P.A. Cox - Inorganic chemistry (793955), страница 8
Текст из файла (страница 8)
Metallic elements form solidsthat are good conductors of electricity, and have structures with many near neighbors and where bonding is not stronglydirectional. Non-metallic elements form molecules or covalent solids, which are generally poor conductors ofelectricity and where bonding is markedly directional in character. This distinction on the basis of physical properties isfairly clear-cut and is shown in the periodic table in Fig. 1. All elements of the s, d and f blocks are metallic (excepthydrogen), non-metallic ones being confined to hydrogen and to the upper right-hand part of the p block. The most obviousatomic parameter that determines this behavior is electronegativity (see Topic B1, especially Fig. 1).Different types of chemical behavior are associated with the two kinds of element.• Typical metallic elements are good reducing agents (for example, reacting with water to produce dihydrogen) andform hydrated cations in aqueous solution (Na+, Mg2+, etc.).
They have solid halides and oxides, which are welldescribed by the ionic model. The oxides are basic and either react with water to produce hydroxide ions (OH−)or, if insoluble under neutral conditions, dissolve in acidic solutions. Their hydrides are solids with some ionic (H−)character.• Typical non-metallic elements form ionic compounds with electropositive metals. They form anions in water,either monatomic (e.g. Cl−) or oxoanions (e.g. NO3−,). They have molecular hydrides and halides. Theiroxides are either molecular or polymeric covalent in structure, and are acidic, reacting with water (as do halides) toproduce oxoacids (H2CO3, H2SO4, etc.)Fig.
1. Periodic table showing metallic and (heavily shaded) non-metallic elements.B2—CHEMICAL PERIODICITY31It must be recognized that this classification has many limitations, and borderline behavior is common. In addition totheir typical cationic behavior, most metallic elements form some compounds where bonding is predominantly covalent(see, e.g. Topic H10).
Some form anionic species such as MnO4− or even Na− (see Topic G2). Many metals in latergroups are much less electropositive than the typical definition would suggest, and the metal-nonmetal borderline in thep block involves a continuous gradation in chemical behavior rather than a discontinuous boundary (see Topic G6).
Nonmetallic elements close to the metallic borderline (Si, Ge, As, Sb, Se, Te) show less tendency to anionic behavior andare sometimes called metalloids.Horizontal trendsThe major horizontal trends towards the right in any block are a general increase of ionization energy (which is reflectedin an increase in electronegativity), a contraction in size, and an increase in the number of electrons in the valence shell.In main groups, the effect of changing electronegativity is obvious in determining the metal-nonmetal borderline.
Thenumber of valence electrons has a clear influence on the stoichiometry of compounds formed (NaF, MgF2, AlF3, etc.).Main group elements commonly form ions with closed shell configurations: hence cations (Na+, Mg2+, Al3+) inwhich all electrons have been lost from the valence shell, and anions (F−, O2−) in which the valence shell has been filled.This observation suggests some ‘special stability’ of filled shells, but, as in atomic structure (see Topic A5), such aninterpretation is misleading.
The stoichiometry of stable ionic compounds depends on the balance between the energyrequired to form ions and the lattice energy, which provides the bonding (see Topic D6). Such an approach providesa better understanding not only of why closed-shell ions are often found, but also of cases where they are not, as happensfrequently in the d block (see Topics H1 and H3).In covalent compounds some regularities in stoichiometry can also be understood from the increasing number ofvalence electrons.
Thus the simple hydrides of groups 14, 15, 16 and 17 elements have the formulae EH4, EH3, EH2 andEH, respectively, reflecting the octet rule. Filling the valence shell creates progressively more nonbonding electronsand limits the capacity for bonding. Such nonbonding electrons also influence the geometrical structures of themolecules (see Topics C1 and C2).The general increase of electronegativity (or decline in electropositive character) and contraction in size is apparentalso in d-block chemistry. The formation of closed-shell ions (Sc3+, Ti4+, etc.) is a feature of only the early groups. Asionization energies increase more electrons are prevented from involvement in bonding.
Non-bonding d electrons alsoinfluence the structures and stabilities of compounds, but because of the different directional properties of d orbitalscompared with p, these effects are best understood by a different approach, that of ligand field theory (see Section H2).Vertical trendsThe general decrease of ionization energy down a group is reflected in the trend towards metallic elements in the pblock. Another change is the general increase in radius of atoms down a group, which allows a higher coordinationnumber. Sometimes this is reflected in the changing stoichiometry of stable compounds: thus ClF3, BrF5 and IF7 arethe highest fluorides known for elements of group 17.
In other groups the stoichiometry is fixed but the structurechanges: thus the coordination of the metallic element by fluorine is four in BeF2, six in MgF2 and eight in CaF2.Although exceptions occur (see Topics G4 and H5) this is a common trend irrespective of different modes of bonding.One further general feature of vertical trends is important, and reflects the analogous trends in atomic propertiesmentioned in Topic A5. For each block (s, p, d) the first series involved has somewhat distinct chemistry compared withsubsequent ones. Hydrogen (1s) is non-metallic and very different from the other s-block elements.
The 2p-serieselements (B-F) have some peculiarities not shared with the rest of the p block (e.g. a limitation in the number of32SECTION B—INTRODUCTION TO INORGANIC SUBSTANCESvalence-shell electrons in molecules, and the frequent formation of multiple bonds; see Topic F1). In the d block, theelements of the 3d series also show characteristic differences from the 4d and 5d series (e.g. forming many morecompounds with unpaired electrons; see Topics H1–H5).Section B—Introduction to inorganic substancesB3STABILITY AND REACTIVITYKey NotesIntroductionEnthalpy and Hess’ LawEntropy and free energyEquilibrium constantsReaction ratesRelated topicsStability and reactivity can be controlled by thermodynamic factors(depending only on the initial and final states and not on the reactionpathway) or kinetic ones (very dependent on the reaction pathway).Both factors depend on the conditions, and on the possibility ofdifferent routes to decomposition or reaction.Enthalpy change (ΔH) is the heat input to a reaction, a useful measureof the energy change involved.
As ΔH does not depend on thereaction pathway (Hess’ Law) it is often possible to constructthermodynamic cycles that allow values to be estimated for processesthat are not experimentally accessible. Overall ΔH values forreactions can be calculated from tabulated enthalpies of formation.Entropy is a measure of molecular disorder. Entropy changes (ΔS)can be combined with ΔH in the Gibbs free energy change (ΔG),which determines the overall thermodynamic feasibility of a reaction.As with ΔH, ΔG can be estimated from thermodynamic cycles andtabulated values, the latter always referring to standard conditions ofpressure or concentration.The equilibrium constant of reaction is related to the standard Gibbsfree energy change. Equilibrium constants change with temperaturein a way that depends on ΔH for the reaction.Reaction rates depend on the concentrations of reagents, and on arate constant that itself depends on the energy barrier for the reaction.Reaction rates generally increase with rise in temperature.
Catalystsprovide alternative reaction pathways of lower energy.Inorganic reactions andLattice energies (D6)synthesis (B6)Industrialchemistry:Bond strengths (C8)catalysts (J5)IntroductionWe tend to say that substances are ‘stable’ or ‘unstable’, ‘reactive’ or ‘unreactive’ but these terms are relative and maydepend on many factors. Is important to specify the conditions of temperature and pressure, and what other substancesare present or could act as potential routes to decomposition.
Thermodynamic and kinetic factors can also beimportant.34SECTION B—INTRODUCTION TO INORGANIC SUBSTANCESThermodynamics deals with overall energy and entropy changes, and their relation to the direction of reaction andthe position of equilibrium.
Such quantities depend only on the initial and final states, and not at all on the reactionpathway. It is often possible to assess the thermodynamic feasibility of a reaction without any knowledge of themechanism. On the other hand, the rate of a reaction does depend on the pathway; this is the subject of chemicalkinetics, and thermodynamic considerations alone cannot predict how fast a reaction will take place.Many known substances are thermodynamically stable, but others are only kinetically stable. For example,the hydrides B2H6 and SiH4 are thermodynamically unstable with respect to their elements, but in the absence of heat ora catalyst (and of atmospheric oxygen and moisture) the rate of decomposition is extremely slow.
To assess why somesubstances are unknown, it is important to consider different possible routes to decomposition. For example, theunknown CaF(s) is probably thermodynamically stable with respect to the elements themselves, but certainly unstable(thermodynamically and kinetically) with respect to the reactionThermodynamic and kinetic factors depend on temperature and other conditions.
![](https://s.studizba.com/z.php?f=/templates/si/i/noavatar.png&w=100&h=100&t=1"){kind=avatar}
