Trends in Ionic Charge

3 Key excerpts on "Trends in Ionic Charge"

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  eBook - ePub

  Foundations for Teaching Chemistry

  Chemical Knowledge for Teaching

  • Keith S. Taber(Author)
  • 2019(Publication Date)
  • Routledge

    (Publisher)

  Two other concepts, however, can be linked more directly to a model of atomic structure – at least for the first 20 elements. The periodic table is arranged into rows known as periods and columns known as groups. That geometric notion links to atomic structure (so an element in period 2 has atoms with two occupied shells of electrons; an atom of an element in group 3 has three electrons in its outer shell – and so the element in period 2 and group 3, B, has an electronic configuration of 2.3). This also relates to periodicity. Elements in the same group are said to share common properties, whilst there are often clear trends in properties moving across periods.

  Common group properties

  Those common group properties and periodic trends reflect certain patterns abstracted from a vast catalogue of data. The patterns of stoichiometry, such as for the period 2 chlorides shown earlier, is strong but has limits. In period 3 some elements, even in the p-block, show more than one valency (e.g., PCl3 and PCl5 ; IF, IF3 , IF5 , and IF7; etc.) – due to the ‘expansion of the octet’, a complication that cannot be readily explained with the simple ‘shells’ models of electronic structure. This links to a common alternative conception that students acquire – that electron shells (not just the L shell) become ‘full’ at eight electrons. Unfortunately, this is something implied (or even stated) in some school textbooks (Taber, 1998a).

  Elements in the d-block commonly show multiple valencies, but the actual patterns of valencies found are more difficult to explain. The alkali metals are all relatively soft, low-melting, highly reactive metals; the halogens are coloured non-metals; but then nitrogen, arsenic, and bismuth do not so obviously have much in common. This is not to suggest that chemists are unable to explain patterns using the periodic table, but rather that whilst some patterns can be readily related to the periodic table using a simple ‘shells’ model of atomic structure of the kind commonly met in introductory chemistry, other patterns need to be related to more sophisticated structural models or interpreted in terms of complex patterns of factors (see the example of ionisation energy in Chapter 12 ).

  Trends in the periodic table

  The periodic table subsumes the important chemical concept of trends in elemental properties. The trends concept is used more widely – for example in economics – so students may be familiar with it from outside the chemistry classroom. News reports may refer to trends in unemployment or in industrial production or in professional sportspeople’s earnings. In everyday usage, a trend refers to a pattern of change over time , and periodic or group patterns are only trends by analogy, in the sense that we conventionally read a periodic table according to an assumed ordering (see Figure 9.1

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  eBook - ePub

  Introduction to Geochemistry

  Principles and Applications

  • Kula C. Misra(Author)
  • 2012(Publication Date)
  • Wiley-Blackwell

    (Publisher)

  Fig. 2.8 ).

  2.4 Chemical behavior of elements

  The chemical behavior of an element is governed by its electronic configuration because the energy level of the atom is determined by the spatial distribution of its electron cloud. It is only the most loosely bound electrons in the outermost orbitals that take part in chemical interaction with other atoms. For example, the alkali elements (group IA of the Periodic Table), all of which have one electron in the outermost orbital, exhibit similar chemical properties; so do the alkaline earth metals (group IIA of the Periodic Table), all of which have two electrons in the outermost orbital.

  2.4.1 Ionization potential and electron affinity

  Two concepts are useful in predicting the chemical behavior of elements: ionization potential (or ionization energy) ; and electron affinity . Ions are produced by the removal of electron(s) from or the addition of electron(s) to a neutral atom. The energy that must be supplied to a neutral atom (M) in the gas phase to remove an electron to an infinite distance is called the ionization potential (I) . In other words, the ionization potential is the difference in potential between the initial state, in which the electron is bound, and the final state, in which it is at rest at infinity; the lower the ionization potential, the easier it is to convert the atom into a cation. This is the reason why the ionization potential generally increases from left to right in a given period and from top to bottom within a given group of the Periodic Table (Fig. 2.9 ). The first ionization potential (I 1 ) refers to the energy required to remove the first (the least tightly bound) electron, the second ionization potential (I 2

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  Introduction to Modern Inorganic Chemistry, 6th edition

  • R.A. Mackay(Author)
  • 2017(Publication Date)
  • CRC Press

    (Publisher)

  Table. 2.14 that these increase towards the right of the Periods and decrease down the Groups. In addition, they reflect the other, smaller, variations which have been remarked; for example, the changes in the Main Groups are more pronounced than in the Transition Groups, and the discontinuity in properties of the elements from gallium to bromine when compared with the rest of their respective Groups is reflected in their electronegativity values.

  8.5 Chemical behaviour and periodic position

  The detailed chemistry of the elements is discussed in the succeeding chapters. In this section, the skeleton of the periodic properties is outlined to provide a framework for the more detailed account which follows.

  Those elements where the outermost electrons are in a new quantum level, after a rare gas configuration, normally react by losing these loosely bound electrons and forming cations. This mode of behaviour is typical of the elements of the lithium, beryllium and scandium Groups together with the lanthanide elements, which have the respective valency shell configurations, s1 , s2 and d1 s2 . All these elements, with the exception of beryllium itself, lose these outer electrons completely with the formation of cations; M+ in the lithium Group, M2+ in the beryllium Group, and M3+ for scandium, yttrium and the lanthanides.

  The elements of the boron, carbon, nitrogen, oxygen, and fluorine Groups, where the outermost electrons are in p orbitals, show more complicated behaviour.

  (a) Elements with electron configurations close to the rare gases can acquire electrons to form anions with complete rare gas shells. Thus the elements of the halogen Group all form X– ions, and we also find stable compounds containing O2− , S2− and N3−

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