The Chemical Bond

Characteristics ionic compounds

. Crystalline solids at room temperature.

2. They have HIGHER melting points and boiling points compared to covalent compounds.

3. They conduct electrical current in molten or solution state.

4. They are extremely polar bonds.

5. Most of them are soluble in water but not soluble in non-polar solvents.

Covalent Bonds

Covalent bonds are formed as a result of the sharing of one or more pairs of bonding electrons. Each atom donates half of the electrons to be shared. This sharing of electrons is as a result of the electroneagativity (electron attracting ability) of the two bonded atoms are either equal or the difference is no greater than 1.7 then the higher electronegative atom has an electron attracting ability large enough to force the transfer of electrons from the less electronegative atom. This would then be an IONIC BOND. As long as the electronegativity difference is no greater than 1.7 the atoms can only share the bonding electrons.

SINGLE COVALENT BOND

A single covalent bond would be the sharing of two electrons between the two bonded atoms. Examples of this are:

H-H

H-Cl

F-F

(The dash is to represent a bonding pair).

DOUBLE COVALENT BOND

A double covalent bond is two pairs of electrons being shared.

Examples are:

O=O

C=O

C=C

TRIPLE COVALENT BOND

A triple covalent bond is the sharing of three pairs of electrons.

Examples are:

TRIPLE BOND BETWEEN TWO NITROGEN ATOMS

TRIPLES BOND BETWEEN TWO CARBON ATOMS

TRIPLE BOND BETWEEN A CARBON AND A NITROGEN

The Polarity of the Covalent bond

Two atoms with the same electronegativity will share the bonding pairs equally. As a result the bonding electrons will be evenly distributed between the bonded atoms. There will be no accumulation of bonding electrons on any one atom and the bond dipole will be zero. Such a covalent bond will be called a non-polar bond. The bond between two Hydrogen as in two Oxygens as in or two Hydrogens like will all be non polar bonds.

However, if the two bonded atoms have a different electronegativity then the bonding pairs of electrons will be shared unequally. The atom with the higher electronegativity will attract the bonding electrons closer to itself. As a result, the electron distribution will be unequal and a bond dipole moment will be formed. For example the single bond between a Hydrogen and a Chlorine as in H-Cl will have the bonding pair closer to the higher electronegative atom (Chlorine). As a result the Chlorine end will be partially negative since the electrons are closer to the Chlorine. The hydrogen end will be partially positive since the bonding pair is further from the Hydrogen. This two pole condition is called a dipole and it generates a dipole moment.

Co-ordinate bonding (Dative covalent bonds)

A Covalent bond is formed by two atoms sharing a pair of electrons. The atoms are held together because the electron pair is attracted by both of the nuclei.

In the formation of a simple covalent bond, each atom supplies one electron to the bond - but that doesn't have to be the case. A dative bond is a covalent bond (a shared pair of electrons) in which both electrons come from the same atom.

Example: THE REACTION BETWEEN AMMONIA AND HYDROGEN CHLORIDE.

If these colourless gases are allowed to mix, a think white smoke of solid ammonium chloride is formed.

Ammonium ions, NH4+, are formed by the transfer of a hydrogen ion from the hydrogen chloride to the LONE PAIR OF ELECTRONS on the ammonia molecule.

When the ammonium ion, NH4+, is formed, the fourth hydrogen is attached by a dative covalent bond, because only the hydrogens nucleus is transferred from the chlorine to the nitrogen. The hydrogens electron is left behind on the chlorine to form a negative chloride ion.

Once the ammonium ion has been formed it is impossible to tell any difference between the dative covalent and the ordinary covalent bonds. Although the electrons are shown differently in the diagram (above), there is no difference between tem in reality.

Representing co-ordinate bonds

In simple diagrams, a dative covalent bond is shown by an arrow. The arrow points from the atom donating the lone pair to the atom accepting it.

Electronegativity

The affinity for electrons. The atoms of the various elements differ in their affinity for electrons.

This diagram changes the conventional periodic table of the elements so that the greater the electronegativity of an atom, the higher its position in the table.

Although fluorine (F) is the most electronegative element, it is the electronegativity of runner-up oxygen (O) that is exploited by life.

The shuttling of electrons between carbon (C) and oxygen (O) atoms powers life.

. Moving electrons against the gradient (O-C) - as it occurs in PHOTOSYNTHESIS -requires energy and stores it.

2. Moving electrons down the gradient (C-O) - as occurs in CELL RESPIRATION -releases energy.

The relative electronegativity of two interacting atoms also plays a major part in determining what kind of chemical bond forms between them.

EXAMPLE: Sodium and Chlorine = IONIC BOND.

* There is a large difference in electronegativity, so the chlorine atom takes an electron from the sodium atom.

* So, that is an ionic bond.

* The atoms become ions Na and Cl and are held together by their opposite electrical charge to form a crystal lattice of table salt (NaCl)

* Each sodium ion is held by 6 chloride ions while each chloride ion is, in turn, held by 6 sodium ions.

* So, a cubical crystal of common table salt.

Intermolecular Forces

Intermolecular forces are WEAKER than covalent bonds.

* Only 16kJ/mol of energy is required to overcome the intermolecular attraction between HCl molecules in the liquid state (i.e. the energy required to vaporise the sample).

* However, 431kJ/mol of energy is required to break the covalent bond between the H and the Cl atoms in the HCl molecule.

So, when a molecular substance changes states, the atoms within the molecules are unchanged.

The temperature at which a liquid boils reflects the kinetic energy needed to overcome the attractive intermolecular forces (likewise, the temperature at which a solid melts).

So, the strength of the intermolecular forces determines the physical properties of the substance.

Attractive forces between neutral molecules.

* Dipole-dipole forces.

* London dispersion forces.

* Hydrogen bonding forces.

Typically, dipole-dipole and dispersion forces are grouped together and termed VAN DER WAALS forces (attractive force between molecules, which is caused by the induced temporary polarisation of molecules by the dipole moments of molecules).

Attractive forces between neutral and charged (ionic) molecules.

Ion Dipole forces

Note that all of these forces will be ELECTROSTATIC in nature.

Ion Dipole.

* Involves an interaction between a charged ion and a polar molecule (i.e. a molecule with a dipole).

* Cations are attracted to the negative end of dipole.

* Anions are attracted to the positive end of a dipole.

* The magnitude of the interaction energy depends upon the charge of the ion (Q), the dipole moment of the molecule (u) and the distance (d) from the center of the ion to the midpoint of the dipole.

* Ion dipole forces are important in solutions of ionic substances in polar solvents (e.g. a salt in aqueous solvent)

Ion dipole-dipole forces

A Dipole-Dipole force exists between neutral polar molecules.

* Polar molecules attract one another when the partial positive charge on one molecule is near the partial negative charge on the other molecule.

* The polar molecules must be close to one another for the dipole-dipole forces to be significant.

* Dipole-dipole forces are characteristically weaker than ion-dipole forces.

* Dipole-dipole forces increase with an increase in the polarity of the molecule.