Order in the elements

Most of the chemistry in this book revolves around elements and isotopes (see Box 1.1). It is therefore helpful to understand how the atomic number (Z) of an element, and its electron energy levels allow an element to be classified. The electron is the component of the atom used in bonding (Section 2.3). During bonding, electrons are either donated from one atom to another, or shared; in either case the electron is prised away from the atom. One way of ordering the elements is therefore to determine how easy it is to remove an electron from its atom. Chemists call the energy input required to detach the loosest electron from atoms, the ionization energy. As explained in Box 1.1, the number of positively charged components (protons, Z) in an atom is balanced by the same number of negatively charged electrons that form a 'cloud' around the nucleus. Although electrons do not follow precise orbits around the nucleus, they do occupy specific spatial domains called orbitals. We need only think in terms of layers of these orbitals. Those electrons in orbitals nearest the nucleus are tightly held by electrostatic attraction-forming core electrons that never take part in chemical reactions. Those further away from the nucleus are less tightly held and may be used in 'transactions' with other atoms. These loosely held electrons are known as valence electrons. Electrons normally occupy spaces available in the lowest energy orbitals such that energy dictates the electron distribution around the nucleus. The valence electrons reside in the highest occupied energy levels and are thus the easiest to remove. For example, the element sodium (Na) has a Z number of 11. This means that sodium has 11 electrons, 10 of which are core electrons, and one valence electron. It is this single valence electron that dictates the way sodium behaves in chemical reactions.

Plotting the expected first ionization energy—i.e. that required to detach the loosest valence electron from the atom—against atomic number (Fig. 2.1a), shows that as atomic number increases the energy required to detach valence electrons decreases from Z = 1 (H) to Z = 20 (Ca). In this diagram the increasing nuclear charge between hydrogen (H) and calcium (Ca) has been disregarded. The clear downward steps in energy mark large energy gaps where electrons occupy progressively higher energy orbitals further away from the nucleus. The steps in Fig. 2.1a predict a marked difference in atomic structure between helium (He) and lithium (Li), between neon (Ne) and sodium (Na) and between argon (Ar) and potassium (K). Although much simplified, this periodic repetition of the elements has long been used as the basis to tabulate the ordering of elements on a grid known as the Periodic Table (Fig. 2.2), first published in its modern form by Mendeleev in 1869.

If the ionization energy is corrected to account for nuclear charge (Fig. 2.1b) —because increasing nuclear charge makes electron removal more difficult—the energy pattern in each period becomes more like a ramp. Each 'period' begins with an element of conspicuously low ionization energy, the so-called alkali metals (Li, Na and K). Each of these elements readily lose their single valence electron to form singly charged or monovalent ions (Li+, Na+ and K+). The periods of elements are depicted as 'rows' in the Periodic Table (Fig. 2.2), and when these rows are stacked on top of one another a series of 'columns' result (Fig. 2.2). Column Ia depicts the alkali metals. Moving up the energy ramps in Fig. 2.1b, the alkali metals are followed by the elements beryllium (Be), magnesium (Mg) and calcium (Ca), each with two, relatively easily removed valence

Fig. 2.1 (a) Expected first ionization energy plotted against atomic number (Z), up to Z = 20. This plot disregards the effects of increasing nuclear charge with Z. (b) Variation of measured first ionization energy with atomic number Z, up to Z = 20. The ramped profile (arrows) of increasing ionization energy following the abrupt drops reflects the increasing nuclear charge. After Gill (1996), with kind permission of Kluwer Academic Publishers.

Fig. 2.1 (a) Expected first ionization energy plotted against atomic number (Z), up to Z = 20. This plot disregards the effects of increasing nuclear charge with Z. (b) Variation of measured first ionization energy with atomic number Z, up to Z = 20. The ramped profile (arrows) of increasing ionization energy following the abrupt drops reflects the increasing nuclear charge. After Gill (1996), with kind permission of Kluwer Academic Publishers.

electrons. These elements form doubly charged or divalent ions (Be2+, Mg2+ and Ca2+) and are known as alkali earth metals (column IIa in the Periodic Table). Continued progression up each energy ramp in Fig. 2.1b results in predictable patterns. For example, Mg is followed by aluminium (Al) which has three valence electrons, and then silicon with four valence electrons. Progressively more energy

a subgroups--b subgroups

Ia

Ila

Ilia

IVa

Va

Via

Vila

VIII

Ib

llb

IIIb

IVb

Vb

VIb

VIIb

O

Fig. 2.2 Periodic Table of the elements and their Z numbers. Note that the periodic pattern is complicated by the transition metals between columns II and III. *La and the lanthanides are known as the rare earth elements (REE). The table has been constructed using conventional terminology and further details can be found in basic chemistry textbooks. Gill (1996) gives an accessible summary with a strong applied earth science stance. Elements in bold are those most abundant in environmental materials (see Fig. 2.3). After Gill (1996), with kind permission of Kluwer Academic Publishers.

Fig. 2.2 Periodic Table of the elements and their Z numbers. Note that the periodic pattern is complicated by the transition metals between columns II and III. *La and the lanthanides are known as the rare earth elements (REE). The table has been constructed using conventional terminology and further details can be found in basic chemistry textbooks. Gill (1996) gives an accessible summary with a strong applied earth science stance. Elements in bold are those most abundant in environmental materials (see Fig. 2.3). After Gill (1996), with kind permission of Kluwer Academic Publishers.

is required to remove these electrons due to the increasing nuclear attraction. This means that aluminium will form trivalent cations whereas silicon typically will not: instead it shares its electrons in covalent bonds (Section 2.3), except in one special case (Section 2.3.2). At the top of each energy ramp are the elements He, Ne and Ar that cling tenaciously to all of their electrons. These elements have no valence electrons and therefore no significant chemical reactivity. These chemically inert elements are often called the inert or noble gases and form column O on the far right of the Periodic Table (Fig. 2.2).

Although the periodic pattern becomes more complicated above Z values of 20, the overall ordering persists. Complications arise in the so-called transition elements that occupy a position between columns II and III of the Periodic Table (Fig. 2.2). These elements have between one and three valence electrons. Importantly, however, the electrons in the orbital below the valence electrons have almost the same energy as the valence electrons themselves. In some compounds, usually depending on oxidation state (see Box 4.3), these lower orbital electrons act as additional valence electrons. For example, the element iron (Fe) exists in compounds in a reduced (Fe2+ or ferrous iron) and oxidized (Fe3+ or ferric iron) state. In general, the transition metals are less regular in their atomic properties when compared to the main groups, which also makes their behaviour more complicated to predict in nature.

It is clear from the discussion above, and by looking at the Periodic Table (Fig. 2.2) that some elements are classed as metals, some as semi-metals and some as non-metals. In each row of the Periodic Table the degree of metallic character decreases progressively from left to right, i.e. up the energy ramps of Fig. 2.1b. In essence this is because those elements with low ionization energy hold electrons loosely. In an applied electrical voltage these excited electrons will flow, conducting the electricity, whereas in non-metals there is a gap in the electron configuration that will not allow passage of excited electrons. In the case of semi-metals the gap in electron configuration is small enough that excited electrons can jump through, but only when activated by an external energy source. In effect the semi-metal flips between being an insulator (when not stimulated by external energy) and a conductor (when stimulated by external energy). Semi-metals such as silicon are also known as semi-conductors, and are used in various industrial applications to speed up electrical processes, most famously as the key component of the 'silicon chip' in computer microprocessors.

There have been many attempts to further classify the elements geologically and environmentally. In Fig. 2.3 we show the most abundant elements in four of the main environmental materials of the Earth. A glance at this figure shows that

Fig. 2.3 Distribution of elements in the four main environmental materials, lithosphere, hydrosphere, atmosphere and biosphere. The elements are shown in their actual form as compounds, ions or molecules as appropriate. The main components of each material are shown in boxes, other major constituents are shown outside the boxes.

Fig. 2.3 Distribution of elements in the four main environmental materials, lithosphere, hydrosphere, atmosphere and biosphere. The elements are shown in their actual form as compounds, ions or molecules as appropriate. The main components of each material are shown in boxes, other major constituents are shown outside the boxes.

oxygen (O), and to a lesser extent hydrogen (H), are superabundant in most Earth surface materials such as air, water, organic matter and silicate minerals. In the lithosphere, silicon (Si) and aluminium (Al) are next most abundant forming the silicate minerals feldspar and quartz (see Chapter 4). In the hydrosphere it is the dissolved ions in seawater (see Chapter 6) that dominate the chemistry, particularly chloride (Cl-) and sodium (Na+), while the main atmospheric gases are nitrogen (N2), oxygen (O2), argon (Ar) and carbon dioxide (CO2), along with water vapour (see Chapter 3). The organic matter of the biosphere is made principally of carbon and hydrogen bonded in various combinations (Section 2.7), along with lesser amounts of oxygen and the nutrient elements nitrogen (N) and phosphorus (P). Based on the information in this diagram it might be tempting to conclude that we need only understand the behaviour of these elements in nature to understand environmental chemistry. In fact the reverse is true. Paradoxically, it is often the elements present in trace amounts in the solids and fluids of the environment that tell us most about chemical processes.

2.3 Bonding

Many elements do not normally exist as atoms, but are bonded together to form molecules. The major components of air, nitrogen and oxygen for example, are present in the lower atmosphere as the molecules N2 and O2. By contrast, argon is rather unusual because as an inert element (or noble gas — Section 2.2) it is found uncombined as single argon atoms. Inert elements are exceptions and most substances in the environment are in the form of molecules.

2.3.1 Covalent bonds

Molecular bonds are formed from the electrostatic interactions between electrons and the nuclei of atoms. There are many different electronic arrangements that lead to bond formation and the type of bond formed influences the properties of the compound that results. It is the outermost electrons of an atom that are involved in bond formation. The archetypical chemical bond is the covalent bond and we can probably best imagine this as formed from outer electrons shared between two atoms. Take the example of two fluorine atoms that form the fluorine molecule:

In this representation of bonding, the electrons are shown by dots. In reality the bonding electrons are smeared out over the entire molecules, but their most probable position is between the nuclei. The bond is shown as the two electrons between the atoms. The bond is created from the two electrons shared between the atoms. In simple terms it can be argued that this arrangement of electrons achieves a structure similar to that of argon, i.e.:

Thus bond formation can be envisaged as a result of attaining noble-gas-type structures that have particularly stable configurations of electrons. Symbolically this covalent bond is written F-F. We can think of the bonding electrons, which tend to sit between the two nuclei, as shielding the repulsive forces of the protons in the nucleus.

Oxygen and nitrogen are a little different:

For oxygen the argon-like structure requires two electrons from each atom and the double bond formed is symbolized O=O. For nitrogen we have:

symbolizedN = N (a triplebond)

Gases in the atmosphere, water and organic compounds (Section 2.7) are typically formed with these kinds of covalent bonds.

2.3.2 Ionic bonding, ions and ionic solids

Unlike oxygen (O2) and nitrogen (N2), where individual atoms bond by sharing electrons, many crystalline inorganic materials bond by donating and accepting electrons. In fact, it can be argued that these structures have no bond at all, because the atoms entirely lose or gain electrons. This behaviour is usually referred to as ionic bonding. The classic example of an ionic solid is sodium chloride (NaCl):

As in equations 2.1-2.3 the dots represent electrons.

The theory behind this behaviour is that elements with electronic structures close to those of inert (noble) gases lose or gain electrons to achieve a stable (inert) structure. In equation 2.4, sodium (Na; Z = 11) loses one electron to attain the electronic structure of neon (Ne; Z = 10), while chlorine (Cl; Z = 17) gains one electron to attain the electronic structure of argon (Ar; Z = 18). The compound NaCl is formed by the transfer of one electron from sodium to chlorine and the solid is bonded by the electrostatic attraction of the donated/received electron. The compound is electrically neutral.

Crystalline solids, for example NaCl, are easily dissolved in polar solvents such as water (see Box 4.1), which break down the ionic crystal into a solution of separate charged ions:

(Note: Most equations in this book will not show the electrons (•) on individual ions and atoms, only their charge (+/-).)

Positively charged atoms like Na+ are known as cations, while negatively charged ions like Cl- are called anions. Thus, metals whose atoms have one, two or three electrons more than an inert gas structure form monovalent (e.g. potassium, K+), divalent (e.g. calcium, Ca2+) or trivalent (e.g. aluminium, Al3+) cations.

Similarly, non-metals whose atoms have one, two or three electrons less than an inert gas structure form monovalent (e.g. bromine, Br-), divalent (e.g. sulphur, S2-) and trivalent (e.g. nitrogen, N3-) anions. In general, the addition or loss of more than three electrons is energetically unfavourable, and atoms requiring such transfers generally bond covalently (Section 2.3.1).

The silicon ion Si4+ is an interesting exception. The high charge and small ionic radius make this cation polarizing or electronegative (see Box 4.2), such that its bonds with the oxygen anion O2- in silicate minerals (see Section 4.2) are distorted. This produces an appreciable degree of covalency in the Si-O bond.

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