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Chapter: Essentials of Inorganic Chemistry : Introduction

Nowadays, chemical bonding is mainly explained by the application of wave mechanics and described either by the valence bond (VB) or molecular orbital (MO) theory.

Basic inorganic principles

It is important to understand the basic inorganic principles in order to evaluate the full potential of inorganic compounds in clinical applications. In the following sections, aspects such as atomic structures, chemical bonds and the set-up of the periodic table will be discussed.



1. Introduction

Historically, the formation of ionic species was seen as a result of the transfer of electrons between atoms. The result is the formation of an anion (negatively charged partner) and a cation (positively charged partner), which form a strong bond based on electrostatic attraction. The bonding situation in covalently bonded molecules was described as a sharing of valence electrons. The valence electrons involved belong to neither of the atoms involved completely. Nowadays, chemical bonding is mainly explained by the application of wave mechanics and described either by the valence bond (VB) or molecular orbital (MO) theory.

A chemical bond is defined as an attraction between atoms, which leads to the formation of chemical substances containing two or more atoms. The bond is a result of the electrostatic attraction between opposite charges, such as electrons or nuclei or dipole attraction.

The forces behind this attraction are electrostatic forces of varying degrees, ranging from weak dipole interactions to strong attraction between opposite charges. Strong bonds are ionic bonds and covalent bonds, whilst weak chemical bonding can be observed where dipole interaction and hydrogen bonding is seen.

Lewis structures are used to simply describe how valence electrons are arranged in molecules and how they are involved in chemical bonds. Basically, dots are used to visualise the number of valence electrons, whereas the elemental symbol represents the nuclei. As a basic rule, electrons should be ruled. Paired electrons are sometimes also represented by a line, which can be interpreted as a single covalent bond. An element with a single electron represents a radical. Electron pairs not contributing to any bonds are called lone pairs. The Lewis structure of water can be found in Figure 1.11. Oxygen (O) has six valence electrons and hydrogen (H) has one valence electron each, adding up to a total of eight electrons around the O centre. Two electrons each are used for the formation of a covalent single bond between the O and the H nucleus, whilst O keeps two electron pairs as lone pairs.

Double and triple bonds can also be symbolised using Lewis structures. Using the Lewis structure nitrogen (N) as an example (see Figure 1.12), it can be seen that three pairs of electrons form the N—N bond, whilst each N centre keeps one electron pair as lone pair. Again, eight electrons in total directly surround each N centre – two from the lone pair and six from the triple bond.

Each atom (at least of atomic numbers <20) tends to form molecules in this way so that it has eight electrons in its valence shell. This gives them the same electronic configuration as noble gases. This rule of thumb is often used for main group elements and is called the Octet Rule. Lewis structures are often used to visualise the valence electrons. Note that valence electrons used in bonds are counted twice, once for each bond partner.

The VSEPR (valence shell electron pair repulsion) rules are a set of rules used to predict the shape of a molecule. The basic principle is that valence electrons around the centre atom repel each other and therefore will form an arrangement in which they are situated furthest from each other. Also, lone pairs are included in this electrostatic repulsion. Double and triple bonds are seen the same as single bonds. The resulting geometry depends on the number of binding and nonbinding electron pairs, which can be summarised as the steric number. 

Figure 1.13 gives a selection of geometries found in typical main group molecules. Note that E denotes lone pairs, whereas B represents a covalent bond atom or group. Lone pairs occupy slightly more space than binding electron pairs, which explains the smaller angle in water of 104.9 rather than the standard tetrahedral angle of 108.9.


2. Covalent bond

The simplest way to describe a covalent bond is based on the picture of sharing electron pairs between two atoms. Each atom would contribute electrons to this bond. In a dative covalent bond, one bond partner would donate all the electrons needed to form the bond. This type of binding is very important for biological actions of various elements.

A covalent bond is defined as a chemical bond that is based on the sharing of electrons. Often, this leads to full outer shells for the binding partner to obtain the noble gas configuration.

Within homonuclear species (chemical bond occurs between two atoms of the same element), the binding electron pair is evenly distributed between the two partners. In a heteronuclear species (chemical bond occurs between two atoms of different elements), the electrons are more attracted/polarised to one partner than the other, depending on the so-called electronegativity (EN).

Electronegativity describes the tendency of an atom to attract electrons or electron density towards itself.

EN depends on the atomic number and the distance of the valence electrons from the nucleus. There are different scales for calculating the relative values, but the Pauling scale is the most commonly used one. Visualising EN trends on the periodic table shows an increase of EN within the row and a decrease within the group (Figure 1.14).

Using a more advanced approach, chemical bonding is nowadays mainly explained by the application of wave mechanics and described either by the VB or MO theory. The VB theory describes the formation of a covalent bond as the overlapping of two half-filled valence AOs from each binding partner, which contains one electron each. The simplest example is probably the molecule H2. Hydrogen has only one valence electron (1s1), and therefore two hydrogen 1s orbitals filled with one electron (1s1), can overlap and form a chemical bond (Figure 1.15).

In a HF molecule, the 1s orbital of H and the 2pz orbital of F, each filled with one unpaired electron, overlap and form a covalent bond (Figure 1.16).

MO theory approaches chemical binding from a more advanced point of view, where MOs are formed covering the whole molecule. The principal idea is that the AOs (as discussed in Section of both binding partners are combined and form binding and nonbinding MOs. Those MO are filled with electrons contributed from both binding partners (Figure 1.17).


3. Ionic bond

Ionic bonds are strong bonds based on the transfer of electrons between the atoms and the resulting electro-static attraction between the negatively and positively charged bond partners (Figure 1.18).

The term ion is used for atoms or molecules in which the total number of electrons is different from the number of protons and therefore carries a positive or negative charge. An anion is an atom or molecule that has a negative charge (has more electrons than protons). In contrast to this, a cation is defined as an atom or molecule with a positive charge, that is, it has more protons than electrons.

Ionic bonds are strong but short-range bonds and they have no defined direction in space, as they are ‘only’ the result of electrostatic attraction. Therefore, these attractions are not only limited to one directional partner each but also with each ion around them. As a result, a whole network of ions will be formed with anions and cations occupying specified spaces. The result is called a salt, which typically have high melting points (Figure 1.19).

 Figure 1.19 Example of a salt structure (NaCl) Reproduced with permission from Gerard J. Tortora and Bryan H. Derrickson, Principles of Anatomy and Physiology: Organization, Support and Movement, and Control Systems of the Human Body, 13th Edition International Student Version. Copyright © 2011, John Wiley & Sons, Ltd

4. Metallic bond

A metallic bond is most commonly described as a type of chemical bond where the metal atom donates its valence electrons to a ‘pool’ of electrons that surrounds the network of metal atoms. Electrons are not anymore identified with one particular atom but are seen as delocalised over a wide range. This is a very strong type of chemical bond. Electrical and thermal conductivity as well as malleability of metals can be explained using this model.


5. Intermolecular forces

After discussing intramolecular forces such as covalent bonding, it is also important to be aware of the inter-actions between molecules, the so-called intermolecular forces. These forces are much weaker than any of the types of binding discussed previously.

van der Waals forces are the weakest forces occurring between molecules. They can be found between molecules that do not have a permanent dipole. Electrons are not static but move around, and therefore moments occur when the electron distribution is not even, leading to the formation of temporary dipoles. These temporary dipoles cause a polarisation in the neighbouring molecules. As a result, molecules are (weakly) attracted to each other.

Dipole–dipole interaction is an electrostatic interaction of permanent dipoles. In heteronuclear molecules, a polarisation of the bond is caused by the difference in electronegativity between the two atoms forming the covalent bond. This leaves a partially positive charge (δ+) on the less electronegative partner and, formally, a partially negative charge (δ) on the more electronegative partner. These molecules are attracted to each other as they can align in such way that opposite charges are next to each other, leading to intramolecular forces between the molecules (Figure 1.20).

Hydrogen bonding is the strongest of the intermolecular forces discussed here and seen as the particularly strong electrostatic interactions occurring between molecules of the type H—X. X is an electronegative atom such as F, O or N. O and N have also the advantage of possessing lone-pair orbitals. Hydrogen bonding can be seen as a strong and specialised form of dipole– dipole interaction and is the reason for the high boiling point of water (Figure 1.21).

 Figure 1.21 Hydrogen bonding in water  (Reproduced with permission from . Copyright © 2009, John Wiley & Sons, Ltd.)

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