Atomic size exceptions

We will examine the applicability of this model to metallic solutions by using the refined version which considers that differences in atom sizes will give rise to local inhomogeneities of structure. Because of the negligible error, we will omit all terms having binary products of p, d, and 0, except for p2. 

The prototype hard metals are the compounds of six of the transition metals Ti, Zr, and Hf, as well as V, Nb, and Ta. Their carbides all have the NaCl crystal structure, as do their nitrides except for Ta. The NaCi structure consists of close-packed planes of metal atoms stacked in the fee pattern with the metalloids (C, N) located in the octahedral holes. The borides have the A1B2 structure in which close-packed planes of metal atoms are stacked in the simple hexagonal pattern with all of the trigonal prismatic holes occupied by boron atoms. Thus the structures are based on the highest possible atomic packing densities consistent with the atomic sizes. 

These three structures are the predominant structures of metals, the exceptions being found mainly in such heavy metals as plutonium. Table 6.1 shows the structure in a sequence of the Periodic Groups, and gives a value of the distance of closest approach of two atoms in the metal. This latter may be viewed as representing the atomic size if the atoms are treated as hard spheres. Alternatively it may be treated as an inter-nuclear distance which is determined by the electronic structure of the metal atoms. In the free-electron model of metals, the structure is described as an ordered array of metallic ions immersed in a continuum of free or unbound electrons. A comparison of the ionic radius with the inter-nuclear distance shows that some metals, such as the alkali metals are empty i.e. the ions are small compared with the hard sphere model, while some such as copper are full with the ionic radius being close to the inter-nuclear distance in the metal. A consideration of ionic radii will be made later in the ionic structures of oxides. 

It is found that for metals, low temperature field evaporation almost always produces surfaces with the (1 x 1) structure, or the structure corresponding to the truncation of a solid. A few such surfaces have already been shown in Fig. 2.32. That this should be so can be easily understood. For metals, field penetration depth is usually less than 0.5 A,1 or much smaller than both the atomic size and the step height of the closely packed planes. Low temperature field evaporation proceeds from plane edges of these closely packed planes where the step height is largest and atoms are also much more exposed to the applied field. Atoms in the middle of the planes are well shielded from the applied field by the itinerant electronic charges which will form a smooth surface to lower the surface free energy, and these atoms will not be field evaporated. Therefore the surfaces produced by low temperature field evaporation should have the same structures as the bulk, or the (lxl) structures, and indeed with a few exceptions most of the surfaces produced by low temperature field evaporation exhibit the (1 x 1) structures. 

H 1,0 Li =1.3 Na = 2.2 K - 2.2 Rb = 22 Cs 22 Tlie result of the opposing tendencies or n and Z is that atomic size increases as one progresses down Group IA (1). This is a general property of the periodic chart with but few minor exceptions, which will be discussed later. 

The radii of the metals increase with increasing atomic number and their atomic sizes are the largest in their respective periods. Such features lead to relatively small first ionization energy (/1) for the atoms. Thus the alkali metals are highly reactive and form M+ ions in the vast majority of their compounds. The very high second ionization energy (I2) prohibits formation of the M2+ ions. Even though the electron affinities (T) indicate only mild exothermicity, M- ions can be produced for all the alkali metals (except Li) under carefully controlled conditions. 

N Atomic Size Ionisation energies N, O show oxidation state Melting/Boiling points Electronegativity 3 to +5 Bi show increase from N As M.P./B.P. decrease oxidation state of+3 only As Bi Elements excepts N, Bi... 

Group-18 Elements Noble Gases Increasing Decreasing Exceptions Trends Trends He Atomic size... 

Loss of the Valence Electron. The Is valence electron may be lost to give the hydrogen ion H+ which is merely the proton. Its small size (r 1.5 x 10 13 cm) relative to atomic sizes (r 10 8 cm) and its charge result in a unique ability to distort the electron clouds surrounding other atoms the proton accordingly never exists as such, except in gaseous ion beams in condensed phases it is always closely associated with other atoms or molecules. 

The quantum idea Bohr used in his model of the hydrogen atom was born in 1900. In that year Max Planck (1858-1947), the leading physicist at the University of Berlin, described the universal constant h, called Planck s constant. This constant establishes the scale of quantum phenomena. The extreme smallness of h explains why quantum effects are masked except at atom-size scales. 

As the magnitude of the positive charge of the nucleus increases, its "pull" on all of the electrons increases, and the electrons are drawn closer to the nucleus. This results in a contraction of the atomic radius and therefore a decrease in atomic size. This effect is apparent as we go across the periodic table within a period. Atomic size decreases from left to right in the periodic table (see Figure 3.7). See how many exceptions you can find in Figure 3.7. 

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