the periodic nature of the ele- ments, it remained for the discoveries of scientists of the 20th Century to explain why the properties of the elements recur periodically. In 1897 J.J. Thomson showed that cathode rays generated by applying a large voltage across a low density gas – first discovered in 1821 – were neg- atively charged. He found that they could be obtained from the discharge of any gas while using different materials for the electrodes allowing him to surmise that the negatively charged particles, now called are con- stituents of all atoms. By studying the amount of deflection of the electrons produced by electric and magnetic fields, he determined the mass-to-charge

20 The Elements

ratio me/e. Later, Millikan (1909) was able to determine just the charge of these electrons from his famous oil-drop experiment. In 1897 Becquerel found that gave off a radiation that could blacken a photographic plate even when the plate was separated from the source of radiation by thick paper. This radiation was separable by a into two electrically charged beams (α- and β-rays) and a neutral beam called γ-rays. Rutherford, in 1903, showed that the α-rays were actually a stream of helium ions, He+2, and the β-rays were later determined to be a stream of high-speed electrons. The α-rays were clear evidence that uranium atoms had disintegrated to give off helium. No longer could atoms be considered uncharged particles. Geiger and Marsden sent the He+2 ions into thin metal foils and found that all but a few of the ions passed straight though the foil. The ones that did not tended to be scatted by large angles. Rutherford explained this scattering, in 1911, on the basis of the nuclear model of the atom; most of the mass of an atoms is concentrated in a positively charged center called the nucleus, around which negatively charged electrons moved. Nuclei have 15 10 diameter of about 10− m, where as atomic diameters are about 10− m— a hundred thousand times larger. Atoms are mostly empty space. In 1919 Rutherford then discovered that nuclei appear to form when α particles collide with nitrogen atoms. Subsequently it was shown that the same thing happened with other elements besides nitrogen. These experiments clearly showed that atomic nuclei had structure and appeared to contain hydrogen nuclei, or protons. In 1930 Bothe and Becker obtained radiation that easily penetrated many layers of matter when they bombarded beryllium atoms with α particles. In 1932 this radiation was shown to consist of neutral particles with a mass approximately equal to the proton, and to be generated by many other elements when bombarded with α particles. So atoms have a nucleus that is made up of neutrons and protons, and sur- rounded by enough electrons to make the atom charge neutral. The , Z, is the number of protons in the nucleus. The mass number is the total number of protons and neutrons in the nucleus. Sodium (Na), for example, has an atomic number of 11 and an atomic mass of 23 (11 protons 23 and 12 neutrons), and is notated as 11Na. Due to the discovery, in 1913 by Thompson, that atoms of the same element could have several different atomic masses (isotopes), it became apparent that mass number (or atomic weight) was not the significant player in the periodic law as Mendeleev, Mey- ers and others had proposed, but rather, the properties of the elements varied

21 Chapter 2 periodically with atomic number. The question of why the periodic law exists was answered as scientists developed an understanding of the electronic structure of the elements be- ginning with Niels Bohr’s studies of the organization of electrons into shells.

2.3 Atomic

Chemists began studying colored flames in the eighteenth century. When sodium compounds are sprayed into a flame, the flame burns with a bright yellow color. Potassium gives a violet flame and lithium and strontium give a red flame. A flame test is a simple way to identify an element. Although the red flames of lithium and strontium appear similar, the light from each can be separated into distinctly different colors. The red Strontium flame when resolved with a prism shows a cluster of red lines and blue lines, while lithium shows a red line, a yellow line and two blue lines (see Figure 2.5). Each element has a characteristic line spectrum consisting of a unique set of frequencies. Lines appear in the ultraviolet and infrared, as well as in the visible frequency range, and they form series in which the lines become more closely spaced with decreasing wavelength, approaching a series limit at which the line spacing converges to zero. These spectra can be used not only to identify elements, but also to learn about the element’s internal structure, in particular, the electrons. The spectroscopists Balmer and Rydberg had been able to fit the fre- quencies of hydrogen, and of some of the simpler spectra, with very accurate empirical formulas. Balmer’s formula for hydrogen was

1 1 Frequency = R 4 − n2 ! " where R is a constant and n is an integer, equal to 3,4, ... This Balmer formula was similar to Rydberg’s formula, which worked for the alkali metals and some other cases. Rydberg’s formula was

1 Frequency = R const − (n d)2 ! − " where n again is an integer, d is a constant, and R is, remarkably enough, the same constant found in Balmer’s formula. This constant, called the Rydberg

22 The Elements

Figure 2.5: The emission spectra of some elements. The lines correspond to visible light emitted by atoms.

23 Chapter 2 frequency, equal to 3.29x1015 Hz, pointed to a far-reaching relation between the spectra of different≈ chemical elements. In both Balmer’s formula and Rydberg’s formula the frequency appears as a difference of two quantities, and Ritz, in 1908, showed that this concept was a general principle. It was possible to write

1 1 Frequency = R (n = n" +1,n" +2, ...) n 2 − n2 ! " " where for Balmer and Rydberg’s formula n" = 2. In the ultraviolet Lyman found the n" = 1 series, and Paschen found the n" = 3 series in the infrared. It was clear that there were discrete sets of lines that were now becoming, at least for hydrogen, well categorized. The problem with this scheme was that it seemed to be in violent dis- agreement with the implications of Rutherford’s nuclear atom picture. If one studies the motion of a charged particle, like the , moving according to classical mechanics in an inverse-square field such as the nucleus must pro- vide one finds that the electron will continually radiate energy, of a frequency equal to its rotational frequency about the nucleus. By studying the dynam- ics of a particle moving according to an inverse-square attraction, we easily find that the orbit becomes smaller and the frequency of rotation in the orbit greater, as the energy decreases. If, then, an electron rotating in such an or- bit radiates energy away, it will move into orbits of successively smaller radii and successively higher frequency of rotation, and will continue to radiate more and more energy, of higher and higher frequency, until it falls into the nucleus. This catastrophe would have to happen to a classically constructed atom consisting of a nucleus and electrons. It clearly cannot be happening with atoms of our experience; they radiate fixed frequencies, sharp spectral lines, and have permanent existence. What prevents the catastrophe? In 1913 Niels Bohr produced a set of hypotheses which circumvented the above objection and which enabled him to derive the Rydberg formula for the spectrum of hydrogen. Bohr made three basic assumptions:

1. The electrons move in circular orbits about the center of mass of the atom

2. The only allowed orbits are such that the of the atom about its center of mass is an integral multiple of h/2π.

24 The Elements

3. Radiation occurs only when an electron ”jumps” from one of the al- lowed orbits to one of lower energy. The difference in energy ∆E is then radiated as a photon of frequency ν = ∆E/h. In addition, Bohr pointed out that the hydrogen atom evidently contains only one electron, a fact not so obvious then as it is now. Following these assumptions it is possible to calculate the radius of each allowed orbit, 2 2 4π&0n ! r = , n mZe2 and then calculate the energy, which is a sum of the kinetic energy of the nu- cleus and the electron plus the negative electrostatic potential energy. (This is all done in exactly the same way that planetary orbits are calculated using the center of mass concept.)

2 4 mrZ e En = − 2 2 2 2 , 32π &0n ! where mr is the reduced mass of the two particle system (electron and nu- cleus). In the Bohr picture, therefore, the possible energies of a hydrogen atom may be represented by an diagram like that of Figure 2.6, which shows the energy corresponding to the various values of n. Arrows on the diagram indicate transitions from one level to another, leading to emission of photons which produce the spectral lines. The lowest level E1 is the level which is normally occupied by the electron, and its energy is therefore equal in magnitude to the energy required to free the electron from the proton. This energy, in electron volts, is numerically equal to the ionization potential, in volts, of atomic hydrogen. The energy levels merge into a continuum for positive values of E; these are energy levels of the free electron, which are not restricted to discrete values. The energy equation above describes the spectrum of any one-electron system – for example, 2H (or deuterium, D), 3H (or tritium, T), singly ionized helium, doubly ionized lithium. The lines of deuterium differ from the lines of ordinary hydrogen only because of the difference in the reduced mass mr. Furthermore, we should note that Bohr’s assumptions explained many other observations as well. For example, it is clear from his theory (and Figure 2.6) why the frequencies of emission and absorption lines are observed to be the same. We can also understand why not all emission lines are observed in

25 Chapter 2

Figure 2.6: Energy levels of the hydrogen atoms, according to Bohr, showing transitions which give rise to the Lyman, Balmer ,and other series of spectral lines. absorption: in a cold gas, the atoms will all be found in their ground states, and therefore only those frequencies can be absorbed which correspond to transitions from the ground state up to some excited states. The was eventually found to have some significant limitations. The largest one for the single electron hydrogen atom was the lack of fine structure that was seen in the emission line spectra. Of larger significance was the inability of the model to explain much of the emission spectra for multi- electron atoms. The resolution of these issues fell into place with the full quantum mechanical treatment (particle/wave duality) of the constituents of the atom. The full solution of the Schr¨odingerequation for the hydrogen atom is one of the canonical problems of a course, and we will not repeat it here. We will simply use the results to finish the story of the atomic spectroscopy and help understand the categorization scheme of the periodic table. Using wave mechanics, every electron in an atom is characterized by four quantum numbers. The size, shape, and spatial orientation of an electron’s probability density are specified by three of these quantum numbers. Bohr energy levels separate into electron subshells, and quantum numbers dictate the number of states within each subshell. Shells are specified by a principle

26 The Elements n, which may take on integral values beginning with unity; sometimes these shells are designated by the letters K, L, M, N, O and so on, which corresponds respectively, to n =0, 1, 2, ,3, 4, 5, .... Note also that this quantum number, and it only, is also associated with the Bohr model. This quantum number is related to the distance of an electron from the nucleus. The second quantum number, l, signifies the subshell, which is denoted by a lowercase letter – an s, p, d, or f; it is related to the shape of the electron subshell. In addition, the number of these subshells is restricted by the magnitude of n. l < n, so the number of subshells available is n 1. The number of energy states for each subshell is determined by the third− quantum number, ml. ml can take on values from l to l, so for a given l there are 2l + 1 energy states available. For an s subshell,− there is a single energy state, where as for p, d, and f subshells, three, five and seven states exist, respectively. In the absence of an external magnetic field, the states within each subshell are identical. However, when a magnetic field is applied these subshells states split each state assuming a slightly different energy. Associated with each electron is a moment, which must be oriented either up or down. Related to this spin moment is the fourth quantum number, ms, for which two values are possible (+1/2 and 1/2), one for each of the spin orientations. − A schematic energy level diagram for the various shells and subshells using the quantum mechanical model is shown in Figure 2.7. Several features of the diagram are worth noting. First, the smaller the principle quantum number, the lower the energy level; for example, the energy of a 1s state is less than that of a 2s state. Second, within each shell, the energy of a subshell level increases with the value of the l quantum number. Finally, there may be overlap in the energy of a state in one shell with states in an adjacent shell, which is especially true of the d and f states; for example, the energy of the 3d state is greater than that for the 4s. The discussion up to this point has dealt primarily with electron states – values of energy that are permitted for electrons. To determine the manner in which these states are filled with electrons, we use the Pauli exclusion principle, another quantum mechanical concept. This principle stipulates that no two electrons can have the same set of quantum numbers. This means that each electron state (a specific n, l, and ml set of quantum numbers) can hold no more than two electrons, which must have opposite spins (so that ms is not the same). This, s, p, d and f subshells may each accommodate, respectively, a total of 2, 6, 10 and 14 electrons.

27 Chapter 2

Figure 2.7: Schematic representation of the relative energies of the electrons for the various shells and subshells

Of course, not all possible states in an atom are filled with electrons. The electrons fill up the lowest possible energy states in the electron shells and subshells, two electrons (having opposite spin) per state. When all the electrons occupy the lowest possible energies in accord with the foregoing restrictions, an atom is said to be in its ground state. However, electron transitions to higher energy states are possible, and are responsible for elec- trical and optical properties of a material. The electron configuration of an atom represents the manner in which these states are occupied. In the con- ventional notation the number of electrons in each subshell is indicated by a superscript after the shell-subshell designation. For example, the electron configurations for hydrogen, helium and sodium are, respectively, 1s1,1s2 and 1s22s22p63s1. At this point, comments regarding these electron configurations are nec- essary. First, the valence electrons are those that occupy the outermost shell.

28 The Elements

These electrons are extremely important as they participate in the bonding between atoms to form atomic and molecular aggregates. Furthermore, many of the physical and chemical properties of solids are based on these valence electrons. In addition, some atoms have what is termed stable electron configura- tions; that is, the states within the outermost or shell are completely filled. Normally this corresponds to the occupation of just the s and p states for the outermost shell by a total of eight electrons, as in neon, argon, and ; on exception is helium, which contains only two 1s electrons. These elements are the inert, or noble, gases which are virtually unreactive chemically. With this knowledge of electron configuration, it is possible to arrange the elements into a periodic table with increasing atomic number, in seven horizontal rows called periods (connected to the principle quantum number n). The arrangement is such that all elements arranged in a given column or group have similar valence electron structures, as well as chemical and physical properties. These properties change gradually, moving horizontally across the and vertically down each column.

2.4 The Modern Periodic Table

The modern explanation of the pattern of the periodic table is that the elements in a group have similar configurations of the outermost electron shells of their atoms: as most chemical properties are dominated by the orbital location of the outermost electron. The last major changes to the periodic table resulted from Glenn Seaborg’s work in the middle of the 20th Century (see Figure 2.8). Starting with his discovery of in 1940, he discovered all the transuranic elements from 94 to 102. He reconfigured the periodic table by placing the series below the series. In 1951, Seaborg was awarded the Nobel Prize in chemistry for his work. Element 106 has been named (Sg) in his honor. The elements positions in Group 8A (or sometimes called Group 0), the rightmost group, are the inert gases, which have filled electron shells and stable electron configurations. Group VIIA and VIA are one and two elec- trons deficient, respectively, from having stable structures. The Group VIIA elements (F, Cl, Br, I and At) are sometimes termed the halogens.1 The

1The term halogen originates from 18th century scientific French nomenclature based

29 Chapter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c17)H+* !GA?DFB* !@@* ?1 M#1+)H+* VDFAW* D C D 9 C D 9 C D 9 C D 9 C D 9 9 D D 9 9 D !9 !9 !9 !9 ED !9 !9 DB !9 ED DB 8 G* 4 K(174/* !D?@!@A* !C* 5$ N)&).4/* D9?@9FF* ED* ;" S#1+(/)H+* AD?GC* F@* 52 5)/* !!9?A!@* 9D* !+ 6#(0* D@A?D* !!C* L9 [/H/\H(0)H+* VD9BW* GA* :% =4&+)H+* !GC?BE@ED* BB* 03 c)/-,#)/)H+* VDFDW* D E D 9 E D 9 E D 9 E D 9 E D 9 9 D D 9 9 D !9 !9 !9 !9 ED !9 !9 D9 !9 ED D9 !E !C !F !G !A :::% :<% <% <:% <::% , < - ,,P`aa$$$?0(>(2?.4+aP#1)40).a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c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

30 ! D E C F G A "4,#`*52#*-H7814HP*/H+7#1-* !b!9*$#1#*(04P,#0*)/*!B9C*7>* ,2#*:/,#1/(,)4/(&*[/)4/*4Q*3H1#* (/0*%PP&)#0*K2#+)-,1>?*52#* /(+#-*4Q*#&#+#/,-*!!Db!!9* (1#*,2#*6(,)/*#\H)X(&#/,-*4Q* ,24-#*/H+7#1-?*

Figure 2.8: Modern periodic table, (from www.dayah.com/periodic) The Elements alkali2 and alkaline earth3 metals (Li, Na, K, Be, Mg, Ca, etc...) are labeled as Groups IA and IIA, having, respectively, one and two electrons in excess of stable structures. The elements in the three long periods, Group IIIB through IIB are termed the transition metals, which have partially filled d electron states and in some cases one or two electrons in the next higher energy shell. Groups IIA, IVA and VA (B, Si, Ge, As, etc.) display char- acteristics that are intermediate between metals and nonmetals by virtue of their valence electron structures.

2.4.1 Group naming schemes There are three ways of numbering the groups of the periodic table, one using Hindu-Arabic numerals and the other two using Roman numerals. The Roman numeral names are the original traditional names of the groups; the Arabic numeral names are those recommended by the International Union of Pure and Applied Chemistry (IUPAC) to replace the old names in an attempt to reduce the confusion generated by the two older, but mutually confusing, schemes. There is considerable confusion surrounding the two old systems in use (old IUPAC and CAS) that combined the use of Roman numerals with letters. In the old IUPAC system the letters A and B were designated to the left (A) and right (B) part of the table, while in the CAS system the letters A and B were designated to main group elements (A) and transition elements (B). The former system was frequently used in Europe while the latter was most common in America. The new IUPAC scheme was developed to replace both systems as they confusingly used the same names to mean different things. The periodic table groups are as follows (in the brackets are shown the old systems: European and American): Group 1 (IA,IA): the alkali metals or hydrogen family/lithium family • on erring adaptations of Greek roots; the Greek word halos meaning ”salt”, and genes meaning ”production” referring to elements which produce a salt in union with a metal. 2The word ”alkali” is derived from Arabic Al Qaly = ”the calcined ashes”, referring to the original source of alkaline substances. Ashes were used in conjunction with animal fat to produce soap, a process known as saponification. 3The alkaline earth metals are named after their oxides, the alkaline earths, whose old- fashioned names were beryllia, magnesia, lime, strontia and baryta. These oxides are basic (alkaline) when combined with water. ”Earth” is an old term applied by early chemists to nonmetallic substances that are insoluble in water and resistant to heating–properties shared by these oxides.

31 Chapter 2

Group 2 (IIA,IIA): the alkaline earth metals or beryllium family • Group 3 (IIIA,IIIB): the scandium family • Group 4 (IVA,IVB): the family • Group 5 (VA,VB): the family • Group 6 (VIA,VIB): the family • Group 7 (VIIA,VIIB): the family • Group 8 (VIII): the family • Group 9 (VIII): the cobalt family • Group 10 (VIII): the nickel family • Group 11 (IB,IB): the coinage metals (not an IUPAC-recommended • name) or copper family

Group 12 (IIB,IIB): the family • Group 13 (IIIB,IIIA): the boron family • Group 14 (IVB,IVA): the carbon family • Group 15 (VB,VA): the (not an IUPAC-recommended name) • or nitrogen family

Group 16 (VIB,VIA): the or oxygen family • Group 17 (VIIB,VIIA): the halogens or fluorine family • Group 18 (Group 0): the noble gases or helium family/neon family •

32 The Elements

2.5 Trends through the periodic table

2.5.1 Atomic radii tends to decrease on passing along a period of the periodic table from left to right, and to increase on descending a group, see Figure 2.9. The electrons in an atom are arranged in shells which are, on average, further and further from the nucleus, and which can only hold a certain number of electrons. Each new period of the periodic table corresponds to a new shell which starts to be filled up, and so the outermost electrons are further and further from the nucleus as a group is descended. Passing along a period from left to right, the nuclear charge increases while the electrons are still entering the same shell: the effect is that the physical size of the shell (and hence of the atom) decreases in response. The increasing nuclear charge is partly counterbalanced by the increasing number of electrons in a phenomenon that is known as shielding, which is why the size of atoms usually increases as a group is descended. However, there are two occasions where shielding is less effective: in these cases, the atoms are smaller than would otherwise be expected.

Lanthanide contraction The electrons in the 4f-subshell, which is progressively filled from (Z = 58) to (Z = 71), are not particularly effective at shielding the increasing nuclear charge from the sub-shells further out. The elements im- mediately following the have atomic radii which are smaller than would be expected and which are almost identical to the atomic radii of the elements immediately above them. Hence has virtually the same atomic radius (and chemistry) as , and has an atomic radius similar to , and so forth. The effect of the lanthanide con- traction is noticeable up to (Z = 78), after which it is masked by a relativistic effect known as the inert pair effect. d-Block contraction The d-block contraction is less pronounced than the lanthanide contraction but arises from a similar cause. In this case, it is the poor shielding capacity of the 3d-electrons which affects the atomic radii and chemistries of the elements

33 Chapter 2

Figure 2.9: Periodic table of empirically measured atomic radius in picome- tres (pm) (from www.dayah.com/periodic). immediately following the first row of the transition metals, from gallium (Z = 31) to bromine (Z = 35).

2.5.2 The ionization energy I is defined as the energy that must be supplied in order to completely remove an electron from an atom. The first ionization energy is the energy required to remove one electron from the neutral parent atom. The second ionization energy is the energy required to remove a second valence electron from the univalent ion to form the divalent ion, and so on. Successive ionization energies increase. As can be seen from Figure 2.10 The decrease in atomic radius also causes the ionization energy to increase when moving from left to right across a period. The more tightly bound an element is, the more energy is required to remove an electron. Group I elements have low ionization energies because the loss of an electron generates an ion with a stable closed shell. Working down a group, each successive element has a lower ionization energy because

34 The Elements

Figure 2.10: Periodic table of measured first ionization energies in kJ/mol (1kJ/mol = 0.010364 eV/atom (from www.dayah.com/periodic) it is easier to remove an electron since the atoms are less tightly bound. Note the ionization energy of hydrogen is 13.6 eV. The energy of the 1s state discussed earlier. There are two exceptions to the general trends. The first is that the energy to remove the first electron to fill the p subshell is less than the energy to remove the last electron to fill the s subshell. For example, the first ionization energy of Be is greater than that of B. The second is that after all the p subshell orbitals have been filled with one electron (which requires 3 electrons) the addition of the fourth electron will be to a state that already has an electron of opposite spin. The electron-electron repulsion of this paired state makes it easier to remove the outermost, paired electron (this is Hunds rule, which we cover in the chapter). This last exception holds for periods 2, 3 and 4.

2.5.3 Electron affinity Electron affinity, A, reflects the ability of a neutral atom to accept an elec-

35 Chapter 2 tron. It is the energy change that occurs when an electron is added to a gaseous atom. Atoms with stronger effective nuclear charge have greater electron affinity. Some generalizations can be made about the electron affini- ties of certain groups in the periodic table, see Figure 2.11. The Group IIA elements, the alkaline earths, have low electron affinity values. These elements are relatively stable because they have filled s subshells. Group VIIA elements, the halogens, have high electron affinities because the ad- dition of an electron to an atom results in a completely filled shell. Group VIII elements, noble gases, have electron affinities near zero, since each atom possesses a filled shell and will not accept an electron readily. Elements of other groups have low electron affinities. As for ionization potential, there are two exceptions to the general trends. The first is that the Groups 13-17 the elements in the first period have lower electron affinities than the elements below them in their respective groups. The second is that elements with particularly stable electron configurations, filled or exactly half filled subshells (Be, Mg, Mn Tc, Zn, Cd, N, etc...) also have very low electron affinity.

2.5.4

Electronegativity is a chemical property which describes the power of an atom in a molecule to attract electrons towards itself. Electronegativity is not strictly an atomic property, but rather a property of an atom in a molecule: the equivalent property of a free atom is its electron affinity. It is to be expected that the electronegativity of an element will vary with its chemical environment, but it is usually considered to be a transferable property, that is to say that similar values will be valid in a variety of situations. First proposed by Linus Pauling in 1932, electronegativity cannot be di- rectly measured and must be calculated from other atomic or molecular prop- erties. Several methods of calculation have been proposed and, although there may be small differences in the numerical values of the electronegativ- ity, all methods show the same periodic trends between elements. The most commonly used method of calculation is that originally pro- posed by Pauling. This gives a dimensionless quantity, commonly referred to as the Pauling scale, on a relative scale running from 0.7 to 4.0 (hydrogen = 2.2). When other methods of calculation are used, it is conventional (al- though not obligatory) to quote the results on a scale which covers the same

36 The Elements

Figure 2.11: The electron affinity of elements. The data are quoted in kJ/mole (1kJ/mole = 0.010364 eV/atom). Elements with a value of zero are expected to have electron affinities close to zero on quantum me- chanical grounds. Elements with no values presented are synthetically made elements - elements not found naturally in the environment (from www.dayah.com/periodic). range of numerical values: this is known as an electronegativity in Pauling units. In general, electronegativity increases on passing from left to right along a period, and decreases on descending a group, see Figure 2.12. Hence, fluorine is undoubtedly the most electronegative of the elements while cesium is the least electronegative, at least of those elements for which substantial data are available. There are some exceptions to this general rule. Gallium and germanium have higher than aluminum and silicon respectively because of the d-block contraction. In inorganic chemistry it is common to consider a single value of the elec- tronegativity to be valid for most ”normal” situations. While this approach has the advantage of simplicity, it is clear that the electronegativity of an element is not an invariable atomic property and, in particular, increases

37 Chapter 2

Figure 2.12: Periodic table of electronegativity using the Pauling scale with the of the element. Allred used the Pauling method to calculate separate electronegativities for different oxidation states of the handful of elements (including tin and lead) for which sufficient data was available. However, for most elements, there are not enough different cova- lent compounds for which bond dissociation energies are known to make this approach feasible. This is particularly true of the transition elements, where quoted electronegativity values are usually, of necessity, averages over several different oxidation states and where trends in electronegativity are harder to see as a result.

2.5.5 Other Periodic Tables Since 1869, numerous table designs have been proposed to demonstrate the periodic law. I have collected a couple here.

Physicists Periodic Table • Figure 2.13 shows Timmothy Stowe’s physicists periodic table. This table depicts periodicity in terms of the four quantum numbers. As you move down the planes, the principle quantum number, n, increases. In

38 The Elements

Figure 2.13: The physicists periodic table, by Timmothy Stowe.

each plane the color depicts l and the direction in and out of the page depicts ml (which is labeled m in the legend. ml = 0 is the only choice for l = 0 or s subshell. For the d subshell (pinkish color in the table) ml = 2, 1, 0, 1, 2. The left and right direction depict ms which is labeled− as −s in the caption. Notice that you can only ever have 2 values of ms for any other set of quantum numbers, so for any given plane (n) and color (l), and front to back position (ml) there are only two elements indicated, each with a different ms.

Triangular Periodic Table • Figure 2.14 shows an example is based on the work of Emil Zmaczynski and graphically reflects the process of the construction of electronic shells of atoms.

39 Chapter 2

Figure 2.14: The triangular periodic table by Emil Zmaczynski

Spiral Periodic Table • Figure 2.15 was devised by Theodor Benfey and depicts the elements as a seamless series with the main group elements radiating from the center with the d- and f-elements filling around loops.

The Mayan Periodic Table • Figure 2.16, named for its similarity to the ancient Mesoamerican cal- endar, is based on electron shells. The shells are shown as concentric circles. Each row in the tabular form is shown as a ring. The noble gases are in the vertical column above the center of the chart. The elements to the right of a have one extra electron, the elements directly on the left need one electron to fill their outer shell. What does this chart show better then the traditional table? (1) The reactivity of the elements. The elements closest to the noble gases are so close to becoming a noble gas they can taste it. This makes them more eager, more reactive, since all they need is to gain or lose one electron. As you move away from the noble gases along the concentric circles, the elements get less and less reactive, since they are so far from being a noble gas, they don’t really think it is worth the effort. (2) The proportions of compounds. The proportion of elements can be guessed by looking at the ’hops’ that an element must take to get to the noble

40 The Elements

Figure 2.15: Spiral periodic table.

gases. The guideline is that for elements to combine, one should be from the left and one from the right. The number of hops an atom takes to get to the vertical line must equal the number its partner on the other side takes, since one is gaining an electron and one is losing one. Aluminum needs three hops left to get vertical, and Oxygen needs just two. Since the number of hops needs to be the same on both sides, we need two Aluminum atoms to make the journey and three Oxygen to make it equal. This implies Al(2)O(3) is a good possibility for a compound. This guideline only works for elements fairly near the noble gases. This chart is a only a rough guide since the inner shells of elements are not always filled before the outer shells.

41 Chapter 2

Figure 2.16: Mayan Periodic Table. Radioactive elements are underlined in red. Synthetic elements are underlined in blue. The Noble gases are in bold.

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