P.A. Cox - Inorganic chemistry (793955), страница 5
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Older (and contradictory) numbering systemsare still found. Some groups of elements are conventionally givennames, the most commonly used being alkali metals (group 1),alkaline earths (2), halogens (17) and noble gases (18).Many-electron atoms (A3)Trends in atomic properties(A5) Chemical periodicity(B2)HistoryAs more elements were discovered in the 19th century chemists started to note similarities in their properties.
Earlyattempts to order the elements in a regular fashion were hampered by various difficulties, especially the fact (only laterrealized) that atomic masses do not increase regularly with atomic number. Mendeleev published the first satisfactoryform of the periodic table in 1869, and although many details of layout have evolved since then, his basic idea has beenretained, of ordering elements horizontally in periods so that they fall in vertical groups with similar chemical properties.Mendeleev was forced to leave some gaps for elements not yet discovered, and his ability to predict their propertiesvindicated his approach.16A4—THE PERIODIC TABLEThe first satisfactory determination of atomic number (as opposed to atomic mass) came from Moseley’s studies ofX-ray spectra in 1917. By determining the wavelength, and hence frequency, of X-rays emitted from differentelements, Moseley observed different series of X-ray lines.
In each series the frequency (ν) of each line varied with atomicnumber (Z) according to the formula(1)where C and σ are constants for a given series. Moseley’s law can be understood from the quantum theory of manyelectron atoms. X-rays are produced when atoms are bombarded with high-energy electrons. These knock outelectrons from filled orbitals, thus providing ‘vacancies’ into which electrons can move from other orbitals and emit Xray photons. Different series of lines come from different vacancies; for example, the highest-energy K series is excitedwhen a 1s electron is removed.
Equation 1 then expresses the energy difference between two types of orbital, with Cdepending on the values of n involved, and σ on the screening constants (see Topic A3).Using Moseley’s law allowed the remaining uncertainties in the structure of the periodic table to be resolved. Atabout the same time the theoretical ideas of the quantum theory allowed the structure of the table to be understood.Bohr’s aufbau (or building up) principle (see below) was developed before the final version of the theory wasavailable; following Schrödinger’s equation (1926) the understanding was complete. The periodic table with itstheoretical background remains one of the principal conceptual frameworks of inorganic chemistry. A complete table isshown inside the front cover of this book.Building upAccording to the aufbau principle, the ground-state electron configuration of an atom can be found by putting electronsin orbitals, starting with that of lowest energy and moving progressively to higher energy.
It is necessary to take intoaccount both the exclusion principle and the modification of orbital energies by screening and penetration effects (seeTopic A3). Thus following He (1s)2, the electron configuration of Li is (1s)2(2s)1, as the 2s orbital is of lower energythan 2p. Following Be, the 2p orbitals are first occupied in B (see Table 1). A total of six electrons can be accommodatedin these three orbitals, thus up to Ne.Table 1. Electron configuration of ground-state atoms up to K (Z=19)SECTION A—ATOMIC STRUCTURE17Following completion of the n=2 orbitals, 3s and then 3p shells are filled. The electron configurations of theelements Na-Ar thus parallel those of Li-Ne with only a change in the principal quantum number n.
An abbreviated formof the configurations is often used, writing [He] for the filled configuration (1s)2 and [Ne] for [He](2s)2(2p)6. The innershell orbitals denoted by these square brackets are too tightly bound to be involved in chemical interactions: it is thevalence or outer electrons that determine chemical properties.
The group structure of the periodic table depends onthe fact that similar outer electron configurations are reflected in similar chemical behaviour.It might be expected that 3d orbitals would fill after 3p, but in fact this does not happen, because the extrapenetration of s compared with d orbitals significantly lowers the energy of 4s. This fills first, so that following Ar thefirst two elements of the fourth period K ([Ar](4s)1) and Ca ([Ar](4s)2) have configurations parallel to Na and Mg,respectively.
The 3d orbitals then fill, giving the 10 elements Sc-Zn, followed by 4p. The fifth period follows similarly,5s, 4d then 5p. In the sixth period another change takes place, with filling of the 4f shell after 6s and before 5d. Theseventh incomplete period begins with 7s followed by 5f and would be expected to continue in the same way, but theseelements become increasingly radioactive and hard to make or study (see Topic I2).The order of filling of shells is conveniently summarized in Fig. 1.
It is important to note that it reflects the order ofenergies at the appropriate point, and that this order changes somewhat as more electrons are added. Thus followingcompletion of the 3d shell, increasing atomic number stabilizes these orbitals rapidly so that they are no longerchemically active; in an element such as Ga ([Ar](3d)10(4s)2(4p)1) the valence orbitals are effectively only the 4s and 4p,so that its chemistry is similar to that of Al ([Ne](3s)2(3p)1).
The same is true following completion of each d and f shell.Block structureThe filling of the table described above leads to a natural division of the periodic table into blocks according to the outerelectron configurations of atoms (see Fig. 2). Elements of the s block all have configurations (ns)1 or (ns)2. In periods 2and 3 these are followed immediately by the p block with configurations (ns)2(np)x. Lower p block elements are similaras the (n−1)d orbitals are too tightly bound to be chemically important. The s and p blocks are collectively known asFig. 1. Showing the order of filling of orbitals in the periodic table.18A4—THE PERIODIC TABLEFig.
2. Structure of the periodic table, showing the s, p, d and f blocks.main groups. d-block elements of periods 4, 5 and 6 have ns and (n−1)d outer electrons, and are known astransition elements. Their configurations show some complexities as the s and d orbitals are similar in energy (seeTopic H1). The f-block elements are known as the lanthanides (4f) and actinides (5f). For ease of presentationthey are generally shown as separate blocks below the main table. In the case of the lanthanides, this procedure ischemically justified as the elements have very similar properties (see Topic I1).Group numbers and namesThe numbering of groups in the periodic table has a confused history reflecting developments in understanding andpresenting the table itself. In the current nomenclature used in this book, groups are numbered 1–18, with thelanthanides and actinides all subsumed into group 3.
Older numberings based on 1–8 are still found, with a division intoA and B subgroups which unfortunately differs according to the continent. In the UK, the s- and early d-block elementsare numbered 1A–8A (the last encompassing modern group numbers 8, 9 and 10), followed by numbers 1B (now 11)to 8B. In the USA, 1A–8A refer to main groups, with d-block elements numbered B. This confusion is resolved by thenewer system.Some groups of elements are conventionally given names. Group 1 elements (not hydrogen) are called alkalimetals and those of group 2 alkaline earths.
Groups 17 and 18 are the halogens and noble gases, respectively.Sometimes group 16 are called chalcogens although this normally excludes the first element oxygen: thus the termchalcogenide refers to compounds with sulfur, selenium and tellurium. Lanthanides were previously called rareearths; although the term is no longer used by chemists it is still common in geochemistry (where it often includesyttrium in group 3 in the previous period, not a lanthanide but chemically very similar).Section A—Atomic structureA5TRENDS IN ATOMIC PROPERTIESKey NotesEnergies and sizesHorizontal trendsVertical trendsStates of ionizationRelativistic effectsRelated topicsTrends in orbital energy and size reflect changes in the principal quantumnumber and effective nuclear charge.
They are seen experimentally intrends in ionization energy (IE) and apparent radius of atoms.Increasing nuclear charge causes a general increase of IE and a decrease ofradius across any period. Breaks in the IE trend are found following thecomplete filling or half filling of any set of orbitals.A general increase of radius and decrease in IE down most groups isdominated by the increasing principal quantum number of outer orbitals.Effective nuclear charge also increases, and can give rise to irregularities inthe IE trends.IEs for positive ions always increase with the charge.
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