R EQUERIMIENTOS F UNCIONALES

A

X

_B

_Y

Highly ionic compound

Large cationic size

Small anionic size

Highly covalent compound

small cationic size

large anionic size

• The covalency properties of a molecule is dependent on the cation and anion where they can be explained qualitatively via

• Polarisation power of cation

• Polarisability of anion

3.6.1.1 Polarisation Power of Cation

Polarisation Power of Cation – measure the ability of a cation to polarise the electron cloud of the anion.

2 factors determining the polarisation power of cation

Charge of cation Size of cation

⇒ Greater the charge of ion, higher the

effective nuclear charge of cation, hence it will be able to attract the neighboring

electron density of anion. This will caused the polarization power of cation increase, hence increase the covalent characteristic of cation.

⇒ Smaller the size of cation, closer the

neighboring anion to the nucleus of cation, hence easier for the cation to polarise the anion and result an increment in the

polarization power of cation, and increase the covalent characteristic of cation.

♦ Both factors can be explained in another term called as charge density where Charge Density = Charge / Ionic Radius

♦ From the equation above, Charge Density will have a greater value, provided that cation has a high charge and small cationic radius.

♦ Greater the charge density, higher the polarization power, greater the covalent characteristic of the cation.

3.6.1.2 Polarisability of Anion

Polarisability of an anion ~ ability of the anion to allow the electron density to be polarised by cation.

2 factors determining the polarisability of an anion

Unlike cation, anion does not have a term that combined both factors of charge and ionic radius. However, information of polarisability of anion enable the prediction of the covalent characteristic of a molecule, since in order to form a covalent bond, it depend on both polarisation power of

cation and polarisability of the anion

Charge of anion Size of anion

⇒ Greater the charge of anion, lower the

effective nuclear charge of anion. This will weakened the electrostatic attraction forces between nucleus and the outermost electron in anion, and increase the polarisability of the anion, hence increase the covalent

characteristic of anion

⇒ Larger the size of anion, further the outermost electron from the nucleus of the anion, easier for the cation to

polarise the anion, and cause the

polarisability to increase, hence increase the covalent characteristic of anion.

3.6.2 Prediction of Chemical Bond :Fajans’ Rule

In 1923, Kazimierz Fajans formulated an easy guidance to predict whether a chemical bond will be covalent or ionic, and depend on the charge on the cation and the relative sizes of the cation and anion. They can be summarized in the following table

Based on these guidance, the bonding of a few compounds shall be discussed to understand the application of Fajans’ Rule in the chemical bonding

Ionic compound Low positive charge Large cation Small anion

Covalent compound High positive charge Small cation Large anion

Lithium halide (LiX)

Lithium ion, Li⁺ (1s²) has a small size due to only 1 shell present in its ion. But since it has a low charge, so its charge density is not too high. That is why, all lithium halide are ionic compound. The

covalency of lithium halide varies from a highly ioniccharacteristic to highly covalency, depending on the polarisability of the anion next to Li⁺

When a group of halide, F^– ; Cl^–; Br^–; I^– is put close to Li⁺, the

covalency of lithium halide increase when going down to Group 17 halide. LiF is highly ionic, since the fluoride ion has small ionic size and low charge, hence has low polarisability. Ionic size

increase with the increasing shell when going down to Group 17 halide, hence increase the polarisability, which allowed lithium ion to polarise the anion’s electron density, hence increase the

covalency

Li

F

Br

Cl

Aluminium halide (AlX₃) and aluminium oxide (Al₂O₃)

Aluminium ion (Al³⁺) has high charge density, due to its high charge unit and its small ionic radius. So, depending on the anion, aluminium has a high tendency to form covalent compound. For example, when going down to Group 17 halide, aluminium fluoride (AlF₃) forms ionic compound (since F- has a low polarisability), while aluminium trichloride (AlCl₃), aluminium tribromide (AlBr₃) and aluminium iodide (AlI₃) form covalent compound (since chloride, bromide and iodide have high polarisability). This explained why aluminium fluoride has a high melting point (1040⁰C), while aluminium trichloride and tribromide are 192⁰C and 78⁰C respectively.

As for aluminium oxide (Al₂O₃), it is an ionic compound with high covalent characteristic, as aluminium ion has high covalent characteristic due to its high charge density. This explained the high melting point of Al₂O₃ (2050⁰C) yet it is insoluble in water. It also explained the amphoteric properties of aluminium oxide where aluminium oxide can act as an acid (covalent characteristic), as well as a base (ionic characteristic).

Metallic Bonding

The properties of metals cannot be explained in terms of the ionic / covalent bond. In ionic / covalent compound, electron are not

free to move under the influence of applied potential (charge) difference. Therefore, ionic solid and covalent compound are insulator.

In metal, electron are delocalised and metal atoms are effectively ionised.

Metallic bond ~ electrostatic attraction between the positively charged metal ion and the electron delocalised.

Because of this, electron now can freely move from cathode to anode when a metal is subjected to an electrical potential. The mobile electron can also conduct heat by carrying the kinetic energy from a hot part of the metal to a cold part. This electron delocalised can also use to explain the electrical and thermal conductivities of metal

The Band Theory : Overlapping of Orbital

The number of molecular orbitals produced is equal to the number of atomic orbitals that overlap.

In a metal, the number of atomic orbitals that overlap is very large.

Thus the number of molecular orbital produced is also very large.

The energy separations between these metal orbitals are extremely small. So, we may regard the orbital as merging

together to form a continuous band of allowed energy state. This collection of very closed molecular orbital energy levels is called an energy band. This theory for metal is called band theory

Electrical Conductors

Molecular orbital model == 2 group of energy level.

Lower energy level – valence band → form from overlap of outer most orbital containing valence electron of each atom.

Higher energy level – conduction band → energy level filled with mobile electron

But there are some case where valence band can also serve as conduction band (caused by the movement of delocalised

molecular orbital)

Electrical conductivities decrease when temperature increase – vibration of the lattice of ion impedes the free movement of

electron in conduction band.

conduction band valence band

Insulator

Difference between conductors, semi-conductors, and insulator depend on the energy gap between the 2 bands.

Conductor – 2 bands overlaps so conduction band always partly filled.

Insulator – gap between the band is large and no electron exist in the conduction band. E.g. insulator – diamond

When 2s and 2p orbital of C is combine to form 2 energy bands, valence band is filled with electron.

In insulator, the energy gap between the band is large. Under

normal condition, few electrons in valence band can jump across to conduction band. If electron cannot reach conduction band

across the gaps, the electrical conduction cannot take place.

Semiconductor

There’s still energy gaps between 2 bands in semiconductor, but it is smaller than insulator.

In semiconductor, some electrons have sufficient energy to jump across the energy gaps and electron can move freely in

conduction band thus enable electrical conduction.

Still, the electrical activity is not as good as metal (conductor) Increasing temperature can help to improve the conductivity because electron gain thermal energy and are able to reach conduction band.

It can also improve its effectiveness by adding small amount of substance. This adding is what we called doping. It can help to increase electrons to fill in valence band.

Example of doping is Si dope P (n-type). Si dope Ge (p-type)

Depend on the needs, this process can help to create the various type of semiconductor in electronic characteristic.

7.1 Van der Waals forces

Van Der Waals forces are the intermolecular forces formed between covalently bond molecules which exist as simple molecules.

There are 2 types of Van Der Waals forces namely

♥ Permanent Dipole – Permanent dipole forces

♥ Temporary dipole – induced dipole forces

7.1.1 Dipole-dipole attraction forces

1. Polar molecule possessed dipole moment. Each of the polar molecules have an overall magnitude. For example in hydrogen chloride

H –––– Cl

δ+ δ–

2. The dipole inside polar molecules is permanent and the forces

between the molecule form as the positive end of dipole will attract to the negative end of another molecule’s dipole.

3. This kind of forced are called permanent dipole-dipole forces.

4. The strength of the attraction depends on two factors : dipole moment and relative molecular mass

5. Higher the dipole moment – the more polar the molecule – stronger the Van Der Waals forces

6. Comparisons were made between 4 molecules that have nearly equaled of molecular mass, but with different dipole moment

7. Methyl cyanide exhibit the highest boiling point among the 3 molecules as it has the highest dipole moment among these

molecules, which makes the attraction between the dipole-dipole attraction become stronger, and required a higher temperature to break the attraction forces among CH₃CN---CH₃CN.

Compounds RMM DM Boiling point (°C)

Propane , CH₃CH₂CH₃ 44 0.1 - 18.0

Methyl methoxide, CH₃–O–CH₃ 44 1.3 4.0

Chloromethane 50.5 1.9 6.0

Methyl cyanide, CH₃CN 41 3.9 56.0

8. Another factor which influence the strength of permanent dipole-dipole forces, are the factor of relative molecular mass.

9. Higher the mass, stronger the forces of attraction ( Van Der Waals forces ), higher the boiling point or melting point of the substance

RMM Melting

point (°C) Boiling point (°C) Hydrogen chloride, H – Cl 36.5 - 114 - 85 Hydrogen bromide, H – Br 81.0 - 87 - 66

Hydrogen iodide, H – I 128 - 51 - 35

7.1.2 Temporary dipole – induce dipole forces

Non-polar molecules have a dipole moment = 0. Basically, they won’t have any attraction between the molecules as there are no significant poles with charge in the molecule, so how they interact ??!!!

For non-polar molecules, they may have a chance to form