The fact that all substances have liquid or solid states and melting or boiling points is evidence that there are forces which bind the particles of those substances together, and which require energy to overcome. Without such forces all substances would be gases right down to a temperature of absolute zero.
We will examine each of these forces in turn, before considering in the next section what their implications are for the structure and physical properties of substances.
The covalent bond
A covalent bond is formed by a shared pair of electrons. The two electrons are localized between the nuclei of the bonded atoms, and the electrostatic attraction between the negatively-charged bonding electrons and the positively-charged nuclei holds the atoms together strongly.
Definition: a covalent bond is the electrostatic attraction between a bonding pair of electrons and the two nuclei of the bonded atoms.
Covalent dot-and-cross diagrams are used to show covalent bonds, which are typically but not exclusively found between non-metal atoms. A single covalent bond is found, for example, in a chlorine molecule where we would represent the bond in a displayed formula as Cl-Cl:
It is also acceptable to omit the outer shell and show the dot-and-cross diagram in this form:
Before the bond was formed, the electrons were found in atomic orbitals of the two atoms. Now they occupy a new region of space between the two bonded atoms, and this region of space is called a molecular orbital. When two electrons form a single covalent bond, the resulting orbital is cylindrically symmetrical (i.e. a sort of stretched sphere) in shape. We call this a sigma (σ-)orbital, and may also refer to a single covalent bond as a sigma-bond. Because the σ-bond is along the axis between the two atoms, rotation of either atom can take place without the bond being broken.
Atoms forming covalent bonds may also share two electrons each to form two covalent bonds between them (a double covalent bond), for example in an oxygen molecule, where we would show the displayed formula as O=O:
In a bond such as this where there are two shared pairs of electrons in the bond, the first pair occupies a σ-orbital along the axis of the covalent bond as shown before. The second pair occupy a molecular orbital that is formed by the sideways overlap of two p-orbitals in the bonded atoms. As a result this molecular orbital is above and below the σ-orbital, and comprises two separate regions of space in which the bonding pair will be found. We call this a pi (π-)orbital.
A double covalent bond therefore consists of a σ-bond plus a π-bond. Because this π-bond is not along the axis between the two carbon atoms, rotation is restricted. Theπ-bond would have to be broken for either of the carbon atoms to be able to rotate around the σ-bond.
Rarely, a triple covalent bond may be formed between two atoms (examples: H-C≡C-H and N≡N). This would comprise a sigma-bond and two pi-bonds with one pi-bond above and below the sigma-bond and the other one in front of and behind it.
A quadruple covalent bond would require too many negatively charged electrons in the same space between the two atoms, and does not occur.
Although it is usual for bonding to result in an electron arrangement isoelectronic with a noble gas, there are a number of atoms that can accept more than a “full” shell during covalent bond formation. While atoms of elements in the second period never have more than 8 electrons in their outer shell, those in period 3 and subsequent periods often do – this is called having an expanded octet.
SF6 sulphur hexafluoride
XeF4 xenon tetrafluoride
Incomplete outer shells
In other cases, covalent bond formation increases the number of electrons around an atom, but not all the way to a filled shell.
Example: BH3 borine
The two unbonded atoms in the N atom’s outer shell in ammonia are not just two random electrons, but a non-bonding pair of electrons called a LONE PAIR. You’ll also find two
lone pairs on each of the O-atoms of O2. These lone pairs are important – they can get involved in reactions, and they help determine the shape of molecules.
Dative covalent bonds
In some situations, one atom may contribute both electrons to form the covalent bond. The bond formed is in all respects identical to a normal covalent bond, and is called a dative covalent bond (also called a co-ordinate bond).
A dative covalent bond can be formed where one atom has a lone pair of electrons available to donate, and another atom has a vacancy for two electrons in its outer shell.
e.g. a dative covalent bond is formed when an ammonia molecule acts as a base (proton acceptor) and accepts a hydrogen ion to form an ammonium ion:
NH3(g) + H+(aq) → NH4+(aq)
This is possible because the N has a lone pair to donate and the H+ ion has an empty n=1 shell able to accept the two electrons.
You will always be able to tell a dative covalent bond in a dot cross diagram because both electrons in the intersection of shells will be the same symbol (i.e. from the same atom).
Dative bonds can also form part of a set of double or triple bonds.
The bonding in CO is :
The Ionic Bond
Definition: An ionic bond is the strong electrostatic attraction between oppositely-charged ions.
An ionic bond is formed as a result of the transfer of electrons between atoms. This transfer results in oppositely-charged ions being formed, with the electron-deficient ion having a positive charge (more protons than electrons) and the electron-rich ion having a negative charge. A strong electrostatic attraction between the positive ions and the negative ions results.
Most usually, ionic bonds would be expected between a metal ion and a non-metal ion, but there are exceptions.
Ionic dot-and-cross diagrams are used to show bonds between positive and negative ions. For the positive ion, either an empty outer shell (preferred) or the next filled shell closer to the nucleus can be shown.
We can also have ionic bonding involving compound ions, such as hydroxide, carbonate, sulphate, nitrate or ammonium ions. Inside the compound ion the bonding is covalent, however by accepting or donating electrons the compound ion as a whole becomes electron deficient or electron rich and can then form an ionic bond with an oppositely charged ion. This situation is very common in salts, since most acids are covalently bonded molecules that dissociate in solution into a hydrogen ion and a negatively charged compound ion.
Sodium hydroxide is a simple example of an ionic bond involving a compound ion…
while ammonium nitrate is a much challenging example, with a dative covalent bond inside each compound ion.
Bonds with covalent and ionic character
Ionic and covalent bonding are two extremes of the same thing. In a perfect covalent bond the two bonding electrons are equally shared between the two bonded atoms, but in most covalent bonds one atom attracts the bonding pair more strongly than the other. These are polar bonds, which we will meet again in the context of intermolecular forces. The more one atom attracts the bonding electron pair compared to the other, the closer we get to one atom simply having both bonding electrons entirely: an ionic bond.
The Metallic Bond
A different kind of bonding is found in elemental metals. Metal atoms lose their outer-shell electrons to form positive ions, but rather than transferring them to other atoms, these electrons form a ‘sea’ of delocalised electrons surrounding the metal ions.
Definition: metallic bonds are the strong electrostatic attractions between positively charged metal ions and the surrounding delocalised electrons.
A dot-and-cross diagram is not needed to represent metallic bonding: the generic labeled diagram shown here suffices.
Note that the strength of the metallic bond (and hence how high the melting and boiling point is) depends on how highly charged the metal ions are, and how many electrons they have donated to the sea of delocalized electrons. e.g. Al has a higher melting point than Na.