Bonding And Solubility.notes for o level chemistry

Qualitative Consideration.

A solute will dissolve in a solvent if the new bonds between the solute and solvent particles are stronger than those between solvent and solvent particles and those between solute and solute particles. This is called entropy factor.

Entropy factor is the degree of disorderliness and randomness of a system.

Ionic Compounds.

These are soluble in compounds which are polar.eg. water because strong coordinate bonds can form between the ion and solvent molecules. This is true if the lattice energy is not too large.eg. for NaCl solubility can be illustrated as;

Diagram.

If lattice energy of the ionic compound is too large compared to the hydration energy then the ionic compound will not dissolve in water. The energy required to break one mole of an ionic compound into its constituent gaseous ions is called lattice energy. The heat energy liberated when each ion is hydrolysed by water molecules is called hydration energy

Molecular Substances.

These are soluble in polar solvents if they themselves are polar because they will be attracted by solvent.eg. lower members of alcohols, carboxylic acids, amines, carbonyl compounds, etc.

Molecular substances that are not polar.

These are soluble in non-polar solvents are non-polar because they will be attracted by van der waal’s forces.

Molecular substances having hydrogen bonding.

These are soluble in water with which hydrogen bonds can be formed.eg. lower members of alcohols, amines, etc.

NB: like dissolves like; this means a substance which has same intermolecular forces of attraction as the solvent will dissolve in it.

Shapes Of Molecules.

In covalent molecules, a bond must constitute a pair of electrons. How the electron pairs around the central atom are arranged in a molecule affects the shape of the molecule. The interaction of the outer most electrons of one atom with those of another atom brings about the overall shape of the molecule.

NB: paired electrons on the valence shell of an atom don’t get involved in covalent bonding. They are called (lone-pairs) non-bonding pairs. It is the unpaired that get involved. After a covalent bond has been formed, the shared pair of electrons is under the influence of two nuclei and it will be held further away from the central atom compared to the lone pairs which are under the influence of only one nucleus. Electron pairs, both bonding and non-bonding on the valence shell of an atom repel each other. Repulsion of electron pairs is in the order lone-lone> lone-bond> bond-bond.

Valence Shell Electron Repulsion Theory (VSEPR Theory)

It states that electron pairs around a central atom in a molecule do repel each other so as to attain a stable arrangement in space where there’ll be minimum repulsion between them and maximum angle of separation between the bond pairs.

According to the VSEPR theory, multiple bonds act as a single pair of electrons.

Molecules with one electron pair shared.

The shape of these molecules is linear because the electrons in one bond pair around the central atom repel each other as fully as possible, giving bond angle of 180 degrees. Eg. H₂; H-H, HCl; H-Cl, etc.

Molecules with two electron pairs all shared.

The shape of these molecules is linear because the two bonded pairs around the central atom repel each other as fully as possible giving bond angle of 180.eg. BeCl₂; CL-Be-Cl, HgCl₂; Cl-Hg-Cl, CO₂; , CS₂; .

Three electron pairs all shared.

As these electron pairs repel each other equally, the molecule will attain a trigonal planar shape with bond angle of 120 degrees. Eg.

• BCl₃ or BF₃
• B (CH₃) ₃
• SO₃
• CO₃^2-
• NO₃^-

Three electron pairs including one lone pair.

The molecule is V-shaped or bent with bond angle less than 120 degrees because of the increasing repulsion caused by the lone pair which brings the bond pairs together.eg. Tin (II) chloride (SnCl₂).

• SnCl₂
• SO₂
• O₃

Four electron pairs all shared.

The four electron pairs are distributed tetrahedrally around the central atom and these account for the tetrahedral shape of the molecule with a bond angle of 109 degrees 28 minutes as in methane. This shape is also shown by ammonium ion, sulphate ion, chromate ion, permanganate ion, phosphate, trichloromethane ion.

Four electron pairs including one lone pair.

The three bond pairs and one lone pair will orientate themselves around the central atom such that the resulting molecule XY₃ will be trigonal pyramidal in shape. The increased repulsion caused by the lone pair will lead to the bond angle being smaller than the tetrahedral angle of 109 degrees 28 minutes found in methane, ammonia and the rest of the hydrides of the group V elements have the formula XY₃. They all have this shape.

NB: the hydrides of group V elements show a general decrease in bond angle as the group is descended as indicated in the shapes below.

Diagrams.

This trend is due to the electronegativity decrease down the group of group V elements and hence nuclear attractions of the elements in the group.

Phosphorus has a lower electronegativity than nitrogen so the shared electrons (bond pairs) are closer to the nitrogen atom in ammonia than they are in the phosphorus atom in phosphine. The fact that the shared pairs of electrons are closer to each other and therefore they experience greater repulsion giving rise to bigger bond angle.

In SbH₃.ie. antimony hydride has much lower electronegativity and nuclear attraction and the shared pair of electrons are therefore further from the central atom and also from each other making them repel each other to a small extent and thus reducing the bond angle. The species PCl₃, SO₃^2- and CLO₃^-.

Four electron pairs including two lone pairs.

In ammonia, the presence of one lone pair increase the repulsion between bond pairs and the lone pair, such that the bond angle decreased from 109 degrees 28 minutes in methane to 107 degrees. In water, the molecule H₂O has repulsive effect of two lone pairs on two pairs and this reduces the bond angle from the expected 109 degrees 28 minutes in methane to 104 degrees, giving the molecule a bent or a V-shape.

• H₂O
• Cl₂O
• ClO₂^-

Five electron pairs.

The bond pairs will orientate themselves around a central atom such that the resulting molecule XY₅ will be a trigonal bipyramid.eg. in PCl₅.

Six electron pairs including one lone pair.

The five bond pairs and one lone pair are distributed tetrahedrally around the central atom. The shape of the molecule is a square pyramid.eg. in IF₅.

Six electron pairs.

The six pairs orientate themselves around a central atom such that the resulting molecule is XY₆. The shape of the molecule is octahedral.eg. SF₆ (sulphur hexafluoride).

Six electron pairs with two lone pairs.

The shape of the molecule is a square planar.eg. in ICl₄^-.

Types Of Solids Structures.

A solid is a substance which possesses a constant shape and volume at constant temperature.

Structure of solids are determined from diffraction patterns made by passing x-rays through the solid. The x-rays reveal;

• Arrangement of particles.ie. ions, atoms, molecules, etc. in a solid.
• Electron density in different parts of the crystal.

From this information, the shapes of the ions, molecules and structural formulae, and bond length are obtained.

There are four types of crystal structures based on the kind of structure particles that make up the crystal and the forces that bind them together. They include;

• Molecular crystal structures.
• Giant crystal.
• Giant covalent/network.
• Giant metallic.

NB; a giant structure refers to one that consists of an infinite assembly of particles while molecular is one that consists of discreet (countable) molecules.

A giant structure can be in 3 or 2 dimensions.

Space lattice/unit cell is the building body which repeats itself over and over to form the entire molecule or crystal. It is the basic unit of a crystal. There are two types of space lattice.ie. cubic and hexagonal lattice.

The cubic lattice consists of the following;

• A simple cube; particles are formed at the corners.
• Face centred cubic lattice; the particles are on each face of a cube.eg. NaCl structure.
• Body centred cubic lattice; one particle is at the centre of a cube.eg. Caesium chloride (CsCl).

Coordination Number.

This is the number of groups (ions, atoms, molecules) surrounding a particular ion or atom in a crystal or a complex ion.

Ionic Crystals.

They consist of positive and negative ions being held together in a crystal lattice by electrostatic attraction.ie. NaCl, KNO₃ barium oxide. The space lattice is a face centred cubic lattice.

Diagram.

Sodium chloride has a giant ionic structure which consists of £D infinite assembly of Na⁺ and Cl^-. The coordination number for each ion is 6. The forces of attraction holding the ions together are electrostatic forces. It has high melting point, the crystals are brittle and conduct electricity in molten/fixed state and aqueous state only.

Molecular crystals.

These are covalent compounds. Molecules occupy lattice position in crystals.ie. Discrete molecules occupy the lattice position. Intermolecular forces are weaker so they are soft with low melting points. Intermolecular forces are van der waal’s for crystals made up of non-polar molecules.eg. iodine, naphthalene, benzene, etc. molecular crystals do not conduct light in solid or liquid state except a few compounds.eg. Water which dissociates to a very slight extent.eg. Iodine molecules.

Giant covalent (network) crystals.

Eg. Diamonds.

In diamond, each carbon atom (Sp³ hybridized) is bonded tetrahedrally to four other carbon atoms. (4:4 coordination).

This symmetrical 3D structure accounts for its extremely hard physical nature. As all the outer electrons are used in forming identical bonds. The crystal is colourless, doesn’t conduct electricity and is a poor conductor of heat.

Diagram of tetrahedral arrangement of carbon atoms.

Graphite.

The carbon atoms may be considered to be Sp² hybridized and forma planar arrangement of fused hexagonal rings. The planes are held together by weak van der waal forces enabling them to be easily cleaved. The electrons unaffected by hybridisation lie at right angles to the plane of carbon atoms and form layers of delocalized electrons accounting for the good electrical conductivity of graphite.

Because these electrons are weakly held in the structure they absorb light from most visible parts of the visible spectrum, giving graphite its physical black appearance. The separation between the planes results in graphite being much softer than diamond and is widely used as a lubricant. Notice that in a perfect graphite crystal, conduction would only occur along the planes and resistance to shear at right angles to the planes would be very high.

Giant metallic crytals.eg. copper.

In copper, most cubic metallic crystals, face centred cubic lattice is the structure. It’s one of the compact arrangements of packing spheres of the same size. Coordination number is 12. Binding force in metallic crystals depends on the valency of the metal.eg. silver, copper, sodium, potassium and iron.

Effects of structure on physical properties.

• The differences in volatility.ie. boiling point and molar heat of vaporisation can be used as a guide to structural types. It indicates the amount of energy required to break down regular arrangement of particles in the crystals.
• General high melting point is associated with an infinite arrangement of atoms or ions.eg. oxides of period 3.

Oxide	Na2O	MgO	Al2O3	SiO3	P2O5	SO3	Cl2O
Melting point	1193	9075	2300	1728	563	30	-91

The high melting point of sodium oxide to aluminium shows the strength of ionic bonds in the crystal. The melting point of silicon (IV) oxide is higher than that of sodium oxide because of the strength of giant covalent structure.

The rest are molecular solutes. The differences are due to the size.ie. phosphorus (V) oxide is a larger molecule than sulphur (VI) oxide and Cl₂, so has higher van der waal forces hence high melting point.

Chlorides.

Chlorides	NaCl	MgCl2	AlCl3
Melting point	801	708	sublimes

There’s rapid decrease in ionic character and rapid increase in covalent character due to increase in electronegativity. This can be detected by;

• Conductivity measurement.

NaCl	MgCl2	AlCl3
134	29	0.000015

Low conductivity shows covalent properties.

• Hardness.

The resistance to penetration or distraction. It involves forcing apart the molecules of a solid and this is more difficult. The stronger the forces holding the molecules together.eg. Solid with infinite 3D atoms/ions joining covalent or ionic bonds are very hard.eg. Diamond, silicon (IV) oxide, aluminium oxide and sodium oxide because of giant structure. Molecular structures are held by van der waal’s forces.eg. candle wax, paraffin, wax, naphthalene are soft and greasy while molecular crystals held by hydrogen bonding.eg. ice, sucrose are brittle and hard.

• Conductivity in fused state.
• Solubility; It depends on strength of forces holding particles together, nature of particles and nature of solvent.

Delocalized Electrons.

• In metallic bonding; described as giant structure of metal ions held together by delocalized valency electrons. This accounts for high thermal electro conductivities and have high melting points except group I metals which have pure delocalized electrons.
• In graphite.
• In molecules.eg. benzene; here each atom has 3 delocalized bonds, 2 to a carbon atom and 1 to a hydrogen atom. This leaves a singly occupied p orbital.
• In benzene (C₆H₆), there are 6 delocalized valence electrons moving about the ring showing that electrons aren’t fixed.
• Others are nitrate ion, carbonate ion, sulphate ion have delocalized electrons in molecules. Special property of compounds containing delocalized bonds is that they are more stable than similar compounds which only contain localised bonds.

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