Electronic Theory of Valency

(GCSE Level recap)

The octet rule states that all compounds want to attain the stable noble gas structure with 8 electrons in the outer shell. In ionic bonding they gain and loose whole electrons, in covalent bonding they share electrons to gain a share in the stable noble gas structure. This works reasonably well in elements up to calcium. However, there are exceptions.

Ions such as Cu2+, Fe2+ and Fe3+.do not have 8 outer shell electrons. SO2 has 10 electrons in its outer shell, whereas BeCl2 has only four.

Ionic Bonding

Where electrons are completely lost and gained to form +ve ions (Cations) and -ve ions (Anions).

DH = the amount of energy required.

The energy needed to loose an electron is the ionisation energy. The energy needed to loose one electron is the first ionisation energy, the second electron loss is the second ionisation energy. These are different values, and change between the different elements.

e.g. Ca ® Ca+ + e- First I.E.
Ca+ ® Ca2+ + e- Second I.E.

I.E. 2 requires more energy as the positive charge on the ion attracts the remaining electrons and the size of the ion shrinks.

Factors Influencing Ionic Bonding

1. How easy is it for the metal to loose electrons and to form cations.

This depends on the ionisation energies - A lower I.E. leads to the electron being lost more easily and the ion being formed more easily.

Going down the first and second groups, the ionisation energies decrease, leading to an increasing tendency to form ions. This is because of shielding of the outer electrons from the nucleus by the inner electrons, etc. (see GCSE notes)

2. How easy is it for the non-metal to form an anion.

eg. Cl + e- ® Cl-
This energy change is called the electron affinity. This is the amount of energy required to form the ion. Therefore, the lower the electron affinity, the more likely the ion is to form. Many electron affinities are -ve ie. Energy is given out when the electron is taken in.

3. A better ionic bond forms when:

The metal is a big ion with a low charge (e.g. Cs+)
The non-metal is a small ion with a high charge (e.g. O2-)

If the non-metal is big then the electrons furthest out from the nucleus can be attracted away from the nucleus by the pull of the positive ion. This happens especially if the positive ion is smaller, in which case the distance between its nucleus and the electrons of the anion is smaller. This is called polarisation (see photocopied sheet) and leads to some covalent character being introduced into the bond.

4. Lattice Enthalpy.

The creation of the two ions (cation and anion) often takes up more energy than it releases (is endothermic.) The energy required for this is released when the two ions come together to form the ionic compound. The energy released is called lattice enthalpy (energy.) This reaction is always exothermic - the greater the amount of energy released, the more likely the ionic bond is to form.

Theoretical Value

Using the laws of electrostatics, the known sizes of the ions, the charges on the ions and assuming a perfect ionic lattice, theoretical values for the lattice enthalpy can be calculated.

Experimental Value

The value of the lattice enthalpy can be worked out from experimental values of electron affinities, ionisation energies, heats of formation, etc. A Born-Haber cycle is used to calculate the result (See module 3).

If there is a good agreement between the theoretical and experimental lattice enthalpies, then the substance is perfectly ionic (the ions are spherical and electrons are completely transferred.) If there is less agreement, then the ion has some slight covalent character creeping in due to polarisation.

Covalent Bonding.

Sharing of two electrons between two atoms is a single covalent bond, or a sigma (s) bond. Sharing of 4 electrons is a double covalent bond, and consists of a s bond and a p bond. Sharing of six electrons is a triple covalent bond (A s bond and two p bonds).

In a sigma bond, two electron orbitals overlap to form a cigar shaped orbital between the two atoms, as shown in the diagram below:

Example: the H-H bond

In a pi bond, two p orbitals overlap to form two separate clouds of electron density, containing two electrons. The electrons can move between the clouds and the two do not necessarily need to be in different clouds at the same time. This is shown in the diagram below.

Example: The C=C bond.

Electronegativity

In a perfect covalent bond, the pair of electrons is shared equally between the two atoms involved. However, this is not always the case. When one atom attracts the electrons more strongly, it pulls some of the charge towards itself, and thus causes a polar bond. This is shown using lower-case delta symbols, or arrows:

Cd+----Cld- or (less common) C-®--Cl

This shows that the electrons have been more strongly attracted by the chlorine than the carbon. The extent to which an atom attracts another atom in a covalent bond. Small atoms can come closer to the bonding electrons, and therefore are more electronegative. The most electronegative elements are found in the top right hand corner of the periodic table - the least in the bottom left. When a compound is made with a very electronegative atom attached to a very non electronegative atom, the bond is very polar (more ionic character starts to creep into the bond.) This leads to the molecule formed having positive and negative ends (e.g. water), and it also makes the bond length shorter.

Dative Covalent Bonding (Co-ordinate bonding)

A dative covalent bond is a covalent bond where both of the electrons come from the same atom. The bond then functions as a normal covalent bond. The most common examples follow.

  1. The ammonium ion.

Ammonia + Hydrogen ® Ammonium+
NH3 + H ® NH4+ + e-

One of the bonds in the ammonium ion is a dative covalent bond - this bond is formed with both electrons coming from the nitrogen.

2. BF3 will combine with ammonia -

3. Aluminium chloride - Al2Cl6 molecules form in the vapour state.

4. Also in complex ions - e.g. Cu(NH3)4

5. Nitric Acid (HNO3)

6. Carbon Monoxide (CO)

General principles for working out covalent bonds:

1. When bonding, electrons always bond in pairs. There are never an odd number of electrons shared between two atoms. There are single, double and triple bonds, which involve sharing respectively 2, 4 and 6 electrons. One line on a bonding diagram represents two electrons, one coming from each of the two atoms bonding, being shared between the two atoms. If the line has an arrow on it, the electrons both come from the same atom.

2. Elements never bond to themselves, always to atoms of other elements. The exceptions to this rule follow:

(a) Carbon bonds to itself in almost all organic compounds.
(b) Silicon bonds to silicon in silicates.
(c) Oxygen bonds to oxygen in peroxides.
(d) Sulphur occasionally bonds to sulphur, but very rarely in small molecules, except S2O3.

3. There are very few compounds with small rings of atoms in them, except ones where the name begins with cyclo.

4. Only use dative covalent bonds when absolutely necessary.

General principles for working out ionic bonding:

  1. Draw both atoms in the state that they are in before the reaction takes place.
  2. Put in arrows to show the movement of the electrons.
  3. Put square brackets round the atons when drawn for the second time, and but the charge on each outside the brackets. Put numbers of ions on the left hand side of the brackets.

Inter-Molecular Bonding (Bonds between Molecules)

There are three types of intermolecular bonds

  1. Hdrogen bonds
  2. Dipole - Dipole Bonds
  3. Van der Waalls forces

Hydrogen Bonding

See the graph of the boiling points of hydrides for evidence for the existence of hydrogen bonding.

Looking at the boiling points of non metal hydrides, NH3, H2O and HF have anomolously high boiling points. This is because they each contain hydrogen bonds.

Points about hydrogen bonds:

Hydrogen bonds are the strongest intermolecular bonds - however, they are not as strong as normal covalent bonds. Hydrogen bonding explains:

Water is very strongly hydrogen bonded. It has 2 hydrogen atoms joined to oxygen atoms and 2 lone pairs of electrons per atom, giving it the possibility of 2 hydrogen bonds per molecule. The structure formed is similar to that of diamond.

Dipole - Dipole Bonding (Polar Bonding)

If a molecule contains a polar bond that is asymmetrically positioned the molecule overall will have a dipole (a +ve and a -ve end.) This will lead to an attractive force between the molecules.
Note that not all molecules with polar bonds lead to polar molecules. If they are symmetrically polar, then no dipole appears, and there cannot be any dipole - dipole bonds. For example (draw the diagrams to see why):

CH3OCH3 forms a dipole.
C2F4 has no dipole as both ends of the molecule are polar.
CCl2H2 forms a dipole (think of the tetrahedral structure).
C2H5Cl forms a dipole as Cl is on one end.
CF4 does not form a dipole.
CH3CH=CH2 has no polar bonds and so no dipole.

Van der Waals forces

Where there is no stronger inter-molecular bonding, Van der Waals forces are important. They are always present between adjacent molecules, but they are usually rendered insignificant by stronger bonds.

The forces are due to vibrations of and fluctuations in the electron fields round the atoms in molecules. This causes temporary dipoles to be created on atoms, which can cause induced dipoles in neighbouring molecules. This results in weak attractive forces between molecules.

The strength of this type of bond depends on the number of electrons in the atom, and therefore the atomic number. This would be of the order of a few joules per mole.

Examples of places where Van der Waals forces are important.

  1. Boiling points of group 0/8 - increase down the group.
  2. The trend in the boiling points of the group IV hydrides (see graph). These increase up the group.
  3. Melting and boiling points of the alkanes - These increase as the chains of carbon atoms get longer. This is because there are more van der Waals forces between each molecule, since there can be, in theory, one bond between each pair of corresponding pair of molecules on two atoms. Therefore, each molecule has more Van der Waals forces attached to it, increasing the attraction.

Effects of Intermolecular Bonding - Liquids

Substances in the liquid state must be held together by forces, otherwise they would be gasses. These bonds between particles are the intermolecular bonds discussed previously.

This explains the trends in the boiling points of the noble gases - these rise as the gas atoms get larger, due to the greater energy needed to make the heavier atoms vibrate. The fact that it is possible to get noble gasses in the liquid form at all is due to the weak Van der Waals forces between the molecules.

Find a graph of the boiling points of the group 4,5,6 and 7 hydrides (or even better, produce one yourself). This shows a general trend of a rising boiling point, as the period rises, and therefore as the molecules get larger. This again, is caused by Van der Waals Forces between the molecules. However, the boiling points of the 1st period molecules in groups 5, 6 and 7 do not fit with the trend. This is due to hydrogen bonding between the molecules increasing the boiling points.

Structure, Bonding and Properties.

Sodium Chloride (NaCl, Common salt) - A giant ionic structure.

Sodium chloride is an ionic substance - it is made up of ions (atoms that have lost or gained electrons). To gain a stable outer electron shell of 8 electrons, sodium gives up an electron, becoming a 1+ ion. Chlorine then gains this electron, completing its outer shell and becoming a 1- ion. Although these two ions are stable, they are now electrostatically charged, and therefore are attracted to each other. They move together until they touch each other, forming a lattice. The sodium ion is now one shell smaller than the chlorine atom, and so the lattice that they form, although still cubic, looks slightly odd, with the big chlorine atoms effectively sandwiching the smaller sodium ions together.

The electrostatic forces that hold the Na+ and Cl- ions together are very strong, and therefore the ions are held together very strongly, and are very difficult to pull apart. This leads to very high melting and boiling points (777C and 1389C respectively) for these giant ionic structures, as compared with the smaller molecular structures. It also leads to the crystals formed being very hard, as it is difficult to break the ions away from their neighbours. All the ions are bonded, there are no weak points in the structure of chlorine. However, the crystals are also very brittle, as the ions are held very rigidly, without the ability to move. If, however, they are forced to move, they can split down a plane, when this happens, cations are brought next to cations, and visa versa, and they repel each other, leading to the crystal breaking.

When solid, NaCl does not conduct electricity, as there are no ions or electrons free to move. However, when NaCl is in a liquid form (either molten or dissolved in water), then the ions can move, and are able to carry an electric current. This is called electrolysis. NaCl can dissolve in water because of the very strong polarity of the water molecules, which are able to drag the ions out of the crystal form.

Diamond - A giant covalent structure.

Diamond is made up purely of carbon atoms. Each one is bonded to four other carbon atom by covalent bonds. This leads to a perfect tetrahedral structure being formed, making the diamond very pure and uncoloured. This structure, with many three-dimensional bonds from each atom, makes diamond very hard, the hardest naturally occurring substance, as again, with all the atoms all fully bonded, there are no weak points. However, again, for the same reason, it is also very brittle, as the bonds hold the atoms of carbon very strongly in place, and if they are pushed to one side, the bonds break and the crystal fractures. However, it fractures along predetermined planes, which means that it can be cut to shape very accurately. The many bonds that have to be broken mean that it takes a lot of energy to separate out the atoms - although diamond sublimes, it does so at a very high temperature - about 4800C. They also make diamond very chemically unreactive. As there are no free electrons or ions within its structure, it cannot conduct electricity at all.

Graphite - A weird and wonderful giant structure (but also giant covalent).

Graphite is also totally made up of carbon atoms. However, this time they arrange themselves into a structure where there are only 3 covalent bonds per carbon atom. This leads to a hexagonal structure, with the hexagons being arranged in a flat sheet. The second set of pi electrons for each atom, however, are forced out of the atoms and delocalise to above and below the rings of atoms, where they are free to move as they like. The sheets of atoms are arranged parallel to each other, with only the weak Van der Waals forces keeping them together. The delocalised electrons repel each other, keeping the sheets at a fair distance, and allowing them to slide over one another. This makes the graphite into a soft, greasy solid, which can be used as a lubricant. These delocalised electrons also allow it to conduct both heat and electricity, as they can move with the flow of an electric current, and also conduct vibrations along themselves as they bump into one another. However, even though there are only weak bonds between the sheets, the covalent bonds that hold the sheets together are very strong indeed, and it takes a lot of energy to break them. Graphite, like diamond, sublimes, and does so again at a very high temperature (4800C)

The giant metallic structure (applies to all metals, to some degree)

Metals are made up of a lattice of tightly packed cations, which are surrounded by a sea of delocalised electrons. The sea of electrons "glue" the ions together, as they are all strongly attracted to the ions. This creates, due to the electrostatic attraction between the cations and the electrons, a very hard substance, as the bonds are very strong. It also leads to high melting and boiling points, as the ions take a lot of energy to separate. In the same way as in graphite, the delocalised electrons give the metals their characteristic properties. They mean that the atoms can slide over each other, making metals malleable and ductile. They move in a constant direction when a potential difference is applied to them, allowing an electric current to flow. They move rapidly from areas of high energy to areas of low energy, thus conducting heat. They form a more even surface, but a very dense one, thus reflecting light and giving metals their characteristic lustre.

The properties, however, vary between the groups of metals. Group 1 metals, having only one electron in their outer shell, have less delocalised electrons per size of atom, especially lower down in the group. This means that there are less electrons to bind them together, and that therefore the atoms are less strongly bonded together. This leads to the metals being softer, and having lower melting and boiling points. This trend also tends to apply down the groups of all metals.

Solubility of Ionic Compounds

When an ionic compound dissolves in water, there are two processes at work. Firstly, the ionic compound must be split up into its two sets of ions, the cation and the anion. These then are attached to the poles of one or more hydrogen ion. The splitting up of the lattice is endothermic (takes in energy) wheras the solvation of the ions produced (forming dipole-dipole bonds with the water molecules) is exothermic (gives out energy).

It can be seen that, therefore, the energy change when an ionic compound dissolves in water is given by the following formula (Hess Law):

From this formula, it is possible to work out whether the reaction will be exothermic or endothermic. The more exothermic the reaction, the more soluble the substance, and vice versa.

However hydration energies are determined by the same factor that determines lattice energies - namely size of the ions. Smaller ions lead to the ions coming closer to each other or the water molecules, and therefore more energy is released by the process. This means that for this substance, both the lattice energy and the hydration energies will be high. Because of this, it is difficult to predict whether the substance will be soluble or not.

Shapes of Molecules
(Electron Pair Repulsion theory)

(A molecule is a small collection of covalently bonded atoms.) The following three rules can be used to predict the shape of molecules:

  1. Pairs of electrons around an atom repel each other and try to spread as far apart as possible.
  2. A lone pair is, on average, closer to the nucleus than a bonding pair. Therefore it exerts a greater repulsion. In order of increasing repulsion:
    Lone Pair - Lone pair
    Lone Pair - Bonding Pair
    Bonding Pair - Bonding Pair
  3. A double bond repels slightly more than a single bond.

Examples:

2 bonds - Beryllium chloride

Cl - Be - Cl
Linear arrangement; bond angle = 180

3 Bonds - Boron Trichloride


Trigonal Planar - Bond angle = 120

4 Bonds - Methane THIS IS THE MOST COMMON


Tetrahedral - Bond angle = 109 28' ( = 109.5)

5 Bonds - Phosphorus (V) Chloride(Gaseous)


Trigonal Bipyramid - Bond angle = 90 / 120

6 Bonds - Sulphur (VI) Chloride SECOND MOST COMMON


Octahedral - All angles 90

Ammonia (NH3)


Based on a tetrahedral arrangement, but has a lone pair, making the shape pyramidial.. Bond angle = 107.0 This is smaller than the tetrahedron angle, due to the greater repulsion of the lone pair.

Water (2 bonds and 2 lone pairs)


Water is based on tetrahedral angles but the extra repulsion due to the lone pairs reduces the bond angles to 104.

Ethene


The double bond repels slightly more than the single bonds, narrowing the outer bond angles.