Question 18
Explain the terms :
(a) Lone pair of electrons
(b) Coordinate bond
Explain diagrammatically the lone pair effect of:
(a) The nitrogen atom of the ammonia molecule leading to the formation of ammonium ions [NH4]+
(b) The oxygen atom of the H2O molecule leading to formation of hydronium [H3O]+ and hydroxyl ions [OH]-
(a) Lone pair of electrons — They are a pair of electrons not shared with any other atom.
(b) Coordinate Bond — It is a type of covalency which involves one of the combining atoms contributing both of the shared electrons. i.e., a bond formed by a shared pair of electrons with both electrons coming from the same atom.
(a) Formation of ammonium ions [NH4]+
In ammonia, the 'N' atom contains one lone pair of electrons after completing it's octet. This lone pair is accepted by the hydrogen ion of water leading to the formation of a coordinate covalent bond:
(b) Formation of hydronium (H3O)+ and hydroxyl ions [OH]-
In water, the 'O' atom contains two lone pairs of electrons after completing it's octet. These lone pairs are accepted by the hydrogen ion leading to the formation of coordinate covalent bond:
Chapter Overview: Chemical Bonding
Chemical bonding explains how atoms combine to form molecules and compounds. Atoms bond to achieve a stable electronic configuration (octet or duplet). The three primary types of bonds are ionic (electrovalent), covalent, and coordinate (dative) bonds. Ionic bonds form by the transfer of electrons from a metal to a non-metal, creating oppositely charged ions held by electrostatic attraction. Covalent bonds form by the sharing of electron pairs between non-metal atoms. Coordinate bonds are a special type of covalent bond where both shared electrons come from one atom. The type of bond determines the physical properties of a substance: ionic compounds have high melting points and conduct electricity in molten/aqueous state, while covalent compounds generally have low melting points and are poor conductors. Students must draw electron dot (Lewis) structures, understand the octet rule and its exceptions, differentiate between polar and non-polar covalent bonds, and relate bonding type to physical properties. This chapter carries significant weightage and questions frequently involve drawing structures and explaining properties.
Key Definitions
| Term | Definition |
|---|---|
| Ionic Bond | Electrostatic force of attraction between oppositely charged ions formed by electron transfer |
| Covalent Bond | Bond formed by mutual sharing of electron pairs between two atoms |
| Coordinate Bond | Covalent bond where both shared electrons are donated by one atom (donor → acceptor) |
| Electrovalency | Number of electrons lost or gained by an atom to form an ion |
| Covalency | Number of electron pairs shared by an atom in covalent bond formation |
| Polar Covalent Bond | Covalent bond between atoms of different electronegativities; shared pair is displaced towards the more electronegative atom |
| Octet Rule | Atoms tend to gain, lose, or share electrons to achieve 8 electrons in the outermost shell |
| Lone Pair | A pair of valence electrons not involved in bonding |
Must-Know Concepts
- Na (2,8,1) transfers 1 electron to Cl (2,8,7) forming Na+Cl− (ionic bond)
- In H2O, oxygen shares 2 electron pairs with 2 hydrogen atoms (covalent, polar)
- In NH4+, the lone pair on nitrogen forms a coordinate bond with H+
- Single bond (1 shared pair), double bond (2 pairs, e.g. O2), triple bond (3 pairs, e.g. N2)
- Ionic compounds: crystalline solids, high MP/BP, conduct electricity when molten or in solution
- Covalent compounds: low MP/BP, generally insoluble in water, do not conduct electricity
- Electronegativity difference determines bond type: large difference = ionic, small = covalent
Ionic vs Covalent Compounds
| Property | Ionic Compounds | Covalent Compounds |
|---|---|---|
| Bond Formation | Electron transfer | Electron sharing |
| Melting Point | High | Low |
| Solubility | Soluble in water | Soluble in organic solvents |
| Conductivity | Conduct when molten/dissolved | Do not conduct |
| State at Room Temp | Crystalline solids | Gases, liquids, or soft solids |
Important Diagrams to Practice
- Electron dot diagrams for NaCl, MgO, CaCl2 (ionic)
- Electron dot structures for H2O, NH3, CH4, CO2, C2H4, N2 (covalent)
- Coordinate bond formation in NH4+ and H3O+
Common Mistakes
- Drawing electron dot structures without showing lone pairs
- Confusing coordinate bond with ordinary covalent bond (arrow shows direction in coordinate bond)
- Saying ionic compounds conduct electricity in solid state (they do not - ions are fixed)
- Forgetting to show the charge on ions in ionic bond diagrams
- Not recognising that HCl is a polar covalent molecule (not ionic despite containing a metal-like H)
Scoring Tips
- Always show complete electron transfer/sharing with arrows in diagrams
- When comparing ionic and covalent, use a table format for clarity and full marks
- For coordinate bonds, clearly identify the donor (has lone pair) and acceptor (has empty orbital)
- Practice drawing at least 10 electron dot structures till they become automatic
Frequently Asked Questions
Why do ionic compounds have high melting points?
Ionic compounds consist of a lattice of positive and negative ions held by strong electrostatic forces. A large amount of energy is needed to overcome these forces, resulting in high melting points.
Can a molecule have both ionic and covalent bonds?
Yes. For example, NaOH has an ionic bond between Na+ and OH−, while the O-H bond within the hydroxide ion is covalent.
What is the difference between polar and non-polar covalent bonds?
In a non-polar covalent bond, electrons are shared equally (e.g., H2, Cl2). In a polar covalent bond, electrons are shared unequally due to a difference in electronegativity (e.g., HCl, H2O).