School of Education and
Humanities
Faculty of Science & Technology
Athrofa Addysg Uwch Gogledd
Ddwyrain Cymru
North East Wales Institute of Higher Education
Wrexham,
NORTH WALES
These notes are based (loosely) on my series of lectures given in the Introductory Chemistry module (Biology, Environmental Science degree) and in the first year inorganic unit of the Higher National Certificate course at the North East Wales Institute, Wrexham. These notes should be seen as supplementary to those lectures and not as replacing them ! These notes are designed to give a quick guide to bonding - there is a lot more to this subject than is covered here and I would love it if these brief notes encouraged you to study the subject further but, hey, life is short, so if these help you then great. If you find them useful or have any comments to make then please e-mail me. You might also want to take a look at the more advanced material available here - links at the end of these notes.
Chemical bonds form to lower the energy of the system, the components of the system become more stable through the formation of bonds. Everything wants to be more stable - its easy to lie down than stand up, bonding is Nature's way of allowing the elements to lie down.
There are several types of chemical bond. Initially we will concern ourselves with three basic types, ionic bonding, which involves the transfer of electron(s), covalent bonding, which involves sharing of electrons and metallic bonding, which in some ways can be considered as a combination of both.
If we examine the periodic table, we find that the elements in Group VIII (or 18), helium, neon, argon and so on, are particularly stable, so much so that they were once labeled the "inert gases". We now know that these elements are not inert, indeed xenon forms a range of compounds, but, nevertheless, they are very stable (although now we refer to these elements as the noble gases). This stability is the result of their electronic configuration, they have a full valence shell of electrons (ns2, np6) and this imparts stability. G. N. Lewis (1916) suggested that bonds (covalent) formed to enable elements to attain this "noble gas configuration". We can extend this idea to ionic compounds, in a compound such as sodium chloride, one element loses electron(s) to gain this stable electronic configuration whilst the other gains electron(s) to achieve the same result.
We can see that in each case, sharing or transfer of electrons, results in a more stable system.
So can we predict the type of bonds which will form between different elements in the periodic table ? The answer is yes (usually). The elements to the left of the periodic table (Groups I and II) can achieve the noble gas electronic configuration by losing electron(s). Ionization enthalpies show that lose of these is relatively easy (ionization enthalpies are fairly low for the elements in these groups). These elements are termed "electropositive". The opposite is true of the elements to the right of the periodic table. Elements in groups VII and VIII (17, 18) can most easily attain the noble gas configuration by gaining electron(s). These are the electronegative elements. (A measure of their ability to accept electrons can be seen in the electron affinity values for these elements).
This information allows us to predict that compounds formed by the combination of an electropositive element with an electronegative element will involve the transfer of electron(s) from the electropositive element to the electronegative element and hence ionic bonding will occur. The degree of ionic bonding will depend upon how extreme the differences are (in terms of electropositive / electronegative character) between the particular elements. Electropositive character increases down the group whilst electronegative character increases up a group. Combination of the most highly electropositive elements (e.g. Cs, Fr) with the most electronegative element (F) will result in the "most ionic" compounds, other combinations (such as Li with Br) will be much less ionic. In other words, transfer of electron(s) is never total, for example CsF is approximately 97% ionic, there is always some degree of covalency (however small).
When we get combination between two electronegative elements, the result is a sharing of electrons and so covalent bonds form. This is best illustrated by the combination of two atoms of the same element (a homonuclear diatomic). Combination of two fluorine atoms to produce F2, for example is achieved by sharing electrons, each fluorine atom donating one electron to form a two electron (two centre, two atoms) bond. The covalent bond is formed by overlap (or combination) of atomic orbitals of each element to form a molecular orbital. The formation of molecular orbitals lies at the heart of Molecular Orbital Theory and is the subject of another set of notes. In simplistic terms, we can visualise the sharing of two electrons between atoms as forming a single bond between the two atoms, joining them together. If the atoms share two electrons each (i.e. 4 in total) we have a double bond which will be stronger (and shorter) than a single bond. Likewise, sharing six electrons gives us a triple bond.
For combination of two electropositive elements we have metallic bonding. Perhaps the best way to model this is with "band theory" but that is the basis of other notes (when they are written). For now, we can simply visualise this as each metal donating electron(s) to a common "sea" of electrons which are shared by all the ions within the solid.
We can make some general statements about the properties of compounds based on the bonding present. These are general statements and so you will find exceptions to these descriptions but they should prove useful for a basic understanding of the chemistry involved.
Ionic compounds are generally solids at room temperature and have high melting and boiling points. They are hard but brittle solids and are poor conductors of electricity in the solid state (good conductors when molten or in solution). To explain these properties we must examine the nature of the bonding in these compound. Ionic bonds are "strong" and omnidirectional (i.e. the ions are attracted to ions of opposite charge in all directions). The ions are (to an extent) fixed in their position within the crystal lattice and have difficulty moving from these positions, being held in position by their attraction to the surrounding ions. If the ions are moved (say through physical force) this attraction is broken and alignment of like charges can occur resulting in repulsion - hence the solid is brittle. The observation that electrical conduction (however low) does occur is caused by defects in the crystal structure and this subject forms part of another series of my lectures (catch them if you can !). In solution (or when molten) the ions are able to move more easily and hence conduction (through the movement of these ions) is possible. As a side note, there are compounds which are essentially ionic which do conduct electricity rather well. These are called solid electrolytes, this is a rather speciallised area and I don't want to confuse you by spending time on this subject - if you are interested look up "silver iodide" or "ß-alumina" in a solid state chemistry text book.
In general, covalent compounds have the property of being boring at least in terms of their physical properties. They represent most of that branch of chemistry known as organic chemistry, which perhaps explains a lot. I’m joking, organic chemistry is very interesting, especially if you are an organic chemist (I am not!). Covalent compounds are not boring. Maybe a little dull, but definitely not boring. What can we say about covalent compounds ? In general they are gases or liquids at room temperature.
That about says it.
Okay, one or two are solids because of
extended 3 dimensional networks (diamond for example) but most are not solids
(and in any case, most of these interesting ones belong in inorganic
chemistry!). I suppose there’s some interest in polar covalent compounds (polar
means a little ionic) but really that’s about it as far as properties go. There
is some interesting chemistry of covalent compounds (I suppose) and we will look
at that later but for now if you want to find out more about the properties then
look elsewhere (ask an organic chemist - they are instantly
recognisable).
Metals are much more interesting. They are generally solids at room temperature. They conduct electricity and heat well and are malleable. They have high tensile strength and are usually hard. To explain these properties we must examine the bonding present. We can visualise the bonding in terms of the valence electrons of the atoms being shared throughout the metal structure. The model that tends to be applied is Band Theory but for now we can imagine the metal ions held together by this "sea" of electrons. This allows the metal to be bent and distorted without the structure breaking. These electrons are also relatively free to move and this explains the high electrical conductivity of metals.
We have looked briefly at the three major types of bonding but there are other, in many ways, just as important types of bonding.
If a atom has a lone (non-bonding) pair of electrons, this can be donated to form a dative covalent bond (provided the receiving atom can accept them). Atoms (or molecules) which have such lone pairs are called Lewis Bases, atoms (or molecules) which can accept such lone pairs are Lewis Acids. For example ammonia, NH3. The nitrogen has a lone pair of electrons, we can represent this as :NH3, ammonia is a Lewis Base. Boron trifluoride, on the other hand is a Lewis Acid, the boron has a vacant p orbital (there are six valence shell electrons around the boron, two short of the "stable octet"). The boron can accept the lone pair of electrons from the nitrogen and we get the adduct F3B-NH3. (Strange word that adduct - wonder what it means? Why not look it up?)
I have already suggested that our models for covalent and ionic bonds are idealised and that in reality we have bonds which are intermediate between these two types. If we take HCl for example, we might suppose that there is a covalent bond between the hydrogen and the chlorine. In fact the chlorine atom is more electronegative than the hydrogen atom and so attracts more of the electron density than does the hydrogen. The upshot of this is that the chlorine becomes slightly negative (i.e. has more than an equal share of the electrons) and the hydrogen becomes slightly positive (i.e. has less than an equal share). The bond becomes slightly "ionic" in nature and the molecule becomes polar (one end is slightly negative, the other slightly positive). It is fairly simple to decide whether or not a diatomic (two atom) molecule is polar, all we need do is judge whether or not there is a significant difference in the electronegativities of the two elements present. When we get multi-atom molecules the situation is a little more complex and we have to take into account how the atoms are arranged. Carbon tetrachloride (CCl4) for example, each C-Cl bond is polar since the chlorine is more electronegative than the carbon, but the chlorine atoms are arranged symmetrically about the carbon (tetrahedrally arranged) and so overall the molecule is non-polar. Before we can predict whether or not such a molecule is polar we have to know, or be able to predict, the shape of the molecule and this topic forms another series of notes.
The effect of molecule polarity upon the physical properties of a substance can be quite startling. What do you think might happen between different molecules of a substance when that molecule is polar ? The answer lies in the attraction between non-like (opposite) charges. Opposite charges will attract one another and so in polar substance molecules will associate with one another, this is called dipole-dipole interaction. What effect do you think this will have on, say, boiling point ?
One very important example of dipole-dipole interaction has been given its own name, hydrogen bonding. An every day example of this is in water. The oxygen - hydrogen bond is polar, oxygen being the more electronegative element. The molecule is therefore polar (the molecule is not linear but has a bent, V, shape). This is extenuated by the two lone pairs of electrons on the oxygen atom. One end of the molecule is partially negative whilst the two hydrogen atoms become partially positive. The molecules of water are attracted to one another, with the slightly positive hydrogens attracted to the negative "ends" (the oxygens) of other water molecules. This intermolecular attraction is termed "hydrogen bonding", and acts almost like a glue holding the molecules of water together. In the case of water the effect on the physical properties of water are quite astounding, the boiling point of water, for example, is very much greater than would be the case if such bonding did not exist. This fact alone should make the human race (and the rest of life) grateful for hydrogen bonding since water would otherwise be a gas at room temperature. Can you think of another physical property which would be strongly influenced by this "glue"?
It is possible to liquefy the elements known as the noble gases. Maybe you don’t find this surprising, cool anything and it will liquefy. But what holds the atoms of these elements together ? There must be some form of attraction.
Imagine an atom. The atom is not a static thing, electron density fluctuates around that atom. At some point in time it is possible that the electron density on one side of the atom is greater than on the other side. We have an instantaneous dipole. If the mixture is sufficiently cold (the atoms are not moving quickly) then an adjacent atom will experience this dipole and will itself establish a dipole in response. We now have an attraction between the atoms (instantaneous dipole - induced dipole interaction). Such attraction is not strong, rather it is very weak, but nevertheless it exists.
Intermolecular forces of attraction, those not involving ions, are generally referred to as van der Waal’s forces.