Ionic radius copper

The higher ionisation energy and smaller ionic radius of copper contribute to its forming oxides much less polar, less stable, and less basic than those of the alkah metals (13). Because of the relative instabiUty of its oxides, copper joins silver in occurring in nature in the metallic state. 

The viscosities of liquid metals vaty by a factor of about 10 between the empty metals, and the full metals, and typical values are 0.54 x 10 poise for liquid potassium, and 4.1 x 10 poise for liquid copper, at dreir respective melting points. Empty metals are those in which the ionic radius is small compared to the metallic radius, and full metals are those in which the ionic radius is approximately the same as tire metallic radius. The process was described by Andrade as an activated process following an AiThenius expression... 

However, consideration in terms of the ionic radius or the LFSE shows that both factors predict that the maximum stabilities will be associated with nickel(ii) complexes, as opposed to the observed maxima at copper(ii). Can we give a satisfactory explanation for this The data presented above involve Ki values and if we consider the case of 1,2-diaminoethane, these refer to the process in Eq. (8.13). 

The data for the 1,2-diaminoethane complexes now parallels the trends in ionic radius and LFSE rather closely, except for the iron case, to which we return shortly. What is happening Copper(ii) ions possess a configuration, and you will recall that we expect such a configuration to exhibit a Jahn-Teller distortion - the six metal-ligand bonds in octahedral copper(ii) complexes are not all of equal strength. The typical pattern of Jahn-Teller distortions observed in copper(ii) complexes involves the formation of four short and two long metal-ligand bonds. 

Derived from the German word meaning devil s copper, nickel is found predominantly in two isotopic forms, Ni (68% natural abundance) and Ni (26%). Ni exists in four oxidation states, 0, I, II, III, and IV. Ni(II), which is the most common oxidation state, has an ionic radius of —65 pm in the four-coordinate state and —80 pm in the octahedral low-spin state. The Ni(II) aqua cation exhibits a pAa of 9.9. It forms tight complexes with histidine (log Af = 15.9) and, among the first-row transition metals, is second only to Cu(II) in its ability to complex with acidic amino acids (log K( = 6-7 (7). Although Ni(II) is most common, the paramagnetic Ni(I) and Ni(III) states are also attainable. Ni(I), a (P metal, can exist only in the S = state, whereas Ni(lll), a cT ion, can be either S = or S =. ... 

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. 

Identical Ni-S and Cu-S bond distances are observed in these two structures, and the expected reduction of the ionic radius of copper is not reflected in a shorter Cu-S bond. Two possible reasons for this result are (1) a concomitant increase of the ionic radii of the sulfur donors offsets the effect of a shorter Cu(III) radius, and (2) there exists significant covalency in the Cu—S bonds. 

On the basis of ionic radius, cobalt ion should be placed between copper and cadmium ions, and thus it has a higher affinity than expected. Perhaps it substitutes the iron(II) of crystal lattice since the two ionic radii are about the same (0.84 nm). 

In spite of its larger size, with an ionic radius of 0.80A (161), In3+ appears to bind to transferrin with an affinity close to that of Fe3+ (145). In3+ displaces Cu2+ from copper-saturated ovotransferrin, and the In3+ even remains bound in the presence of an added twofold excess of Fe3+. Indium-transferrin also migrates indistinguishably from iron-transferrin (145) and gives the same closed conformation, as judged by small-angle X-ray scattering (105). 

Zinc has a highly concentrated charge in comparison to its relatively small ionic radius (0.65 A) and binds modestly to anions such as carboxylates and phosphates. Its second characteristic is its high affinity for electrons, making it a strong Lewis acid, similar to copper and nickel. However, unlike the other two transition metal ions, it does not show variable valence, which might lead to it being preferred quite simply because it does not introduce the risk of free radical reactions. 

Thallium is a rare element which occurs in the Earth s crust at an estimated abundance of 0.1 to 0.5 ligg (see Part I, Chapter 1). The specific ionic properties of thallium (e.g., ionic radius Tl 0.147 nm) are similar to those of potassium and rubidium (ionic radius K 0.133 nm, Rb" 0.147 nm) thus, thallium occurs ubiquitously as a trace element within the environment, mainly in association with K and Rb. Besides its occurrence in widespread potassium compounds, thallium is a trace component in iron, zinc, copper, and lead minerals (Nriagu 1998). 

---