P.A. Cox - Inorganic chemistry (793955), страница 6
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Electron affinities arethe IEs of negative ions and are always less than IEs for neutral atoms.Deviations from the nonrelativistic predictions become significant for heavyatoms, and contribute to especially high IEs for later elements in the sixthperiod.Many-electron atoms (A3)The periodic table (A4)Chemical periodicity (B2)Energies and sizesThe first ionization energy (IE) of an atom (M) is the energy required to form the positive ion M+:The IE value reflects the energy of the orbital from which the electron is removed, and so depends on the principalquantum number (n) and effective nuclear charge (Zeff; see Topic A3):(1)The average radius of an orbital depends on the same factors (see Topic A2):(2)Smaller orbitals generally have more tightly bound electrons with higher ionization energies.20A5—TRENDS IN ATOMIC PROPERTIESIt is sometimes useful to assume that the distance between two neighboring atoms in a molecule or solid can beexpressed as the sum of atomic or ionic radii.
Metallic, covalent or ionic radii can be defined according to the typeof bonding between atoms, and van der Waals’ radii for atoms in contact but not bonded. Such empirically derivedradii are all different and are not easily related to any simple predictions based on isolated atoms. They are, however,qualitatively related to orbital radii and all follow the general trends discussed below (see, e.g. Topic D4, Table 1. forionic radii).Horizontal trendsIncreasing nuclear charge is accompanied by correspondingly more electrons in neutral atoms.
Moving from left to rightin the periodic table, the increase of nuclear charge has an effect that generally outweighs the screening from additionalelectrons. Increasing Zeff leads to an increase of IE across each period, which is the most important singletrend in the periodic table (see Topic B2). At the same time, the atoms become smaller.As illustrated for the elements Li-Ne in Fig. 1. the IE trend across a period is not entirely regular.
Irregularities can beunderstood from the electron configurations involved (see Topics A3 and A4). Ionization of boron removes an electronfrom a 2p orbital, which is less tightly bound than the 2s involved in lithium and beryllium. Thus the IE of B is slightlyless than that of Be. Between nitrogen and oxygen, the factors involved in Hund’s rule are important.
Up to three 2pelectrons can be accommodated in different orbitals with parallel spin so as to minimize their mutual repulsion. For O(2p)4 and subsequent elements in the period some electrons are paired and repel more strongly, leading to IE values lessthan would be predicted by extrapolation from the previous three elements.The trends shown in Fig. 1 are sometimes cited as evidence for a ‘special stability’ of filled and half-filled shells.
Thisis a misleading notion. The general increase of IE across a period is entirely caused by the increase of nuclear charge.Maxima in the plot at filled shells (2s)2 and half-filled shells (2p)3 occur only because of the decrease after these points.It is the exclusion principle that controls such details, by forcing the next electron either to occupy another orbital type(as in boron) or to pair up giving a doubly occupied orbital (as in oxygen).Fig.
1. Ionization energies (IE) and electron affinities (EA) for the elements Li-Na.SECTION A—ATOMIC STRUCTURE21Vertical trendsThe IE generally decreases down each group of elements. Figure 2 shows this for hydrogen and the elements of group 1,all of which have the (ns)1 outer electron configuration. The main influence here is the increasing value of principalquantum number n. The fall in IE is, however, much less steep than the simple hydrogenic prediction (1/n2; seeTopic A2). There is a substantial increase of nuclear charge between each element, and although extra inner shells areoccupied, they do not provide perfect shielding.
Thus, contrary to what is sometimes stated, effective nuclearcharge increases down the group. In the resulting balance between increasing n and increasing Zeff (see Equation1) the former generally dominates, as in group 1. There is, however, nothing inevitable about this, and there areoccasions in later groups where Zeff increases sufficiently to cause an increase of IE between an element and the onebelow it.Figure 2 also shows the group 11 elements Cu, Ag and Au, where an ns electron is also being ionized. The increase ofIE along period 4 between K (Z=19) and Cu (Z=29) is caused by the extra nuclear charge of 10 protons, partly shieldedby the 10 added 3d electrons.
A similar increase occurs between Rb and Ag in period 5. In period 6, however, the 4fshell intervenes (see Topic A4) giving 14 additional elements and leading to a total increase of Z of 24 between Cs andAu. There is a much more substantial increase of IE therefore, and Au has a higher IE than Ag. (Relativistic effects alsocontribute; see below.) Similarly irregular trends in IE may have some influence on the chemistry of p-block elements(see Topics F1 and G1).Fig.
2. Ionization energies for elements with (ns)1 outer electron configurations.22A5—TRENDS IN ATOMIC PROPERTIESOrbital radii also depend on n2 and generally increase down each group. Because the radius depends on Zeff and noton(see Equation 2) irregular changes in this quantity have less influence than they do on IEs.
(See, however,transition metals, Topics H1 and H5).There is another interesting feature of vertical trends, arising also from the way in which the periodic table is filled.For orbitals of a given l there is a more significant change, both in IE and size, between the first and second periodsinvolved than in subsequent cases. Figure 2 illustrates this for s orbitals, where the IE decreases much more fromhydrogen (1s) to lithium (2s) than between the lower elements.
Such a distinction is reflected in the chemical propertiesof group 1 elements, hydrogen being nonmetallic and the other elements metals (see Topic B2). Similar, although lessdramatic, differences are found with 2p and 3d. Thus period 2 p-block elements are in many ways different from thoselower in the p block, and 3d series elements distinct from those of the 4d and 5d series.States of ionizationThe successive energies required to create more highly charged ions, M2+, M3+ …are the second, third,…IEs. Thevalues always increase with the degree of ionization.
When electrons are removed from the same shell, the maineffect is that with each successive ionization there is one less electron left to repel the others. The magnitude of thechange therefore depends on the size of the orbital, as electrons in smaller orbitals are on average closer together andhave more repulsion. Thus with Be (2s)2 the first two IEs are 9.3 and 18.2 eV, whereas with Ca (4s)2 the values are 6.1and 11.9 eV, not only smaller to start with (see above) but with a smaller difference. The third IE of both elements isvery much higher (154 and 51 eV, respectively) because now the outer shell is exhausted and more tightly bound innershells (1s and 3p, respectively) are being ionized.
The trends are important in understanding the stable valence states ofelements.The electron affinity of an atom may be defined as the ionization energy of the negative ion, thus the energyinput in the process:although some books use a definition with the opposite sign. Electron affinities are always less than ionization energiesbecause of the extra electron repulsion involved (see Fig. 1). As with successive IEs, the difference depends on theorbital size.
Some apparently anomalous trends can be understood in this way. For example, although the IE of F isgreater than that of Cl (17.4 and 13.0 eV, respectively) the electron affinity of F is smaller (3.4 eV compared with 3.6eV) partly because the smaller size of F− provides more repulsion from the added electron.Some atoms have negative electron affinities, meaning that the negative ion is not stable in the gas phase.
Second andsubsequent electron affinities are always negative because of the high degree of repulsion involved in forming a multiplycharged negative ion. Thus the O2− ion is not stable in isolation. This does not invalidate the ionic description ofcompounds such as MgO, as the O2− ion is now surrounded by positive Mg2+ ions which produce a stabilizing effect(the lattice energy; see Topic D6).As expected, ion sizes decrease with increasing positive charge, and negative ions are larger. In most ioniccompounds, anions are larger than cations (see Topics D3 and D4).Relativistic effectsSchrödinger’s equation does not take into account effects that are important when particles travel at a speed comparablewith that of light. There are two important aspects: moving charged particles experience magnetic as well as electricfields; and also the special theory of relativity predicts effects such an enhancement of the mass of fast-movingparticles.
These effects were incorporated into the quantum mechanical wave theory by Dirac’s equation (1928).SECTION A—ATOMIC STRUCTURE23One remarkable prediction is the existence of electron spin (see Topic A3) and the occurrence of spin-orbit splittingin atomic spectra. The energies of orbitals are also altered, especially for electrons close to highly charged nuclei, as it isthen that they are travelling fast. Inner shells are most affected but they are not important in chemistry. For very heavyelements even outer shells show an influence of relativity.
This is true for the 6s shell in gold and mercury, and the 6pshell in subsequent elements of period 6. Relativistic effects increase the binding energy of these electrons. They thuscontribute to the irregularities in group trends, and make an appreciable contribution to the high IEs and hencechemical inertness of some heavy elements especially gold and mercury.Section B—Introduction to inorganic substancesB1ELECTRONEGATIVITY AND BOND TYPEKey NotesDefinationsThe bonding triangleBond polarityRelated topicsElectronegativity is the power of an atom to attract electrons to itself in achemical bond.
Different numerical estimates agree on qualitative trends:electronegativity increases from left to right along a period, and generallydecreases down groups in the periodic table. Elements of lowelectronegativity are called electropositive.Electropositive elements form metallic solids at normal temperatures.Electro-negative elements form molecules or polymeric solids withcovalent bonds. Elements of very different electronegativity combine toform solids that can be described by the ionic model.The polarity of a bond arises from the unequal sharing of electronsbetween atoms with different electronegativities.
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