(G S Manku, email: manku.gs@gmail.com)
(4.2).1 Moseley’s Characteristic X Rays
A major step towards assigning a fundamental number Z to elements was the discovery of characteristic X rays of the elements. When a target of an element is bombarded with high energy electrons, two types of rays are emitted. One type is a continuous distribution of wavelengths [Figure (4.2).1 (1)]. Its high energy end depends upon the energy of the incident electrons. The other type superimposed on this is called the characteristic X ray spectrum of the element (Figure (4.1).1 (2)]. It consists of a few sharp peaks, wavelength of which depends upon the target element and is independent of the energy of the incident electrons. These are emitted only if the energy of the incident electrons is above a certain minimum value that depends on the target element
1 2
Figure (4.2).1 (1) The continuous X ray spectrum of Mo; (2) The characteristic X rays superimposed on the
continuous X rays for Mo. (3) Graph of frequency ν of the
characteristicX rays against the atomic number Z.
Barkla and Sadler (1908 – 1911) showed the presence of two types of characteristic X rays – called the K rays with a higher penetration power than the other one, called the L rays. Same characteristic X rays can be obtained by using an anticathode of containing the element and then passing the emitted rays through a screen made of the same element to absorb the extraneous radiations.
Moseley’s Law: Moseley in 1919 studied the characteristic X rays of 38 elements using a crystal of potassium ferrocyanide as diffraction grating and a photographic screen developed by Bragg. Each ray of Barkla and Sadler was found to consist of two rays (e.g., Kα and Kβ for K rays). The graphs of square root of frequency of characteristic rays ( √ν) against the atomic number Z of the element as determined by Rutherford resulted in straight lines [Figure (4.1).1.(3)] for aluminium (Z = 13) to silver (Z – 47) for K lines, and from zirconium (Z = 40) to gold (Z = 79) for the L lines respectively. Hence, Moseley proposed that
√ν = a (Z – b)
where a is the proportionality constant (= 1 for K line and 7.4 for L line) and b is same for all the lines of a given series Figure (4.2).2. No such relation was observed when Z was replaced by atomic mass A of the elements.
Figure (4.2).2 The graphs of √ν against Z for molybdenum
Therefore, Moseley concluded that nuclear charge, i.e. atomic number or a property closely related to atomic number Z is a fundamental property of an element. Later, other group of characteristic X rays frequencies (like M, N, O series) were also discovered. For all the series, Moseley’s equation is broadly valid. The discrepancies have been of great significance for elucidating the details of atomic structure of elements.
(4.2).2 Origin of Characteristic X Rays
The Bohr’s theory can be used easily to explain the origin of characteristic X rays. The high energy electron just knocks off an electron from the K shell of the target atom. To fill up vacancy in the K shell, an electron from the L shell falls into it, and create a vacancy in the L shell. To fill up this vacancy, an electron from M shell falls to L shell, and a cascade effect ensue, and characteristic X rays are observed in the emitted radiations.
From Bohr’s theory, the frequency of radiation emitted (ν) when the electron from the ni-th shell falls to nf-th shell is given by
ν = (μe4Z2 / 8ε02h2) (1/nf2 – 1/ni2)
where μ is the reduced mass of the electron having the charge e, Z is the nuclear charge, ε0 is the permittivity of vacuum and h is the Planck’s constant. For an electron in a shell, the charge on the nucleus as felt by it will be less than the actual nuclear charge because of the shielding effect of other electrons. Replacing Z by Z – b we get
ν = (μe4(Z – b)2 / 8ε02h2) (1/nf2 – 1/ni2)
For a given line, the term
ν = (μe4 / 8ε02h2) (1/nf2 – 1/ni2) = a2
is a constant. Representing it by a2, the equation reduces to
ν = a2 (Z – b)2
so that we get the Moseley’s equation:
√ν = a (Z – b)
(4.2).3 The Periodic Law
Moseley’s work on the characteristic X rays of elements showed that the atomic number Z, the number of protons in the nucleus of an atom, is a fundamental characteristic property of the elements.
The number of electrons in the extranuclear part of an atom is also equal to Z. As properties of the elements and their compounds depend upon the arrangement of electrons in their extranuclear part, the properties of the elements strongly depend on their atomic number, Z.
Also, for all the successive elements, Z increases exactly by one unit.
According to the periodic law, the properties of elements are a periodic function of their atomic numbers.
(4.2).4 TYPES OF ELEMENTS
Based on the electronic configuration of the incompletely filled shells of electrons, the elements are classified as (1) Main period representative or R elements, (2) Transition or T elements, and (3) Inner-transition or IT elements.
(4.2).4.1 REPRESENTATIVE ELEMENTS
In representative elements or the R families, only the outermost shell has incompletely filled orbitals. Thus, only the outermost n-th shell is the valence shell of R elements. These elements are further classified as:
- s-Block Elements, having configuration of outermost shell of the atom as n1 or ns2. These elements belong to group I A and II A of Mendeleev’s or classical periodic table and groups 1 and 2 of the long periodic table.
- p-Block Elements, having electronic configuration of the outermost valence shell of atom as ns2npx, where x = 1, 2, 3, 4, 5 or 6. In the long periodic table, these are groups 13 to 18 elements (group number = 10 + y, (where y = 2 + x) is the total number of electrons in the outermost shell). These are also called R-3 to R-7 and R-0 families.
Table (4.2).1 The R Elements
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Electronic configuration Group number Family name
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ns1 Group 1 R-1 Alkali metals
ns2 Group 2 R-2 Alkaline earth metals
ns2np1 Group 13 R-3 Boron family
s2np2 Group 14 R-4 Carbon family
ns2np3 Group 15 R-5 Nitrogen family
ns2np4 Group 16 R-6 Oxygen family
ns2np5 Group 17 R-7 The halogens
ns2np6 Group 18 R-0 Group 0, Noble gases
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Strictly speaking, in the atoms of copper and zinc families, ytterbium and nobelium also, only the outermost shell is incomplete, all other underlying subshells are completely filled up with electrons and should also be considered as representative elements:
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Cu Z =29 [Ar] 3s23p63d10 4s1 Group 11 Transition element
Ag Z =47 [Kr] 4s24p64d10 5s1 Group 11 Transition element
Au Z =79 [Xe] 4f14 5s25p65d10 6s1 Group 11 Transition element
Zn Z =30 [Ar] 3s23p63d10 4s2 Group 12 Transition element
Cd Z =48 [Kr] 4s24p64d10 5s2 Group 12 Transition element
Hg Z =80 [Xe] 4f14 5s25p65d10 6s2 Group 12 Transition element
Yb Z = 70 [Xe] 4f14 6s1 4f block Inner transition element
Lu Z = 71 [Xe] 4f14 6s2 5f block Inner transition element
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However, as the copper and zinc families have a close resemblance to the transition elements, they are considered as transition elements. Ytterbium and nobelium are similar to other inner transition elements, and are considered inner transition elements.
On the basis of their chemical reactivity, the s2p6 elements are called group zero or noble elements (formerly called “the inert elements”) because of their extremely low reactivity. The remaining elements however, are called the R-1 to R-7 families as before.
(4.2).4. 2 TRANSITION ELEMENTS
The transition elements have incompletely filled d orbitals in their ground state or in chemically combined state. The configuration of the incompletely filled electronic subshells of these atoms is:
(n – 1)s2(n – 1)p6 (n – 1)dxns2 (x = 1-10)
Thus, atoms of transition elements have two outermost incompletely filled shells, with principal quantum numbers = n and n – 1. Therefore, their valence shells consist of electrons present in two outermost electron shells and these elements can use the (n – 1)d, ns and np electrons for bonding in chemical compounds.
Four series of transition elements are known. These correspond to the filling up of the energy levels corresponding to 3d, 4d, 5d, and 6d orbitals respectively:
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Series Electronic configuration Elements Atomic number Z
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3d [Ar] 3d{1-10}4s1,2 Sc (21) – Zn (30)
4d [Kr] 4d{1-10}5s1,2 Y (39) – Cd (48)
5d [Xe] 4f14 5d{1-10}6s1,2 La (57), Hf (72) – Hg (80)
6d [Rn] 5f14 6d{1-10}7s1,2 Ac (89), Rf (104) – Og (118)
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Strictly speaking, the atoms of lutetium (Z = 71) and nobelium (Z = 103) have completely filled f14 subshell and have electronic configuration similar to those of the transition elements:
Lutetium Z = 71 [Xe] 4f14 5d1 6s2 Completely filled 4f orbitals
Nobelium Z = 103 [Rn] 5f14 6d1 7s2 Completely filled 5f orbitals
Yet they are classified as inner transitions elements because of a close resemblance between the properties of these elements and the preceding inner transition elements.
The transition elements have similar properties: high melting and boiling points, malleability and ductility, tensile strength, metallic character, formation of alloys, coordination compounds, high as well as low oxidation states, formation of paramagnetic and colored ions, polynuclear and metallic clustered compounds, etc. Detailed discussion is in Coordination Compound and Transition elements.
(4.2).4.3 INNER TRANSITION ELEMENTS
The inner transition elements correspond to the filling of electrons in the $f$ orbitals of the atoms.The electronic configuration of these elements can be represented as
(n – 2)s2 (n – 2)p6(n – 2)d10 (n – 2)f(1-14)
(n – 1)s2 (n – 1)p6(n – 1)d1 ns2
The atoms of these elements have three outermost incompletely filled shells, with principal quantum numbers = n, n – 1 and n – 2. Their valence shells, therefore, consist of electrons in these three outermost shells.
Only two series of inner transition elements are known. These correspond to the filling up of the 4f and 5f orbitals, and contain fourteen elements each:
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Name of series Elements and Atomic numbers (Z)
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4f Lanthanides or lanthanons Ce (58) to Lu (71)
5f Actinides or actinons Th (90) to Lr (103)
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As the successive electrons in the lanthanide series are filled up in the low energy inner 4f subshell, they generally, do not participate in chemical combinations. As a result, the lanthanides have almost similar properties and show a regular gradation throughout the series.
In the 5f or the actinide series, however, due to larger atomic sizes and considerably lower energy differences between the 5f, 6d and 7s orbitals, the actinide elements up to Z = 96 can use their 5f as well as 6d electrons for chemical bonding. Actually speaking, thorium (90), protactinium (91) and uranium (92) resemble transition elements titanium, vanadium and chromium respectively. As Z increases, the energy of the lower shell 5f electrons decreases more as compared to that of higher shell 6d electrons, and the relative stability of 5f electrons increases. As a result, the latter 5f elements resemble the lanthanide elements of 4f series.
All the inner transition elements show a uniform valence state of III by losing outermost shell ns and next inner shell electrons. However, due to the additional stability of completely vacant f0, half-filled f7 and completely filled f14 configurations of the f subshell, the elements having one electron less or one electron more show additional states of II and IV respectively also.
(4.2).4.4 Prediction of New Elements and Compounds
Newlands (1864) was first to predict the existence of a missing element between silicon and tin with atomic mass of 73. (Presently, accepted value for atomic mass of germanium, discovered by Winkler in 1886, is 72.59.) Mendeleev’s predictions, made in 1869–1871, were more exhaustive and indicate the depth of his understanding. Of the 26 undiscovered elements of his time between hydrogen and uranium,
Eleven were lanthanides, which are very similar to one another, and were established in 1871 to 1907.
Five are noble gases (helium, neon, argon, krypton, and xenon) present in small amounts in air and isolated around 1894–1898 .
The four radioactive elements were discovered: polonium (Marie Curie, 1898), radium (M Curie and P Curie, 1898), actinium (A Debierne, 1899) and radon (Dorn, 1900).
(4.2).4.5 The super-heavy elements with atomic numbers > 103
A group of trans-uranium elements was synthesized by artificial nuclear reactions during 1940 onwards. The 6d transition series starts with Z = 104 (rutherfordium, Rf) and is complete at Z = 112 (copernicium Cn), Other members of the series are Z = 105 (dubnium, Db), Z = 106 (seaborgium, Sg), Z = 107 (Bohrium, Bh), Z = 108 (hassium, Hs), Z = 109 (meitnerium, Mt), Z = 110 (darmstadtium, Ds) and Z = 111 (roentgenium, Rg). These elements are increasing unstable with respect to α decay or spontaneous fission with half-life of the order of less than 1 sec. It is unlikely that much chemistry can be developed for these elements, though their physicochemical constants have all been predicted.
After this, there follow the six 7p elements for Z = 113 to 118. These are (with their atomic numbers): 113 (nohonium, Nh), 114 (flevonium, Fl), 115 (moscovium, Mc), 116 (livermanium, Lv), 117 (tenessium Ts) and 118 (oganesson, Og). On the basis of present theories of nuclear stability, the 298114Fl may be stable due to island of nuclear stability (number of protons = 114, a magic number for nuclear stability) with a long half-life.
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