Trends in Ionisation Energy

5 Key excerpts on "Trends in Ionisation Energy"

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  eBook - ePub

  Chemical Fundamentals of Geology and Environmental Geoscience

  • Robin Gill(Author)
  • 2014(Publication Date)
  • Wiley-Blackwell

    (Publisher)

  Although the architecture of the Periodic Table can be thought of as an outcome of wave-mechanical theory, it was originally worked out from chemical observation. It was first published in its modern form by the Russian chemist Dimitri Mendeleev in 1869, almost 60 years before Schrödinger published his paper on wave mechanics.

  Ionization energy

  The bonds formed by an atom involve the transfer or sharing of electrons. It therefore makes sense to illustrate the periodicity of chemical properties by looking at a parameter that expresses how easy or difficult it is to remove an electron from an atom. The ionization energy of an element is the energy input (expressed in J mo1−1 ) required to detach the loosest electron from atoms of that element (in its ground state). It is the energy difference between the ‘free electron at rest' state (the zero on the scale of electron energy levels) and the highest occupied energy level in the atom concerned. What this means in the simplest case, the hydrogen atom, is shown in Figure 5.6 . A low ionization energy denotes an easily removed electron, a high value a strongly held one.

  We can picture how ionization energy will vary with atomic number by considering the highest occupied energy level in each type of atom (Figure 5.7 ). In lithium (Li; Z = 3, electronic configuration = ls2 2s1 ) and beryllium (Be; Z = 4, 1s2 2s2 ) it is the 2 s level; in boron (B; Z = 5, ls2 2s2 2p1 ) it is the 2p level; and so on. If we were to disregard the increasing nuclear charge, we would predict that the energy needed to strip an electron from this ‘outermost' level would vary with atomic number as shown in Figure 6.1 a. One would expect a general decline in ionization energy with increasing Z, punctuated by sudden drops marking the large energy gaps between one ‘shell' and the next one up (Figure 5.6 ); the downward series of steps in Figure 6.1 a thus reflects the occupation of progressively higher energy levels in Figure 5.6 . There is no suggestion of periodicity.

  Figure 6.1

  (a) A notional plot of ionization energy against atomic number, predicted without regard to the effect of increasing nuclear charge. (b) The variation of measured ionization energy with atomic number Z among the first 20 elements. (The whole Z

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  eBook - ePub

  General Chemistry for Engineers

  • Jeffrey Gaffney, Nancy Marley(Authors)
  • 2017(Publication Date)
  • Elsevier

    (Publisher)

  2.15 shows a pattern opposite to that seen for atomic radius. Ionization energy shows a pattern with a maximum followed immediately by a minimum and a steady increase back to a maximum. So, the periodic trend in ionization energies is opposite to that observed with the atomic radii. The highest ionization energies are seen for the noble gases in group 18 (atomic numbers 2, 10, 18, 36, 54, and 86), while the lowest ionization energies are seen for the alkali metals in group 1 (atomic numbers 3, 11, 19, 37, 55, and 87). So, the ionization energy increases from left to right across a period. Although opposite to the trend observed in atomic radius, the trend in ionization energy occurs because of the same electronic effects responsible for the trend in atomic radius.

  Ionization energy tends to increase as you move across a period from the alkali metals to the noble gases because the increasing number of protons in the nucleus results in a stronger attraction for the valence electrons. This stronger attraction increases the energy required to remove a valence electron from the valence shell. Also, it has been observed that an atom whose valence shell contains the maximum number of electrons allowed by the Pauli Exclusion Principle, called a closed shell , is especially stable. The alkali metals in group 1 have the lowest ionization energies because their single electron in the valence shell requires less energy to remove. The noble gases have very high ionization energies because their valence shells are completely full and the nuclear attraction for the valence electrons is at a maximum. Notice that helium (atomic number 2) has the highest ionization energy of all the elements.

  Both the maxima (group 18) and minima (group 1) decrease as atomic number increases. So, the ionization energy decreases from top to bottom in a group. As atomic number increases in a group, the electrons are added to electron shells farther from the nucleus (increasing n

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  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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  Biophysical Basis of Physiology and Calcium Signaling Mechanism in Cardiac and Smooth Muscle

  • Tetsuya Watanabe(Author)
  • 2018(Publication Date)
  • Academic Press

    (Publisher)

  p -orbital, tending to attract one more electron from another atom, and then drop to a lower energy state. Helium, neon, and argon have the filled shells. They are inert gases.

  1. (1)  
     Periodic trends of ionization energy: IE is the minimum energy required to remove an electron from the highest occupied atomic orbital unless otherwise specified. Across a row to the right, IE increases since Z increases but n stays the same. The outermost electron is pulled more to the nucleus by the increased number of its protons. As moving a column down, IE decreases. Z increases as going down a column, but larger n suggests that electrons are farther away from the nucleus. A large distance dominates over the increased Z .
  2. (2)  
     Periodic trends of electron affinity (EA ): EA shows the ability of an atom to gain electrons. Chlorine (Cl) tends to gain one electron from another atom and becomes Cl− 1 ion which is more stable than Cl. Acquisition of one electron by Cl (gas) requires − 364 kJ/mol. The minus sign reflects energy release. EA of Cl is 364 kJ/mol. As moving across a row to the right, EA increases. As going down a column, EA decreases.
  3. (3)  
     Periodic trends of electronegativity (χ ): Electronegativity is the ability of an element to attract electrons toward itself from another element. Linus Pauling first proposed the idea in 1932. Mulliken later proposed electronegativity as the average of IE and EA [4] ,
     χ =

     1 2

     IE + E A
       (3.8)
     It is usual to use a linear transformation to transform these absolute values into values that resemble the more familiar Pauling values. For ionization energies and electron affinities in electron volts [5] ,
     χ = 0.187

     IE + E A

     + 0.17
     and for energies in kilojoules per mole,
     χ = 1.97 ×

     10

     − 3

     IE + E A

     + 0.19
     The Mulliken electronegativity can only be calculated for an element for which the electron affinity is known.

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  eBook - ePub

  Introduction to Modern Inorganic Chemistry, 6th edition

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

    (Publisher)

  IG . 8.4 Variations of first ionization potential across the first short Period

  As the ionization potentials measure the energy required to remove the least tightly bound electron from an atom or ion, values reflect the stability of the configuration from which the electron is being removed. Table. 2.8 gives the ionization potentials of the elements.

  The energy gaps between successive levels with the same l value decrease as the n values increase, so that all the atomic orbitals get closer in energy as the atomic number increases. This trend is not completely regular, and larger than average energy gaps occur between the 4p and 5s levels where the first set of d orbitals has been filled, and between the 6s and 7s levels where the first of the f levels comes. These energy jumps reflect the poorer-than-average shielding powers of d and f electrons.

  Apart from the major discontinuities at the rare gases, there is also a gap in energy wherever the outermost electron enters a new atomic orbital. These gaps correspond to stabilization of the filled shell configurations, s2 , d10 s2 and f14 d10 s2 , before the p orbitals are occupied, and also suggest the possibility of transfer of s electrons into the d shell to give the d10 configuration, or of d electrons into the f shell to give the f14 arrangement, which was discussed in the previous section.

  The stability of the rare gas configurations can be seen, both from the high energies required to remove an electron from the rare gases themselves, and from the leap in the values of the potential when the rare gas configuration has to be broken (i.e. when the second electron is removed from an alkali metal, the third electron is removed from an alkaline earth, the fourth electron from a boron Group element, etc.). The very low first ionization potentials of the alkali metals, and, to a lesser extent, the low first and second potentials of the alkaline earths, show how loosely held are the first one or two electrons outside the rare gas configuration.

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