Ionisation Energy

6 Key excerpts on "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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  Basic Concepts of Chemistry

  • Leo J. Malone, Theodore O. Dolter(Authors)
  • 2012(Publication Date)
  • Wiley

    (Publisher)

  The ion- ization energy generally increases across a period but decreases down a group. The ionization energies for the second and third periods are shown in Table 8-2. The energy unit is abbreviated kJ/mol, which stands for kilojoules per mole, the energy for a defined quantity of atoms. Notice that the trends in ionization energy follow from the discussion of atomic radius. That is, the smaller the atom, the harder it is to remove an electron. This makes sense according to Coulomb's law. Coulomb's law 152 111 88 77 74 66 64 37 186 160 143 117 110 104 99 227 197 122 122 121 117 114 Li Be B C N O F H Na Mg Al Si P S Cl K Ca Ga Ge As Se Br FIGURE 8-18 Atomic Radii Radii of these atoms are given in picometers (pm). TABLE 8-2 Ionization Energy of Some Elements ELEMENT I.E. (KJ/MOL) Li 520 Be 900 B 801 C 1086 N 1402 O 1314 F 1681 Na 496 Mg 738 Al 578 Si 786 P 1102 S 1000 Cl 1251 262 CHAPTER 8 Modern Atomic Theory states that the forces of attraction increase as the charges increase but decrease as the distance between the changes increases. In smaller atoms, the positive protons and negative electrons are closer together and are thus attracted to each other more strongly. Notice that ionization energy generally decreases to the left and down. It is no coincidence that this is the same direction as increasing metallic properties. In fact, the most significant chemical property of metals is that they lose electrons relatively easily to form cations in compounds. However, some metals have considerably higher ionization energies than others. The chemical reactivity of a metal relates to this energy. For example, the ionization energy of sodium is 496 kJ/mol and of gold is 890 kJ/mol. As a result, sodium is a very reactive metal (i.e., it reacts dramatically with water), yet gold is called the eternal metal because of its relative unreactivity. Nonmetals also have very high ionization energies and so do not form positive ions in compounds.

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  Theoretical Aspects of Chemical Reactivity

  • (Author)
  • 2006(Publication Date)
  • Elsevier Science

    (Publisher)

  Theoretical Aspects of Chemical Reactivity A. Toro-Labbé (Editor) © 2007 Published by Elsevier B.V. Chapter 8 The average local ionization energy: concepts and applications Peter Politzer and Jane S. Murray Department of Chemistry, University of New Orleans, New Orleans, LA 70148, USA 1. Orbital ionization energies The ionization energy I of an N-electron atom or a molecule having energy E is a well-defined property: I = EN − 1 − EN (1) I has been measured experimentally at a high level of accuracy for a large number of atoms and molecules [1]. For molecules, it can be important to distinguish between the adiabatic and the vertical ionization energies. The former corresponds to the neutral molecule and the ion, both being in the ground states, which may or may not have similar geometries. The latter refers to the situation in which the positions of the nuclei are the same in the ion, as in the neutral molecule, which may not be the ion’s ground vibrational state. Our interest shall be in vertical ionization energies. Computationally, either I can be obtained via (1) by evaluating the appropriate EN − 1 and EN . However, another approach is used very frequently. In Hartree– Fock (HF) theory, it follows directly from the formalism that the vertical ionization energy I i of any electron i would equal the negative of its orbital energy i if all of the orbitals of the system were unaffected by the loss of the electron. Koopmans’ theorem assures the stability of the one from which the electron is lost [2,3], and thus the approximation I i ≈ − i (2) is a common one at the HF level. Then, the molecular vertical ionization energy I corresponds to i for the highest occupied orbital. 119 120 Average local ionization energy In reality, the remaining orbitals do undergo some changes when an electron is removed from one of them. By ignoring this, which stabilizes the positive ion, (2) overestimates I i .

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  Chemistry, 5th Edition

  • Allan Blackman, Steven E. Bottle, Siegbert Schmid, Mauro Mocerino, Uta Wille(Authors)
  • 2022(Publication Date)
  • Wiley

    (Publisher)

  As with atomic radius, Ionisation Energy changes less for elements in the d and f blocks, because increased shielding from the d and f orbitals offsets increases in Z. As a rule of thumb, the trend in the ionisation energies is inverse to that of the atomic radii, that is, smaller atoms have higher ionisation energies. Therefore, the trend in the ionisation energies can be rationalised in similar fashion. FIGURE 4.41 The first Ionisation Energy generally increases from left to right across a period (blue arrow) and decreases from top to bottom down a group (red arrow) of the periodic table. 0 10 20 30 40 50 60 70 80 90 First Ionisation Energy, E i1 (kJ mol –1 ) Atomic number, Z 2000 1500 1000 500 0 Rb Xe Rn Ar H Li Na K 2500 He Ne Kr Cs As Z increases, E i1 increases. As n increases, E i1 decreases. 6 5 4 3 2 1 Row He 2372 H 1312 1 2 2000 1500 1000 2500 First Ionisation Energy (kJ mol –1 ) 500 2000 1500 1000 2500 500 0 0 Group 18 13 14 15 16 17 Ne 2080 F 1681 O 1314 N 1402 C 1086 B 800 Be 899 Li 520 Ar 1520 Cl 1256 S 999 P 1012 Si 786 Al 577 Mg 738 Na 496 Kr 1351 Br 1143 Se 941 As 947 Ge 761 Ga 579 Ca 590 K 419 Xe 1170 I 1009 Te 869 Sb 834 Sn 708 In 558 Sr 549 Rb 403 Rn 1037 At (926) Po 813 Bi 703 Pb 715 Tl 589 Ba 503 Cs 376 Higher ionisations A multi-electron atom can lose more than one electron, but ionisation becomes more difficult as positive charge increases. The first three ionisation energies for a magnesium atom in the gas phase provide an illustration. Process Configurations E i Mg(g) ← → Mg + (g) + e - [ Ne ] 3s 2 ← → [ Ne ] 3s 1 738 kJ mol - 1 Mg + (g) ← → Mg 2+ (g) + e - [ Ne ] 3s 1 ← → [ Ne ] 1450 kJ mol - 1 Mg 2+ (g) ← → Mg 3+ (g) + e - [ Ne ] ← → [ He ] 2s 2 2p 5 7730 kJ mol - 1 178 Chemistry Notice that the second Ionisation Energy of magnesium is almost twice as large as the first, even though each electron is removed from a 3s orbital. This is because Z eff increases as the number of electrons decreases.

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  Extreme Physics

  Properties and Behavior of Matter at Extreme Conditions

  • Jeff Colvin, Jon Larsen(Authors)
  • 2013(Publication Date)
  • Cambridge University Press

    (Publisher)

  The energy required to completely strip the atom of its electrons is the sum of these single-particle ionization energies I j . Figure 7.4 shows the successive ionization potentials for a few elements of the periodic table. These data were obtained from a relativistic Hartree–Fock calculation for isolated atoms. Neutral atoms, ions, and free electrons in thermal equilibrium obey the laws of statistical mechanics; the particular distribution function is proportional to the 194 Ionization Figure 7.4 Ionization potentials as a function of the number of bound electrons; elements are C, Al, Fe, Ag, Au. Boltzmann factor e −I/kT , where I is a potential of some form. We see that the ionization process depends upon temperature, and begins at kinetic temperatures much lower than the ionization potential I. The reason ionization starts for low temperature is that the statistical weight of the free electron is very large. Ionization begins sooner the lower the ionization potential. Successive stages of ionization proceed until all of the bound electrons are liberated, but the onset of the next stage of ionization begins before the previous stage ends, and thus the atoms in a volume of gas have a distribution of ionization states, ranging, perhaps, from the neutral atom to the fully stripped atom. The ensemble of the distribution is characterized by an average ionization level, Z ∗ . If the “material” consists of a mixture of elements then it will contain differently charged ions of each element. We use the term “material” in a broad sense; most commonly one thinks of a gas, but it also applies to warm dense matter even though that state may exhibit characteristics of a liquid or solid-state material. Our discussion proceeds assuming we have a simple gas consisting of atoms of a single element. The ionization level Z ∗ is the effective number of free electrons per ion.

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  Chemistry of the Non-Metals

  With an Introduction to Atomic Structure and Chemical Bonding

  • Ralf Steudel, Frederick C. Nachod, Jerry J. Zuckerman, Frederick C. Nachod, Jerry J. Zuckerman(Authors)
  • 2011(Publication Date)
  • De Gruyter

    (Publisher)

  There is a continuous transition between bond types, and in most compounds more than one of these idealized types is present. 5.1. The Ionic Bond 5. L L Introduction A large number of compounds crystallize in ionic lattices consisting of periodic, regular, three-dimensional arrays of anions and cations. The ions may be elemental or complex (derived from compounds), as: Li + and H~ in LiH NO + and HSO 4 in NOHSO 4 Ca 2+ and F~ in CaF 2 H 3 O + and C1O 4 in H 3 OC1O 4 A1 3+ and O 2 ~ in A1 2 O 3 NHJ and BF; in NH 4 BF 4 Stoichiometric ratios of anions and cations achieve electrical neutrality in the crystal. Crystal type, i.e., geometry and symmetry of the lattice, is dictated by the relative sizes of the ions and their charges. Ions are generated from atoms by supplying the ionization energy, I, or by liberating the electron affinity, A. 60 Atomic Structure and Chemical Bonding 5. L 2. lonization Energy The first ionization of a gaseous atom, X: (49) requires the first ionization energy, which is positive, i.e., energy must be supplied to the system. The value of I! is influenced by the position of X in the periodic system. Metal atoms ionize most easily and noble gas atoms with most difficulty, lonization energies range from ca. 4 to 25 eV, or 400 to 2400 kJ/mol. Figure 21 plots the first ionization energy as a function of atomic number. [eV] 26-24-22-Γ I 18-§>16-c .2 12-CO .g 10-o — 8-6-4-2-O Ra 10 20 30 40 50 Atomic Number —·-60 70 80 90 Fig. 21. First ionization energy, I l5 as function of atomic number. The second ionization energy, I 2 , follows from: -e-(50) and is always considerably larger than l i because the electron must now be removed from a particle already bearing a positive charge. The value of I 2 changes periodic-ally with atomic number, like I l5 and maximizes at the ions having the rare gas configuration (e.g., Na + , K + , etc.).

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