CCEA ADVANCED SUBSIDIARY

CHEMISTRY

MODULE 1

1.2 Bonding

Whenever atoms, ions or molecules approach each other, there are electrostatic forces acting between them. When the net forces are forces of attraction, and they are strong enough to bind the particles together, we refer to them as chemical bonds. When particles are bound together by chemical bonds, the resulting arrangement is known as the structure of the substance concerned.

All the noble gases except He have completely full inner shells and an outer octet of electrons. As these gases are monatomic and rarely enter into chemical combination it is assumed that the outer octet of electrons is a very stable arrangement and therefore when atoms combine they will try to obtain the noble gas configuration.

The three main types of bonds are:

1. Ionic or electrovalent

2. Covalent

3. Metallic

Ionic bonding restricted to elements in groups I, II, VI and VII, the ions of which have a noble gas structure. Dot and cross diagrams. Characteristic properties of ionic compounds. Sodium chloride as a typical ionic crystal.

When you have finished this section you should be able to:

· Understand that ionic bonding involves attraction between oppositely charged ions formed by electron transfer

· Describe the various electron configurations of simple stable ions.

· List properties typical of ionic compounds and indicate how useful each one is in deciding whether a substance is ionic.

Ionic Bonding

Atoms can achieve noble gas electronic configuration by loss or gain of electrons to form ions. Metals (with low electronegativities) lose electrons to form positive ions and non-metals gain electrons to form negative ions.

Sodium can attain the stable electron configuration of neon by losing one electron

Na (1s²2s²2p⁶3s¹) Na⁺ (1s²2s²2p⁶) + e^-

With ten electrons and eleven protons he species formed has a positive charge.

Fluorine is one electron short of the neon electronic configuration. If it obtains one electron (from sodium) it can achieve a full outer shell of eight electrons.

F (1s²2s²2p⁵) + e^- F^- (1s²2s²2p⁶)

The species formed has ten electrons and nine protons; it is a negatively charged fluoride ion.

x x x x x x x x

x x x x x x x

x x x x x x x x

Na atom F atom Na⁺ ion F^- ion

You will soon be aware of the limitations of these diagrams, but they are nevertheless quite useful even in A-level work for describing electron transfer and for checking that all electrons are accounted for. However, electron shell diagrams are rather tedious to draw, and give more information than we usually need when only the outer shell electrons are involved in reactions. We can simplify the above diagram to draw what we call a 'dot-and-cross' diagram showing only the outer shell electrons:

xx + xx -

Na · + x x · x

F x Na x F x

xx xx

Electrons shown by dots and crosses are of course indistinguishable - we use different symbols only to show where they come from. Note also the use of brackets to separate the charge symbol from the symbols for electrons.

In the first exercise you draw some similar dot-and-cross diagrams yourself

Exercise 1

Draw dot-and-cross diagrams for the formation of:

(a) sodium chloride from sodium atoms and chlorine atoms;

(b) magnesium oxide from magnesium atoms and oxygen atoms;

(c ) calcium fluoride from calcium atoms and fluorine atoms.

Properties of Ionic Compounds

1. Crystalline solids : Ionic compounds are crystalline solids. Ionic crystals are quite hard due to the strong electrostatic forces between the ions[GMD1] .

2. High melting and boiling points : All ionic compounds have high melting and boiling points as a large amount of heat energy is required to break the strong electrostatic forces of attraction between the ions.

3. Soluble in water : Many ionic compounds are soluble in water. The water molecules attract the ions and pull the structure apart. They also stabalise the ions once in solution

4. Conduct electricity when molten or dissolved in water : All ionic compounds conduct electricity when molten and are decomposed in the process, which is called electrolysis. Most ionic compounds dissolve in water, which frees the ions so that they can move and carry electrical charge.

Structure of sodium chloride

An ion may be regarded as an electrically charged sphere. A charged sphere is surrounded by a uniform electric field and therefore attracts oppositely charged spheres in all directions.

No particular orientation is favoured, so we often say that ionic bonding is non-directional. What we really mean by this is that the forces of attraction are non-directional.

However, when large numbers of oppositely-charged spherical ions are attracted to each other, the repulsion of the like-charges also comes into play. Mutual repulsion of similarly-charged ions limits the number which can come into contact with an oppositely-charged ion, and effectively fixes their relative positions. Ions, therefore, tend to cling together in large clusters known as ionic lattices in which attractive and repulsive forces are balanced. The particular arrangement of ions depends on their relative charges and sizes.

The ionic lattice of sodium chloride consists of a regular arrangement of alternating sodium and chloride ions extending in three dimensions:

Na⁺ Cl^- Na⁺ Cl^- Na⁺ Cl^- Na⁺ Cl^-

Cl^- Na⁺ Cl^- Na⁺ Cl^- Na⁺ Cl^- Na⁺

Na⁺ Cl^- Na⁺ Cl^- Na⁺ Cl^- Na⁺ Cl^-

Cl^- Na⁺ Cl^- Na⁺ Cl^- Na⁺ Cl^- Na[GMD2] ⁺

The lattice can be imagined as consisting of two inter-penetrating face centred cubic structures of sodium ions and chloride ions.

Na⁺ Na⁺ Cl^- Cl^-

Na⁺ Cl^-

Na⁺ Na⁺ Cl^-Cl^-

Each sodium ion is surrounded by six chloride ions as nearest neighbours in the lattice and each chloride ion by six sodium ions. The lattice is said to have 6:6 coordination.

Stereoscopic picture of sodium chloride. The large circles represent Cl^- ions and the small circles Na⁺ ions. [To view place a piece of card about 15 cm long between the two pictures so that the left eye can only see the left image and the right eye the right image. Your brain will fuse the two images into one 3-D picture]

You can see from the above figure that it is impossible to say that a particular sodium ion ‘belongs’ to a particular chloride ion, as they are an equal distance from six. This means that ‘molecules’ of an ionic compound are not formed. The formula used is the atoms in their combining ratio.

Formation of covalent bonds in terms of the sharing of electron pairs between atoms. Dot and cross diagrams. Multiple bonds exemplified by C₂H₄, N₂ and CO₂. The octet rule and its limitations, eg BeCl₂, BF₃. The coordinate bond as a special case of the covalent bond eg NH₄⁺. Dot and cross diagrams are required to show only the outermost electrons but, where appropriate, charges should be included.

When you have finished this section you should be able to:

· Describe covalent bonding in terms of shared pairs of electrons.

· Draw dot-and-cross diagrams for a variety of covalent molecules.

· Explain the term dative covalent bond.

· Draw dot-and-cross diagrams and structural formulae including dative bonds.

· For polyatomic ions, distinguish between the covalent bonding within the ion and the ionic bonding between ions.

· Draw dot-and-cross diagrams for at least two covalent compounds in which an atom other than hydrogen has fewer than eight outer shell electrons.

· Draw dot-and-cross diagrams for at least two covalent compounds in which an atom has more than eight outer shell electrons.

1.2 Covalent Bonding

Covalent bonds are formed by the sharing of electrons between atoms. The two atoms have to approach sufficiently close to each other for their atomic orbitals to overlap. The shared pair of electrons constitutes a single covalent bond and occupy the same orbital with opposing spins.

Dot-and cross Diagrams

Dot-and cross diagrams are simplified versions of the diagram for the chlorine molecule on the left below, showing only the outer shell electrons

Note that the use of two (or even more) different symbols for the electrons does not mean that the electrons are different - it simply identifies the ‘parent’ atoms.

You may also encounter dot-and-cross diagrams with bond lines added, and also versions showing only the bonding electrons, such as

or Cl Cl

Multiple Covalent Bonds

Multiple covalent bonds may involve the sharing of two or even three pairs of electrons between atoms.

x · ·

or N x N or N N

x · ·

Exercise 2

Draw dot-and-cross diagrams and structural formulae for the following molecules:

(a) H₂ (d) CH₄ (g) C₂H₄ (j) O₂

(b) HI (e) H₂O (h) CHCl₃ (k) CH₃OH

Exercise 3

(a) How many electrons are associated with the outer shell of each atom in the molecules in Exercise 2?

(b) You have learned in your pre-A-level course that hydrogen has a 'valency' or 'combining power' of 1, oxygen 2 and carbon 4. Explain these values in terms of your answer to part (a).

Dative Covalent Bonds (Coordinate Bonds)

A covalent bond is formed when two atoms share an even number of electrons. Usually each atom donates an equal number of electrons to form the bond. A dative covalent or coordinate bond is formed between atoms when one of the atoms donates both electrons (called a lone pair) to the bond.

The dative covalent bond is represented by the symbol . The arrow originates from the atom donating the electron pair (the donor atom) and points to the atom receiving the electron pair (the acceptor atom).

e.g. NH₃ + H⁺ NH₄⁺

Exercise 4

Draw dot-and-cross diagrams and structural formulae for the following molecules and ions. In the structural formulae use an arrow to represent a dative bond.

(a) H₃O+ (i.e. H₂O linked with H⁺)

(b) BF₃NH₃

(d) NO₃^-

(e) H₂SO₄

The Octet Rule

When an atom forms a chemical bond by gaining, losing or sharing electrons, its electronic configuration becomes the same as that of a noble gas either at the end of the same period or at the end of the previous period in the Periodic Table. Noble gases, with the exception of helium, have a stable octet of electrons in the outer shells of their atoms. The formation of chemical bonds to achieve noble gas configuration is called the 'octet rule' and applies to both ionic and covalent compounds.

When elements with fewer than four outer shell electrons per atom form compounds they usually lose those electrons to form ions. However, for small atoms, the relevant ionisation energies may be so high that covalent bonding occurs instead. Since there are fewer than four electrons available for sharing, there will then be fewer than eight outer shell electrons per atom in the resulting compound. Such compounds are often called 'electron-deficient', but you should not imagine there is anything wrong with them!

Exercise 5

Draw dot-and-cross diagrams for

(a) beryllium chloride, BeCl₂

(b) boron trifluoride, BF₃

How many outer shell electrons are there around the central atom?

The molecules above certainly do exist in certain conditions. There is a tendency for the ‘stable octet’ to be reached by means of dative bonding between the neighbouring molecules.

We will now consider some examples where the stable octet is exceeded, so called ‘expanded octet’ compounds.

If the octet rule always applied , no atom could have more than eight electrons in its outer shell. As a result no atom could have more than four covalent bonds (or the equivalent) associated with it. The octet rule applies to elements in the first two periods of the Periodic Table (H to Ar) but many other elements show a covalency greater than four in some of their compounds. The following formulae represent some examples.

PF₅ SF₆ PCl₆ ^-

Exercise 6

Draw dot-and-cross formulae and count the electrons around the central atom in the following:

(a) PF₅

(b) SF₆

The ability of elements in the third period to have more than eight electrons in the outer shell is due to the presence of empty 3d-orbitals.

e.g. phosphorus

Ground state

3s 3p 3d

3 unpaired electrons, normal compound PCl₃.

One of the 3s electrons can be promoted to the 3d orbital.

Excited state

3s 3p 3d

5 unpaired electrons, PCl₅. This tends to be unstable decomposing to form PCl₃ and Cl₂.

The two types of bonding, ionic and covalent, represent two extreme cases. In practice very few bonds are purely ionic or purely covalent.