3.1.3 Bonding - Nature of ionic, covalent and metallic bonds
Electronic configuration and energy
The noble gases have full outer shells and are stable as individual atoms. All of the other atoms have incomplete outer shells and are chemically reactive. The inference is that full outer shells confers stabiity on an atomic structure.
Reactive unstable configuration Unreactive stable configuration
Metal atoms have very few electrons in their outer shells and can gain stability by losing these electrons. The electrons cannot just be ejected into space as their removal actually requires energy. However, if the electrons are transferred to a non-metal atom, forming an ionic pair, then the resultant lowering of energy caused by the mutual attraction of the ions is more than enough to compensate for the energy required to ionise the metal.
Electrons are transferred from metal atom to non-metal atom. In the course of the process both the metal and non-metal atoms attain a noble gas configuration.
The number of atoms that react together is determined by the number of electrons that each atom needs to lose, or gain, to attain a full outer shell. Group II metals need to lose 2 electrons. Both of their outer (valence) shell electrons MUST be transferred to suitable non-metal atoms. If the non-metal can only accept one electron, then two metal atoms are needed (and vice versa).
Negative and positive ions
The electrical charge on an ion depends on the number of electrons gained or lost by an atom. This, in turn, depends on the group number of the atom in the periodic table.
Metals are found on the left hand side of the periodic table, they always form positive ions. The magnitude of the positive charge is the same as the group number of the metal.
- Group I metals form +1 charge ions, eg Li+, Na+ and K+
- Group II metals form +2 charge ions, eg Mg2+, Ca2+ and Ba2+
Non-metals are from the right hand side of the periodic table, they have nearly full outer shells and form negative ions by capturing electrons. The magnitude of the negative charge on the ion is the same as the number of electrons needed to attain a full outer shell.
- Group VI elements require 2 more electrons to attain a full outer shell, the ionic charge is 2-, eg O2- and S2-.
- Group VII elements require 1 more electron to attain a full outer shell, the ionic charge is 1-, eg F- and Cl-.
Remember that ALL compounds are neutral so the positive charges of the ions MUST be balanced by the negative charges of the ions.
Elements that are able to form positive ions transfer electrons to elements that are able to form negative ions. In general, this means reactions between metals and non-metals, but be aware that there are exceptions. When the electronegativity difference between the metallic element and the non-metallic element is small, covalent compounds can form. BeCl2, AlCl3 are both covalent compounds
The process follows these steps:
- Transfer of electrons
- Formation of positive and negative ions with full outer shells
- Electrostatic attraction of the positive and negative ions, forming a giant ionic lattice.
The ionic lattice
An ionic lattice is a regular repeating arrangement of ions in three dimensions. The ions ocupy files and rows and columns. Each negative ion is surrounded by positive ions and each positive ion is surrounded by negative ions. The actual packing structure of the ions depends on the relative sizes of the ions and their respective charges.
Lattice strength and energy
The lattice strength depends on the forces of attraction between the positive and negative ions. This electrostatic force depends on two factors:
- The magnitude of the charge on the ions
- The distance between the ions (the sum of the ionic radii)
The equation for determining the electrostatic force of attraction is:
In this equation the term is just a series of constants and can be combined together to give:
In summary, the force of attraction between two ions is proportional to the product of their magnitudes and inversely proportional to the square of the distance between them.
The strength of a lattice is a function of its lattice enthalpy, the energy required to overcome the forces of attraction keeping the ions togerther in the lattice. This is defined as the energy required to separate one mole of a crystal lattice into gaseous ions at infinite separation. For example sodium chloride lattice enthalpy is given by:
NaCl(s) Na+(g) + Cl-(g) ΔH = +790 kJ mol-1
Ions with a double charge produce lattices with a much higher lattice enthalpy.
Sodium chloride structure
Sodium chloride, NaCl, can be considered a typical ionic compound. The structure often appears in examinations and should be familiar to students. It can be used as an example of the bonding and structure of all of the compounds formed between group I and group VII elements.
The sodium chloride lattice
All ionic compounds exist in a giant ionic lattice of repeating oppositely charged ions. The formula of the ionic compound is the simplest ratio between the ions within the lattice. For this reason the term "relative formula mass" is used when dealing with an ionic compound.
- The strength of the ionic lattice is a function of both the charge on each ion and the radius of the ions.
- Greater charge = stronger lattice
- Smaller size = stronger lattice
This can be demonstrated by both the lattice enthalpy and the melting point of the lattice. The degree of ionic character is determined by the difference in electronegativity between the elements involved in the compound.
- Greater difference in electronegativity = more ionic character
- Decreasing the degree of ionic character decreases the lattice strength and melting point.
The covalent bond
Atoms are held together in covalent bonding by means of shared pairs of electrons.
This constitutes a single covalent bond. The various theories as to how and
why these bonds are formed are discussed below. It is important to recognise
that formation of covalent bonds (or any other type of bond) is an exothermic
process, one that releases energy. Similarly, to break a bond always requires
energy. Any breakage of covalent bonds involves a chemical reaction and new
substances are necessarily formed
Atoms are unstable unless they have fully occupied outer shells of electrons. We know from observation and inference that atoms bond together to make larger structures. In order to explain how this happens, different theories are proposed that explain these observations. All of these theories revolve around the accepted idea that opposite charges attract (electrostatic attraction).
The hydrogen molecule
The hydrogen molecule is the simplest structure formed between atoms. In this case two hydrogen atoms share one pair of electrons between them, forming a diatomic molecule.
- The negative charge clouds (atomic orbitals) overlap, placing a region of negative charge between the two hydrogen nuclei.
- The electrons in each atom is also attracted to the nucleus of the neighbouring atom.
- The atoms are pulled together.
This model works very well for a simple picture of bonding between atoms. The nuclei of the atoms are held to one another by the positive - negative - positive attractions caused by electron pairs between the nuclei.
Full outer energy shells
The only atoms that occur in nature not combined into some kind of structure are the inert gases. These all have full outer (valence) shells. It seems to be that this particular electronic configuration is stable for some reason.
|argon||2, 8, 8|
|krypton||2, 8, 18, 8|
|xenon||2, 8, 18, 18, 8|
Covalent bonding occurs between atoms so that they can attain a full outer shell by SHARING electrons.
In the case of hydrogen molecules, above, the atom would have a full outer shell if it were to attain two electrons. By an arrangement in which two atoms share one pair of electrons, each atom achieves the requirement. The system is now stable, it does not change further.
Covalent bonding usually occurs between non-metal atoms; they attain a full outer shell of electrons by sharing electrons. However, as we shall see in the following section there are exceptions.
Summary of covalent bonding
- Covalent bonding happens when non-metal atoms combine
- Electron pairs are shared to give a full outer shell of valence electrons to each atom
- The final units are molecules of atoms joined together by shared pairs of
Dative covalent bonds
Usually, the two atoms involved in a covalent bond provide one electron each to make the pair. However, occasionally, both of the electrons come from one of the atoms. In this case, the bond is said to be dative (= giving) covalent. The fact that an electron pair is dative has no influence on the final structure.
In ozone the central oxygen atom is bonded to one of the other oxygen atoms by a dative covalent bond (blue pair).
All of the oxygen atoms have a full outer shell (octet) of electrons.
Double bonds consist of two shared pairs of electrons between the bonded atoms.
The two single bonds are actually different in character and are given names to differentiate them, sigma and pi bonds. A double bond always consists of 1 and 1 pi bond.
Electron deficient molecules
There are exceptions to the octet rule, such as NO and NO2, in which the nitrogen atom does not have a full outer shell. The molecule is said to be electron deficient, as it is missing an electron.
Nitrogen(IV) oxide has only seven electrons in its outer shell. It is electron deficient. Molecules such as this are usually very reactive.
Electron deficient molecules can be identified by counting up all of the available valence electrons. If they add up to an odd number then there MUST be one electron left over somewhere, as an unshared single electron.
In the case of the nitrogen(IV) oxide molecule, the formula is NO2. Nitrogen is from group V and has five valence electrons. Oxygen is from group VI and has six valence electrons.
Total valence electrons = 5 + 6 +6 = 17
This is an odd number, so there has to be an unpaired electron somewhere.
The strength of the bond depends very much on the atoms being bonded. However, carbon forms single, double and triple bonds with other carbon atoms and we can use this to determine the relative strength of these bonds.
Single bonds use one pair of electrons to hold atoms together. Greater force of attraction between the electron pairs and the two nuclei draw the atoms closer together reducing the bond length
Delocalisation of electrons
Metal atoms, in common with all other metals apart from the noble gases, cannot exist for very long on their own. Metal atoms aggregate and attract one another in an attempt to stabilise themselves.
Metal atoms have very few electrons in the outer shell (valence electrons) and so cannot achieve a full outer shell by gaining electrons or sharing electrons. They tend to lose electrons, transferring them to non-metal atoms. However, in the absence of non-metal atoms, the only way that they can achieve stability is by sharing all of the outer shell electrons in giant delocalised orbitals. It is these delocalised orbital electrons that give metals their unique characteristics
Formation of ions
Losing the outer electrons to a large delocalised orbital leaves the metal atoms as ions. These ions are then held in place by the attraction of the negative charge in the delocalised orbital. The ions themselves are arranged in a giant lattice (network).
The charge on the ions depends on the number of outer shell electrons. Group 1 metals provide one electron per atom to the delocalised orbital and the ions formed have a 1+ charge. Group 2 atoms have ions with a 2+ charge
The metallic bond
The sea of electrons is a negative charge cloud that attracts all of the positive ions. It is rather like marbles stuck into blu-tack. The metal ions would repel each other without the electron charge cloud, however the force of electrostatic attraction between the electrons and the positive ions holds the whole structure together.
The strength of metallic bonding is a function of the number of electrons provided by the atoms and the consequent charge on the metal ions. The ionic radius also plays a part, as smaller ions exert a greater force of attraction on the negative charge cloud.
- Increasing ionic charge = stronger metallic bonding
- Decreasing ionic radius = stronger metallic bonding
The effect of these two factors can be seen by comparing the melting points (the temperature needed to overcome the forces within the metal structure) down group 1 and across the third period.
|Group 1 metals||Li||Na||K||Rb||Cs|
|m.p. / K||454||371||337||312||302|
It is clearly seen that as the ionic radius increases so the melting point decreases. Caesium would be a liquid on a warm summer's day.
|Period 3 metals||Na||Mg||Al|
|ionic radius /nm||0.098||0.065||0.045|
|m.p. / K||371||922||936|
Although magnesium has a similar radius to lithium, the melting point is far higher, indicating that the effect of doubling the ionic charge is much more significant.
Aluminium has a higher melting point than magnesium although not such a diference as between lithium and magnesium. It is thought that the high charge density of the aluminium 3+ ion pulls electron density back onto the aluminium ions effectively decreasing their ionic charge.
Aluminium is known to do this in its compounds, giving them a high degree of covalent character, so it seems reasonable that similar effects apply to the metallic bond.
The metallic lattice
As the metal ions in a lattice of a metallic element are all the same radius they can easily pack together like marbles in a bucket. One type of arrangement is called hexagonal close packing. It is the most efficient way for spheres to pack close together.
There are two main close packing systems, depending on how the third layer is placed in comparison to the other two. These two packing systems are called ABA and ABC. If the ions of third layer are directly above those of the first layer, it is called ABA. If the ions of the third layer sit in 'holes' that are not directly above any other ion the packing is called ABC. The best way to visualise this is using models