AS unit foundation bonding page


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Foundation - Bonding & Structure

Ionic Forces

Definition
Dot & cross diagrams

Covalent Bonds

Definition
Dot & cross diagrams
VSEPRT

Electronegativity
Intermolecular forces

Van der Waals forces
Dipole-dipole forces
Hydrogen bonds

Bonding & Structure - Ionic Forces

(1) Definition :

When an atom, generally a metal atom, loses one or more of its electrons, a positive ion is formed.

e.g. Na

Na⁺ + e^-

When an atom, generally a non-metal atom, gains one or more electrons, a negative ion is formed.

e.g. Cl₂ + 2e^-

2Cl^-

Just as with the opposite poles of a magnet, when positive and negative ions approach one another a very strong force of attraction is formed. This is sometimes called an ionic bond, though more accurately it is a strong ionic force of attraction.

e.g. Na⁺ + Cl^-

NaCl

(2) Dot & cross diagrams :

Below are two dot and cross diagrams for the valence (outer) shells of an atom of sodium and chlorine. They show that sodium has one electron in its outer shell and chlorine has seven,


Na, 1s²2s²2p⁶3s¹	Cl, 1s²2s²2p⁶3s²3p⁵

When the two elements react together to form an ionic compound, the sodium donates its one spare electron to the chlorine atom. The seven electrons of the chlorine atom become eight, a full 3^rd shell, and the sodium atom also has a full outer shell (the 2^nd shell).

Below are dot and cross diagrams representing the two ions formed,


Na⁺, 1s²2s²2p⁶	Cl^-, 1s²2s²2p⁶3s²3p⁶

N.B.: When drawing dot and cross diagrams of ionic compounds the ions must not touch each other; hence the use of square brackets above to isolate each ion in space.

(3) Lattice structure :

An ionic compound is formed by massive numbers of positive and negative ions. The ions form a very large structure, with positive surrounded by negative ions, and negative ions surrounded by positive ions, in what is called a giant lattice structure.

e.g. for sodium chloride, NaCl -

Bonding & Structure - Covalent Bonding

(1) Definition :

A covalent bond put in its simplest terms is the sharing of two electrons by two atoms. In a normal covalent bond (e.g. C-H, C-C, O-H, etc..) each atom forming the bond will donate one electron to the bond.

If the covalent bond is formed by one of the atoms donating both electrons to the bond then it is called a dative covalent (or co-ordinate) bond. The best example of this is one of the N-H bonds in the ammonium ion, NH₄⁺,

(2) Dot & cross diagrams :

(i) Hydrogen -

(ii) Chlorine -

(iii) Hydrogen chloride -

(viii) Carbon dioxide -

(ix) Ethene -

(3) Valence Shell Electron Pair Repulsion Theory :

Electrons are negatively charged and when they are in covalent bonds the different bonds all try to repel each other. When most atoms form covalent bonds they do so to fill their outer or valence shell of electrons. This generally means an atom desires eight electrons (the so-called eight electron rule) in its valence shell.

This leads to four pairs of electrons being in the valence shell of atoms. These four pairs of electrons can lead to different molecular shapes :

methane, CH₄ (4 bonding pairs) :

With methane the four pairs of electrons are all bonding pairs and so are all of equal strength. The maximum separation of the four pairs in 3D is with an internal bond angle of 109.5°. The shape is a tetrahedron.

ammonia, NH₃ (3 bonding pairs and 1 lone pair) :

The three bonding pairs of electrons are of equal strength and the fourth pair of electrons is a lone pair. The repulsion between the lone pair and bonding pairs is greater than the repulsion between bonding pair and bonding pair. Therefore the H-N-H internal bond angle is 107° and its shape is trigonal pyramidal (a triangular-based pyramid).

water, H₂O (2 covalent bonds and 2 lone pairs ) :

The two bonding pairs of electrons are of equal strength and the other pairs of electrons are lone pairs. The repulsion between the lone pairs and bonding pairs is greater than the repulsion between bonding pair and bonding pair. Therefore the H-O-H internal bond angle is 104° and its shape is bent.

There are some more advanced examples of molecules which don't obey the standard 8 valence electron rule :

boron trifluoride (3 bonding pairs) :

The boron atom is satisfied with only six electrons in its valence shell in compounds. These three bonding pairs of electrons form a trigonal planar shape and have a bond angle of 120°.

phosphorus pentachloride (5 bonding pairs) :

The phosphorus atom in covalent compounds can expand its octet. That is, it can utilize the empty 3d orbitals in its bonding, enabling ten valance electrons to be permitted. These five equal bonding pairs of electrons have a trigonal bipyramidal arrangement, with two different internal bond angles -; 120° in the plane and 90° between the plane and the axial atoms (see diagram below).

sulphur hexafluoride (6 bonding pairs) :

As with the phosphorus atom, the sulphur atom can also expand its octet. It can utilize the empty 3d orbitals in its bonding, this time enabling twelve valence electrons to be permitted. These six equal bonding pairs of electrons will have an octahedral arrangement with an internal bond angle of 90°.

(3) s- and p-bonds :