P.A. Cox - Inorganic chemistry (793955), страница 7
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There is no sharpdividing line between polar covalent and ionic substances.Trends in atomic propertiesIntroduction to solids (D1)(A5)Electron pair bonds (C1)DefinitionsElectronegativity may be defined as the power of an atom to attract electrons to itself in a chemical bond.It is the most important chemical parameter in determining the type of chemical bonds formed between atoms. It ishard to quantify in a satisfactory way, especially as electronegativity is not strictly a property of atoms on their own, butdepends to some extent on their state of chemical combination. Nevertheless several scales have been devised.• Pauling electronegativity is based on bond energies (see Topic C8), using the empirical observation that bondsbetween atoms with a large electronegativity difference tend to be stronger than those where the difference is small.This scale was historically the first to be devised and although it lacks a firm theoretical justification is still widelyused.• Mulliken electronegativity is the average of the first ionization energy and the electron affinity of an atom (seeTopic A5), reflecting the importance of two possibilities in bond formation, losing an electron or gaining one.
Thescale has the advantage that electronegativity values can be estimated not only for the ground states of atoms, but forother electron configurations and even for polyatomic fragments.• Allred-Rochow electronegativity is proportional to Zeff/r2, where Zeff is the effective nuclear charge of valenceorbitals (see Topic A3), and r the covalent radius of the atom. The value is proportional to the effective electrostaticattraction on valence electrons by the nucleus, screened by inner shell electrons.26SECTION B—INTRODUCTION TO INORGANIC SUBSTANCESFig.
1. Pauling electronegativity values for the elements H–K. Elements in the shaded region are metallic (see Topic B2).Each scale produces different numbers and they should not be mixed. The broad general trends do, however, agree:electronegativity increases towards the right and decreases towards the bottom in the periodic table. Itthus follows the same trend as atomic ionization energies (see Topic A5).
Elements in early groups have low values andare called electropositive. Figure 1 shows the Pauling electronegativities of elements up to potassium. Elements ofgroup 18 in early periods do not form any stable compounds, and so the most electronegative element is fluorine.The bonding triangleThe bonding triangle (see Fig. 2) is a useful way of showing how the electronegtivities of two elements A and B (whichmay be the same) determine the type of bond formed between them.
The horizontal and vertical scales show thePauling electronegativities of the two elements. (Other scales would do equally well at this qualitative level.) Pure elements(A=B) appear on the diagonal, and various compounds are shown within the triangle. Three basic regions aredistinguished.• When A and B are both electropositive they form a metallic solid, characterized by high electrical conductivity anda structure where each atom is surrounded by many others (often 12; see Topic D2).
Metallic bonding involves thedelocalization of electrons throughout the solid. The electrons are shared between atoms as in covalent bonding(see below), but in a less specific way and without the directional character of covalent bonds.• When A and B are both electronegative they form covalent compounds. These may consist of individualmolecules (O2, H2O, etc.) or of giant covalent lattices (polymeric solids) with a continuous network ofbonds. Although the dividing line between these types is not sharp, very highly electronegative atoms (F, O, Cl,etc.) have more tendency to molecular behavior in both their elements and their compounds. Covalent solids do notconduct electricity well. The most important feature of this bonding, whether in molecules and solids, is its highlydirectional and specific nature.
Thus the neighbors to any atom are limited in number (e.g. four in the case ofelemental silicon, three for phosphorus, two for sulfur, one for chlorine), and are generally found in specificB1—ELECTRONEGATIVITY AND BOND TYPE27Fig. 2. The bonding triangle, showing a selection of elements and compounds plotted against the Pauling electronegativities.geometrical arrangements. The simplest view of covalent bonding involves the sharing of electrons in specific,localized bonds between atoms (see Topic C1).• When one atom is very electropositive and the other very electronegative, a solid compound is formed that is oftenregarded as ionic.
In this picture there is a complete transfer of one or more electrons, giving cations of theelectropositive element and anions of the electronegative one, which are then held together by electrostaticattraction (see Topics D3, D4 and D6). Solids are formed rather than molecules because the force is not directional,and greatest stability is achieved by packing several anions around each cation and vice versa.Bond polarityA covalent bond between two atoms of the same element is described as homopolar, one between different elementsas heteropolar; the general term bond polarity describes the unequal sharing of electrons between two atoms, andis a feature of heteropolar bonds when the two elements concerned have a different electronegativity. The moreelectronegative atom draws electrons and thus acquires a partial negative charge, with the other atom becomingcorrespondingly positive.
One manifestation of such polarity is the formation of an electric dipole moment, themagnitude of which is equal to the product of the charges and their average separation. The dipole moments decrease ina series of molecules such as HF> HCl>HBr>HI as might be expected from the falling difference in electronegativities.Dipole moments are, however, not always easy to interpret, as they can be influenced by other factors, such as therelative orientation of bonds in polyatomic molecules and the distribution of nonbonding electrons.
Dipole momentsare an important source of intermolecular forces (see Topic C10).28SECTION B—INTRODUCTION TO INORGANIC SUBSTANCESPolar covalent bonds can be regarded as having some degree of ionic character, and the distinction between ‘ionic’and ‘covalent’ bond types is sometimes hard to make.
Some compounds have clear examples of both types of bondingsimultaneously. Thus CaCO3 has well-defined carbonate ionswith C and O covalently bonded together; thecomplex ion also interacts ionically with Ca2+. Such complex ions need not be discrete entities but can formpolymeric covalent networks with a net charge, with ionic bonds to cations (e.g. silicates; see Topics D6 and F4).
Evenwhen only two elements are present, however, bonding may be hard to describe in simple terms.When a compound is molecular under normal conditions it is usual to regard it as covalent (although ‘ionicmolecules’ such as NaCl(g) can at be made by vaporizing the solid compounds at high temperatures). When twoelements of different electronegativity form a solid compound alternative descriptions may be possible. Consider thecompounds BeO and BN. Both form structures in which every atom is surrounded tetrahedrally by four of the otherkind (BN also has an alternative structure similar to that of graphite).
For BeO this is a plausible structure on ionicgrounds, given that the Be2+ ion must be much smaller than O2− (see Topic D4). On the other hand, many of thestructures and properties of beryllium compounds are suggestive of some degree of covalent bonding (see Topic G3).Thus one can think of BeO as predominantly ionic, but with the oxide ion polarized by the very small Be2+ ion so thatelectron transfer and ionic character are not complete. For BN the electronegativity difference between elements ismuch less, and it would be more natural to think of polar covalent bonding.
The tetrahedral structure of BN can beunderstood from its similarity to diamond, where each carbon atom is covalently bonded to four others. The differencebetween two descriptions ‘polarized ionic’ and ‘polar covalent’ is not absolute but only one of degree. Which startingpoint is better cannot be laid down by rigid rules but is partly a matter of convenience.One should beware of using oversimplified criteria of bond type based on physical properties. It is sometimes statedthat ‘typical’ ionic compounds have high melting points and dissolve well in polar solvents such as water, whereascovalent compounds have low melting points and dissolve well in nonpolar solvents.
This can be very misleading.Diamond, a purely covalent substance, has one of highest melting points known and is insoluble in any solvent. Somecompounds well described by the ionic model have fairly low melting points; others are very insoluble in water ongrounds that can be explained perfectly satisfactorily in terms of ions (see Topic E4).Section B—Introduction to inorganic substancesB2CHEMICAL PERIODICITYKey NotesIntroductionMetallic and nonmetallic elementsHorizontal trendsVertical trendsRelated topicsMajor chemical trends, horizontally and vertically in the periodictable, can be understood in terms of changing atomic properties. Thisprocedure has its limitations and many details of the chemistry ofindividual elements cannot be predicted by simple interpolation fromtheir neighbors.Metallic elements are electropositive, form electrically conductingsolids and have cationic chemistry.
Non-metallic elements, found inthe upper right-hand portion of the periodic table, havepredominantly covalent and anionic chemistry. The chemical trend iscontinuous and elements on the borderline show intermediatecharacteristics.Moving to the right in the periodic table, bonding character changesas electro-negativity increases. The increasing number of electrons inthe valence shell also gives rise to changes in the stoichiometry andstructure of compounds. Similar trends operate in the d block.The increased size of atoms in lower periods is manifested instructural trends.
For each block, changes in chemistry between thefirst and second rows concerned are often more marked than thosebetween lower periods.The periodic table (A4)IntroductiontoTrends in atomic propertiesnontransition metals (G1)(A5)Introduction to transitionIntroduction to nonmetalsmetals (H1)(F1)IntroductionThe periodic table was devised by Mendeleev in response to observed regularities in the chemistry of the elementsbefore there was any understanding of their electronic basis (see Topic A4). His procedure was vindicated by his abilityto predict the properties and simple chemistry of the then unknown elements gallium and germanium by simpleinterpolation between known elements in neighboring positions. Chemical periodicity was thus seen to be a powerfultool in the interpretation and even prediction of the chemical properties of elements.Since Mendeleev the range of chemical compounds known has expanded enormously and it has become apparent thatsuch simple interpolation procedures have many limitations.
In a few groups (especially the s block) the chemistry isfairly similar, and most of the observed trends in the group can be interpreted straightforwardly from changes of atomic30SECTION B—INTRODUCTION TO INORGANIC SUBSTANCESproperties such as radius. In the p and d blocks, however, this is not so easy. Complications arise partly from the fact thatatomic trends are themselves less regular (because of the way in which the periodic table is filled), and partly from thegreater complexities in chemical bonding, which respond in a more subtle way to changes in orbital size and energy.The periodic table remains the most important framework for understanding the comparative chemistry of elements,and many major trends can be understood from the atomic trends described in Topic A5. Most elements havepeculiarities, however, which although they can be rationalized in terms of periodic trends, would probably not havebeen predicted if they were not known.Metallic and non-metallic elementsThe most important classification of elements is that of metallic versus non-metallic.
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