Elements of the p Block

edited by Charlie Harding, Rob Janes and David Johnson

Based on 'The Chemistry of Hydrogen, the Halogens and the Noble Gases' by David Johnson (1994) and 'The Chemistry of Groups III-VI' by David Johnson and Lesley Smart (1994)

INTRODUCTION 1

Chemistry is an immensely diverse subject. Nowhere is this more evident than in the properties of the p-Block elements, those in which the outer electronic configuration is s2px. In this Book, we will study the typical elements, focusing on the p-Block elements (Figure 1.1).

Much of the chemistry of the metallic s-Block elements is consistent with an ionic model. With the p-Block elements, this is no longer the case: the proportion of obviously covalent substances is much larger. Before we begin our study of these elements, it is necessary to introduce some essential principles, mainly concerning bonding and thermochemistry, which apply particularly to covalent substances. Following this, we consider the chemistry of the Groups of p-Block elements, beginning with the halogens and noble gases, as these Groups most clearly display covalent properties. These are the two Groups that contain no metallic elements.

1.1 Group numbers and the Periodic Table

Because this Book concentrates upon the typical elements, we shall make particular use of the mini-Periodic Table of Figure 1.1 that contains the typical elements alone. In Figure 1.1, the Groups are numbered in two different ways. In the Mendeléev numbering scheme, which is given in Roman numerals, the Group numbers increase smoothly across each row from I to VIII. These numbers are useful because, for each element, they are almost always equal to the number of outer electrons in the atom, and to the highest oxidation state (see Section 2.1) that is reached in the compounds that the element forms. Use is made of these relationships in this Book. Most modern textbooks and chemistry research journals, however, use a numbering system recommended by IUPAC (International Union of Pure and Applied Chemistry). In this system, the Groups are numbered in Arabic numerals as Groups 1–18 (Figure 1.2). The p-Block elements of Groups III–VIII are, therefore, numbered as Groups 13–18. Because both numbering schemes are important, we include both alternatives in the headings of the Sections and the figures describing the Groups in this Book. Thus the Group of elements headed by boron is designated Group III/13. Notice that for the typical elements of Figure 1.1, the second digit of a Group number in the IUPAC scheme is equal to the Mendeléev number. This provides an easy way of shifting from one system to another.

OXIDATION AND REDUCTION 2

Oxidation may be defined as a loss of electrons, and reduction as a gain of electrons. In a reaction such as

Mg(s) + Cl2(g) = MgCl2(s) (2.1)

the ionic model implies that magnesium dichloride is composed of the ions Mg2+ and Cl-. Consequently, Reaction 2.1 is a redox reaction: magnesium has been oxidized because a magnesium atom loses two electrons; chlorine has been reduced because a chlorine atom gains an electron. However, this argument assumes that MgC12 conforms to an ionic model. While this may be justified for compounds such as MgC12, a difficulty arises when we turn to covalent compounds. Consider the reaction between sulfur and chlorine:

S(s) + Cl2(g) = SCl2(l) (2.2)

Reactions 2.1 and 2.2 appear similar, but at room temperature, SCl2 is a red liquid containing discrete SCl2 molecules, whose atoms are held together by covalent bonds (Structure 2.1); there cannot be complete transfer of two electrons per sulfur atom from sulfur to chlorine. Consequently, our definitions of oxidation and reduction prevent the classification of both Reactions 2.1 and 2.2 as redox reactions. To obtain a broader definition of oxidation and reduction, we must abandon the implication that complete electron transfer takes place.

2.1 Oxidation numbers or oxidation states

The solution that has been devised for this problem is to imagine that the shared electron pairs in covalent bonds between different elements are completely transferred to one of the two bound atoms. The transfer is assumed to occur to the more electronegative of the two.

* Consider Structure 2.1. Which is the more electronegative type of atom?

* Chlorine is more electronegative than sulfur; it lies to the right of sulfur in Period 3 (Figure 1.1).

If the shared electron pair in each bond in Structure 2.1 were transferred to chlorine, each chlorine would have eight outer electrons, one more than in the free atom. It would therefore become Cl-, and carry a charge of -1.

* What number of outer electrons would the sulfur be left with, and what charge would it carry?

* Sulfur would be left with four outer electrons, two less than in the free atom, and would have a charge of +2.

These imaginary charges are called oxidation numbers or oxidation states. The oxidation number of an atom in its elemental form, e.g. graphite, aluminium metal and chlorine gas, is taken to be zero.

Oxidation is now defined as an increase in oxidation number, and reduction as a decrease in oxidation number.

* Why is Reaction 2.2 a redox reaction? What has been oxidized and reduced in it?

* Sulfur has been oxidized because its oxidation number has increased from zero to +2; chlorine has been reduced because its oxidation number has decreased from zero to -1.

To deduce the oxidation numbers of the elements in a compound or ion, it is unnecessary to write down a Lewis structure, as described above. Instead a...