Chapter 3: Chemical BondingLong Questions with Answers

Ionic bond

This chemical bond is formed as a result of the tendency of atoms to lose or gain electron or electrons to acquire the electronic configuration of the nearest noble gas, because this is a more stable electronic structure.

Formation of sodium chloride

Sodium chloride is formed when the elements sodium and chlorine react chemically. Their electronic configurations are Na 2, 8, 1 and Cl 2, 8, 7. An electron from the outermost shell of the sodium atom is transferred to the outermost shell of the chlorine atom, and in doing so both these atoms acquire the electronic configurations of their nearest noble gases.

The electrons which take part in a chemical reaction come only from the outermost shells. Sodium chloride contains the positively charged sodium ions (Na⁺) and negatively charged chloride ions (Cl⁻). These oppositely charged ions are held together by the electrostatic force of attraction. The chemical bond thus formed is called an ionic or electrovalent bond. Calcium, similarly, loses two electrons to form CaCl₂. An ionic bond is therefore a bond formed by the complete transference of electron or electrons from one atom to another atom.

Covalent bond

Some atoms decrease their energy by mutually sharing their electrons. When two atoms approach each other, the electrons of one atom come under the attractive influence of the nucleus of the other atom; this new force of attraction lowers the energy. The electrons and the nucleus of one atom also repel those of the other atom, which increases the energy. The two atoms bring themselves to such a distance that the attractive forces dominate the repulsive forces. The total energy at this distance is minimum and thus a stable molecule is formed.

A covalent bond is therefore a bond formed by the mutual sharing of an electron pair provided by the bonded atoms. This is called a single covalent bond. Atoms may share two electrons each to form a double covalent bond, or three each to form a triple covalent bond.

Sodium chloride contains Na⁺ and Cl⁻ ions held together by the electrostatic force of attraction. These ions then surround each other three-dimensionally to form a crystal lattice.

The NaCl lattice

Properties that follow

In ionic compounds the ions are held together by strong electrostatic forces in the form of a crystal lattice. Since the ions are rigid, such compounds exist as very stable solids with significantly high melting points (NaCl melts at 801 °C). If an external force is applied on the crystal lattice, it breaks easily: ionic solids are highly brittle.

Nature of the metallic bond

Usually metals have low values of ionisation energy. Their atoms can therefore lose their outer electrons easily. In sodium metal each sodium atom is surrounded by eight other sodium atoms. The outer electrons of these atoms move freely between the vacant spaces. No electron remains attached to any particular nucleus; instead all the electrons at the same time get attached to all the nuclei. When all the atoms attract all the electrons collectively, they are bound together. A metal appears to have a sea of electrons in which all the nuclei are submerged.

1. High melting and boiling points

The strength of a metallic bond depends on the number of positive charges on the ions and the number of mobile electrons set free by each atom. Each sodium atom sets free only one electron, so its metallic bond is not very strong; each magnesium atom releases two electrons, so magnesium melts at a higher temperature (650 °C against 98 °C).

2. Conduction of heat and electricity

The presence of freely moving electrons in metals makes them good conductors of heat and electricity.

3. Hard and heavy

In metals the atoms are strongly held and arranged in the form of rows one above the other. This arrangement makes them hard and heavy.

4. Malleable and ductile

When pressure is applied on metals, the upper rows of atoms slip past the lower rows. As a result their shapes are changed. Metals can therefore be easily drawn into wires and sheets.

5. Lustre

Metals usually show metallic lustre: the free electrons absorb and re-emit the light falling on the surface, so metals shine.

Properties of ionic and covalent compounds are adequately explained on the basis of the type of attractive forces present between them.

Structure of graphite

Graphite consists of a layered structure, made of hexagonal rings of carbon. In a layer each carbon atom forms covalent bonds with only three other carbon atoms, using three of its four outer electrons.

Mobile electrons

These layers in graphite have mobile electrons in between them. The fourth outer electron of each carbon atom is not fixed in a bond; these electrons are delocalised and free to move along the layers, like the electrons in a metal.

Conduction

When a voltage is applied, the mobile electrons drift through the layers and carry the current. Graphite is therefore a good conductor of electricity and is also used as an electrode, in battery cells and in electric arc furnaces to produce steel.

In metals, the atoms are strongly held and arranged in the form of rows one above the other. This arrangement makes them hard and heavy.

Hard

Every positive ion is attracted by the whole sea of mobile electrons. In sodium each atom is surrounded by eight others; these strong, all-round attractions hold the rows firmly, so most metals are hard.

Heavy

Because the atoms are packed so closely, many fairly heavy atoms fit into a small volume: metals have high densities, such as iron 7.9 g cm⁻³ and copper 8.9 g cm⁻³.

Exceptions

Sodium and potassium set free only one electron each, so their metallic bonding is weak: they are soft enough to cut with a knife and light enough to float on water (densities 0.97 and 0.86 g cm⁻³).

Why AlCl₃ joins up

A coordinate covalent bond is formed when a species has an electron pair to donate to another species. In AlCl₃ aluminium shares its three outer electrons with three chlorine atoms, so it has only six electrons in its outer shell: like boron in BF₃, it has a partially empty outer shell and can accept an electron pair. Each Cl atom has three lone pairs to donate.

The bridge

A giant covalent structure

Compounds may exist as a network arrangement of their atoms, for example covalent compounds like sand (SiO₂). Sand is mostly silicon dioxide (silica, as quartz). It is not made of small SiO₂ molecules: it is one giant network of atoms joined by covalent bonds, like diamond.

Properties that follow

Why atoms form bonds

Atoms have a tendency to decrease their energy. They can do this by combining with other atoms. It is a natural phenomenon because it increases the stability of atoms.

Clue from the noble gases

The early chemists found the answer when the noble gases He, Ne, Ar, Kr and Xe were discovered. Helium has two electrons in its outermost shell while all other noble gases have eight. Neither their atoms combine with themselves nor with other atoms. The probable reason for this lack of reactivity was their stability.

Duplet and octet rule

Having two electrons (for hydrogen and helium, which have only the first shell) or eight electrons in the outermost shell means stability and hence unreactivity as well. This principle was named the duplet or octet rule. Atoms form bonds because they would like to lower their energy by completing their duplet or octet.

Examples

Definition

Coordinate covalent bond is a type of covalent bond in which the shared electron pair is donated by one atom only. This bond is formed when a species has an electron pair to donate to another species. The species which donates the electron pair is called a donor, while that which accepts it is called an acceptor. An arrowhead (→) pointing towards the acceptor represents this type of bond.

1. Hydroxonium ion (H₃O⁺)

Acids provide protons (H⁺) when dissolved in water. This proton has an empty outer shell and can accept one of the two pairs of electrons present on the oxygen atom in the water molecule. As a result a hydroxonium ion is formed. The positive charge covers the whole of the hydroxonium ion, and all three bonds of oxygen behave exactly alike.

2. Ammonia and boron trifluoride

During the reaction, an electron pair from nitrogen of ammonia fills the partially empty outer shell of boron present in boron trifluoride: H₃N→BF₃.

3. Ammonium chloride

A coordinate covalent bond in ammonium chloride links nitrogen of ammonia and the proton. The positive charge is spread all over the ammonium ion. All the four bonds between nitrogen and hydrogen behave exactly alike.

Covalent vs coordinate covalent bond

This proves that the difference between a covalent bond and a coordinate covalent bond lies in the way they are formed. Once such bonds are formed, there does not remain any difference. When a coordinate covalent bond breaks, both the bonded electrons go to the donor species.

Intermolecular forces

The forces of attraction which are present between the molecules of elements and compounds are named intermolecular forces of attraction. These forces are generally very weak compared with the bonding forces between atoms. Among the three states of matter, they are weakest among the molecules of gases and strongest among the molecules of solids. The melting and boiling points of substances depend on their strength: the stronger the forces, the higher the boiling point.

1. Dipole-dipole forces

These attractive forces are present between the molecules of a polar compound like HCl. Hydrogen and chlorine attract the shared pair with different force (electronegativity). Since chlorine is more electronegative, it attracts the shared pair more strongly and the bond becomes polar: H(δ+)–Cl(δ−). Due to these partial charges the molecules of HCl start attracting each other. These forces are called dipole-dipole forces. Compounds with these forces show relatively higher melting and boiling points.

2. Hydrogen bonding

Hydrogen bonding is a special case of dipole-dipole attractive forces. When hydrogen is covalently bonded to highly electronegative elements like F, O or N, the large difference of electronegativity makes the covalent bond highly polar. As a result strong dipole-dipole attractions are observed among the molecules. In H₂O the O–H bonds are highly polar, so strong attractive forces develop between water molecules. This attractive force is called hydrogen bonding.

The strength of the hydrogen bonds causes water to have relatively higher melting and boiling points compared with H₂S and NH₃.

Non-polar covalent bond

When a covalent bond is formed between two identical atoms, the shared pair of electrons lies at the middle of the two bonded atoms. This gives rise to a pure covalent bond which is essentially non-polar in character. Covalent bonds in hydrogen and chlorine molecules are non-polar.

Polar covalent bond

When a covalent bond is formed between two non-identical atoms, the shared pair of electrons bends towards the atom which has more electronegativity. As a result a partial negative charge is created on the more electronegative atom and an equal partial positive charge on the less electronegative atom. A polar covalent bond is thus formed.

Examples

Lattice and melting point

In ionic compounds the oppositely charged ions are held together by the strong electrostatic force of attraction in the form of a crystal lattice. Since the ions are rigid, such compounds exist as very stable solids with significantly high melting points. Ions surround each other from all sides, so ionic bonds are non-directional.

Brittleness

If an external force is applied on the crystal lattice, it breaks easily: ionic solids are highly brittle.

Conduction

In the solid form ionic compounds do not conduct electricity because the ions are tightly held and cannot move. In the molten state the ions get free and start conducting electricity. In aqueous solution they also conduct, because the free ions move towards their respective electrodes.

Solubility and hydration

Ionic solids are generally soluble in water. Water not only breaks the electrostatic force of attraction but also hydrates the resulting free ions. In the process of hydration, water molecules surround and interact with ions.

Reactions in solution

Ionic compounds generally react in aqueous solution: the positive ions of one compound may react with the negative ions of the other to form a new compound. A white precipitate of silver chloride forms from NaCl and AgNO₃ solutions.

The mutually shared electrons may be shown by a dot or a cross. After mutually sharing their electrons, the bonded atoms acquire the electronic configurations of the nearest noble gas.