Structure and Bonding notes for high school

Bonding.

A chemical bond is a strong force of attraction holding molecules together in an atom, a molecule or a crystal.

Typically, chemical bonds have much higher energies and are distinguishes from the much weaker forces between molecules.eg van der Waal’s forces.

The fundamental types of bonds include;

• Ionic/electrovalent bond.
• Covalent bond.
• Metallic bond.

Ionic Bond.

They are formed by transfer of electrons. It occurs usually between a metal atom and a non-metal atom.

The metal atom transfers its valence electrons to the non-metal atom resulting into the formation of a metal cation and a non-metal anion. The two ions are then (attracted) brought together by electrostatic attraction which holds them together in an ionic crystal.

Properties of ionic structures/compounds.

• They are commonly soluble in water or other ionising solvents and are not soluble in organic solvents.
• They have high melting and boiling points owing to the strong attraction between the oppositely charged ions.
• They are often hard crystalline solids at room temperature because of inter ionic forces (between) within an ionic crystal are usually very strong.
• They are good electrolytes in molten and aqueous state.
• Individual molecules of ionic compounds do not exist because these compounds are made up of an interlocking structure of ions. The formula of an ionic compound simply shows the relative numbers of each ion present.

Covalent Bond.

These bonds are formed by sharing of electrons between atoms. The two bonded atoms contribute equally to the shared electrons.

A typical covalent bond is formed when there is negligible or small electronegativity difference between the bonded atoms.

A particular type of covalent bond is one in which the shared pair of electrons is contributed by only one atom. This type of covalent bond is called a coordinate bond (semi-polar or dative).

A condition for coordinate bond formation is that the atom donating the shared pair of electrons is electron rich/must be electron rich and the one accepting to share the pair must be electron deficient.ie. Must be bonded to a highly electronegative atom like fluorine.eg.

• The bond formed between the ammonia molecule and boron trichloride or trifluoride is a coordinate bond because the electron deficient boron atom accepts to share the lone pair of electrons in the nitrogen atom in the ammonia molecule.

Equation.

• Addition of a proton to an ammonia molecule to form an ammonium ion.

Equation.

A dative bond is weaker than a normal covalent bond.

Properties of covalent bonds.

• Covalent compounds are gases, liquids or soft solids at room temperature because weak intermolecular forces.
• They are commonly soluble in the organic solvents but not in water, except those containing highly electronegative atoms.eg. Nitrogen, oxygen, fluorine or chlorine bonded to hydrogen atoms.
• They have low melting and boiling points.
• They are made up of individual covalent molecules with weak intermolecular forces.

Bond Polarisation.

The bonding inmost compound is intermediate between purely ionic.eg. In caesium chloride or caesium fluoride and purely covalent.eg. In chlorine.

An ionic bond can have some covalent characters. This arises because of influence of an electric field of one ion on its neighbour.

A positive will attract the neighbouring anion and build up of charge cloud between the ions measures up to some degree.

The anion depends on its size and charge in bringing about bond polarization.eg. A large anion in which the outermost electrons are far away from the nucleus is easily deformed.ie. Iodide ion.

If the anion has a high negative charge, the electrons outnumber the protons to a greater extent than the anion with lower negative charge. For cations to form bond polarisation depends on charge and size. However a small cation with high charge is most effective in bringing about ion distortion.eg. Most beryllium compounds are covalent. Aluminium chloride is covalent.

Dipole; is the separation of two opposite charges by a small distance.

A stream of trichloromethane is deflected much more by a charged rod than a stream of tetrachloromethane.

Metallic Bonding.

In metallic structures, the atoms are closely packed in such a way that there are strong atomic forces (metallic bond) holding them together.

The outermost quantum shell electrons in each atom of the metal are delocalized around the atom and the whole metal structure.ie. They are free and mobile electrons. The attraction for these electrons by the nuclear charge brings about the atomic forces whereby each atom attracts the nearby electrons. It is very hard (high melting point) to overcome these metallic bonds.

NB: a metallic bond is different from a covalent bond in that it is non-directional. The more the number of electrons contributed to the pull of delocalised electrons by each atom, the stronger is the metallic bond.

Metallic bonds are distinguished from other types of bonds due to their high conductivity of electricity and heat.

Intermolecular Forces Of Attraction.

Dipole-dipole forces.

They occur between polar molecules. A dipole is formed by unequal sharing of electrons. A dipole-dipole force is caused by attraction of positive and negative poles of molecules to one another.

Van der Waal’s forces.

These are attractive forces between atoms or molecules. They are much weaker than covalent, ionic or metallic bonds. They are the forces responsible for the non-ideal behaviour of gases and for the lattice energy of molecular crystals.

There are three factors causing such forces;

• Dipole-dipole interaction.ie. Electrostatic attraction between two molecules with permanent dipole moments.
• Dipole-induced dipole interactions; dipole of one molecule polarizes a neighbouring molecule.
• Dispersion of forces; arising because of small instantaneous dipole in atoms.

Hydrogen Bond.

This is a dipole-dipole attraction between a hydrogen atom bonded to a highly electronegative element such as fluorine, oxygen or nitrogen of one molecule and electronegative element of a second molecule.

The large electron shift between hydrogen and a highly electronegative element brings about a large dipole moment in the molecule.

The dipole-dipole attraction accounts for;

The unexpected high solubility of some compounds containing oxygen, nitrogen, and fluorine in certain hydrogen containing solvents notably, water.eg. Ammonia, ethanol, methanol, methanoic acid, ethanoic acid, hydrogen chloride, dissolve in water through the formation of hydrogen bonds.

Diagram.

Lower members of alcohols, carboxylic acids, amines and some carboxylic compounds are soluble in water due to hydrogen bonding.

Hydrogen bonding accounts for the structure adopted by some molecules.eg.

• Ice has an open structure; in the ice, hydrogen atoms are held together by hydrogen bonds, specifically each oxygen atom is bonded to two hydrogen atoms by a normal covalent bond and two others by long hydrogen bonds to ice a tetrahedral configuration. This gives ice an extremely open structure such that for a given mass of water, mass: volume ratio in ice (density) is smaller than the mass: volume ratio in water. This is because, during melting, heat is used to progressively break hydrogen bonds in ice which results into the close packing of water molecules hence ice floats in water.

Diagram.

• Ethanoic acid in benzene forms dimers.

Diagram.

• Association of hydrogen fluoride molecules.

Diagram.

Hydrogen bonding accounts for the very high boiling points and molar heats of vaporisation of compounds containing nitrogen, oxygen, and fluorine.eg. The boiling points of hydrides of period 2, particularly water, hydrogen fluoride and ammonia are abnormally high as indicated below. In the plots of boiling points against period number for hydrides of group IV, V, VI and VII.

Graph.

The series H2Se, A5H3, HBr and GeH4 illustrates the expected trend in boiling point for compounds in which there is only van der waal’s forces of attraction.ie. boiling points increase as the molecular size increases.

Boiling points of H2O, HF, NH3 and CH4 are anomalously high implying that they have strong intermolecular forces of attraction specifically hydrogen bonds.

How important is hydrogen bonding to life.

• It has made water to be a liquid at room temperature instead of a gas.
• Has made ice less dense than water.
• The secondary structure; alpha-helix of a proton is due to hydrogen bonding and also the double helix in DNA molecule is due to this.