Comprehensive Notes on Chemical Bonding: Electrovalent, Covalent, and Coordinate Bonds

Electronic Configuration and the Stability of Atoms

• The Drive for Stability: All atoms strive to achieve stability, which is defined as possessing the electronic arrangement of an inert gas. This typically involves having an octet (eight electrons) in the outermost shell.

• Duplet vs. Octet:

  • Duplet: Helium ($He$) has only two electrons in its valence shell, known as a duplet configuration.

  • Octet: Other inert gases such as Neon ($Ne$), Argon ($Ar$), Krypton ($Kr$), Xenon ($Xe$), and Radon ($Rn$) possess eight electrons in their outermost shell.

• Reactivity Trends: Elements with complete outermost shells (noble gases) are found to be chemically unreactive. Atoms with incomplete shells react to redistribute their valence electrons via chemical bonding to attain a stable configuration.

• Table 2.1: Electronic Configurations of Inert Gases:

  • Helium ($He$): Atomic No. 2; K: 2; Valence Electrons: 2.

  • Neon ($Ne$): Atomic No. 10; K: 2, L: 8; Valence Electrons: 8.

  • Argon ($Ar$): Atomic No. 18; K: 2, L: 8, M: 8; Valence Electrons: 8.

  • Krypton ($Kr$): Atomic No. 36; K: 2, L: 8, M: 18, N: 8; Valence Electrons: 8.

  • Xenon ($Xe$): Atomic No. 54; K: 2, L: 8, M: 18, N: 18, O: 8; Valence Electrons: 8.

  • Radon ($Rn$): Atomic No. 86; K: 2, L: 8, M: 18, N: 32, O: 18, P: 8; Valence Electrons: 8.

Introduction to Chemical Bonding

• Chemical Bond Definition: A chemical bond is the force of attraction between any two atoms in a molecule that maintains stability.

• Cause of Combination: Elements combine to acquire the nearest noble gas configuration in their outermost orbit.

• Methods of Achieving Stability:

  1. Transfer of Electrons: One or more electrons move from one atom to another, forming an electrovalent (ionic) bond.

  2. Sharing of Electrons: Mutual sharing of one, two, or three pairs of electrons forms a covalent (molecular) bond.

  3. One-sided Sharing: Shared electron pairs are contributed by only one of the combining atoms, forming a coordinate (dative) bond.

Electrovalent (Ionic) Bonding

• Mechanism: Atoms of metallic elements (1, 2, or 3 valence electrons) lose electrons to non-metallic elements (5, 6, or 7 valence electrons).

• Ion Formation:

  • Cation: A metallic atom that loses electron(s) and becomes positively charged (electropositive elements). Example: $Na - e^{-} \rightarrow Na^{+}$.

  • Anion: A non-metallic atom that gains electron(s) and becomes negatively charged (electronegative elements). Example: $Cl + e^{-} \rightarrow Cl^{-}$.

  • Definition: An ion is a charged particle formed due to the gain or loss of one or more electrons by an atom.

• Definition of Electrovalent Bond: The chemical bond formed due to the electrostatic force of attraction between a cation and an anion.

• Electrovalency: The total number of electrons an atom loses or gains to form an electrovalent bond.

• Conditions for Formation:

  1. Low Ionisation Potential: Higher tendency to lose electrons and form cations.

  2. High Electron Affinity: Higher tendency to gain electrons and form anions.

  3. Large Electronegativity Difference: Facilitates the transfer of electrons; larger differences lead to more ionic character.

• Periodic Trends:

  • Group 1, 2, and 13 metals combine with Group 15, 16, and 17 non-metals.

  • Group 1 elements are most electropositive (metallic character increases down the group).

  • Group 17 elements are most electronegative; Fluorine is the most electronegative.

  • Caesium Fluoride ($CsF$) is the most ionic compound.

• Stability of Ionic Compounds: Despite repulsive forces between like-charged ions, the electrostatic force of attraction between opposite charges is far greater, providing stability.

Structures of Electrovalent Compounds

• Sodium Chloride ($NaCl$):

  • Sodium ($Na$): $2, 8, 1$. Loses $1e^{-}$ to become $Na^{+}$ ($2, 8$), resembling Neon.

  • Chlorine ($Cl$): $2, 8, 7$. Gains $1e^{-}$ to become $Cl^{-}$ ($2, 8, 8$), resembling Argon.

  • Table 2.2: Sodium Atom vs. Sodium Ion:

    • Sodium Atom ($Na$): Silvery white, poisonous, very active, incomplete shell, neutral, combined state existence.

    • Sodium Cation ($Na^{+}$): Colourless, non-poisonous, inactive, complete shell, positively charged, independent existence.

  • Table 2.3: Chlorine Atom vs. Chloride Ion:

    • Chlorine Atom ($Cl$): Yellowish green gas, poisonous, suffocating odour, very active, incomplete shell, neutral, not independent.

    • Chloride Anion ($Cl^{-}$): Colourless, non-poisonous, odourless, inactive, complete shell, negatively charged, independent.

• Magnesium Chloride ($MgCl_2$):

  • Magnesium ($2, 8, 2$) loses two electrons: $Mg - 2e^{-} \rightarrow Mg^{2+}$ ($2, 8$).

  • Two Chlorine atoms each accept one electron: $2[Cl + e^{-} \rightarrow Cl^{-}]$.

  • Ratio of $Mg^{2+}:Cl^{-}$ is $1:2$.

• Calcium Oxide ($CaO$):

  • Calcium ($2, 8, 8, 2$) loses two electrons: $Ca - 2e^{-} \rightarrow Ca^{2+}$ ($2, 8, 8$).

  • Oxygen ($2, 6$) gains two electrons: $O + 2e^{-} \rightarrow O^{2-}$ ($2,8$).

  • Formula is $CaO$ because one Oxygen atom accepts both electrons from one Calcium atom.

Redox Processes in Chemical Bonding

• Definitions:

  • Oxidation: Loss of electrons by an electropositive atom.

  • Reduction: Gain of electrons by an electronegative atom.

  • Redox Reaction: Oxidation and reduction occurring simultaneously.

• Example: Formation of $NaCl$:

  • Oxidation: $2Na \rightarrow 2Na^{+} + 2e^{-}$.

  • Reduction: $Cl_2 + 2e^{-} \rightarrow 2Cl^{-}$.

  • Net: $2Na + Cl_2 \rightarrow 2Na^{+} + 2Cl^{-}$.

• Reducing and Oxidising Agents:

  • Reducing Agent: A donor of electrons (e.g., Sodium).

  • Oxidising Agent: An acceptor of electrons (e.g., Chlorine).

• Example: Hydrogen and Copper(II) Oxide:

  • $CuO + H_2 \rightarrow Cu + H_{2}O$.

  • $Cu^{2+}$ is reduced to $Cu$ ($Cu^{2+} + 2e^{-} \rightarrow Cu$).

  • $H_2$ is oxidised to $H^{+}$ ($2H - 2e^{-} \rightarrow 2H^{+}$).

Covalent (Molecular) Bonding

• Definition: A bond formed by the mutual sharing of electron pairs between two atoms of non-metals.

• Types of Covalent Bonds:

  1. Single Covalent Bond: Sharing of one pair of electrons (e.g., $H-H$, $Cl-Cl$, $H-Cl$, $NH_3$, $CH_4$, $CCl_4$, $H_2O$).

  2. Double Covalent Bond: Sharing of two pairs of electrons. Represented as "=" (e.g., $O=O$ in $O_2$, Ethene $C_2H_4$).

  3. Triple Covalent Bond: Sharing of three pairs of electrons. Represented by three lines (e.g., $N \equiv N$ in $N_2$, Ethyne $C_2H_2$).

• Covalency: The number of electrons an atom contributes to shared pairs (Hydrogen: 1, Oxygen: 2, Nitrogen: 3, Carbon: 4).

• Conditions for Formation:

  • Both atoms should have 4 or more valence electrons (except $H$, $Be$, $B$, $Al$).

  • Both atoms need high electronegativity, high electron affinity, and high ionization energy.

  • Electronegativity difference should be zero or negligible.

• Non-Polar Covalent Compounds:

  • Equal distribution of shared electrons; symmetrical molecule with no charge separation.

  • Examples: $H_2$, $Cl_2$, $O_2$, $CH_4$, $CCl_4$.

• Polar Covalent Compounds:

  • Shared electrons are attracted more by one atom (the more electronegative one), leading to fractional charges ($\delta^+$ and $\delta^-$).

  • Known as "dipole molecules."

  • Examples: $HCl$ ($Cl$ energy = 3.0, $H$ energy = 2.1), $H_2O$, $NH_3$, $HF$.

Detailed Structure of Covalent Molecules

• Hydrogen ($H_2$): Two H atoms (1 electron each) share one pair to reach a duplet.

• Chlorine ($Cl_2$): Two Cl atoms (7 valence each) share one pair to reach an octet.

• Nitrogen ($N_2$): Two N atoms (5 valence each) share three pairs ($N \equiv N$) to reach an octet.

• Water ($H_2O$): Oxygen (6 valence) shares one pair with each of two Hydrogen atoms. Result: Polar molecule with two single bonds.

• Ammonia ($NH_3$): Nitrogen (5 valence) shares one pair with each of three Hydrogen atoms. Result: One lone pair remaining on Nitrogen.

• Carbon Tetrachloride ($CCl_4$): Carbon (4 valence) shares one pair with each of four Chlorine atoms (7 valence each).

• Methane ($CH_4$): Carbon (4 valence) shares one pair with each of four Hydrogen atoms.

Comparison of Electrovalent and Covalent Compounds

• Nature:

  • Electrovalent: Hard solids composed of ions. Strong electrostatic attraction.

  • Covalent: Gases, liquids, or soft solids composed of molecules. Weak intermolecular forces.

• Boiling/Melting Points:

  • Electrovalent: Non-volatile, high MP and BP (requires high energy to break strong ionic bonds).

  • Covalent: Volatile, low MP and BP (weak forces between molecules require less energy to break).

• Electrical Conductivity:

  • Electrovalent: Non-conductors as solids; good conductors in molten or aqueous states due to mobile ions.

  • Covalent: Non-conductors in any state (solid/molten/aq) due to absence of free ions (Note: Polar covalent molecules ionise in water).

• Solubility:

  • Electrovalent: Soluble in water (polar solvent), insoluble in organic solvents.

  • Covalent: Insoluble in water, soluble in organic (non-polar) solvents.

• Rate of Reaction:

  • Electrovalent: Rapid speed in aqueous solutions (fast ion combination).

  • Covalent: Slow speed (requires breaking old bonds and forming new ones).

• Dissociation/Ionisation:

  • Electrovalent: Dissociate into ions ($NaCl \rightarrow Na^{+} + Cl^{-}$).

  • Covalent: Do not dissociate (except polar compounds which ionise, e.g., $HCl + H_2O \rightarrow H_3O^{+} + Cl^{-}$).

Coordinate (Dative) Bonding

• Definition: A bond where the shared pair of electrons is provided entirely by one atom (the Donor) but shared by both atoms (the Acceptor).

• Meaning of Lone Pair: A pair of electrons not shared with any atom in the normal state but available for coordinate bonding.

• Conditions for Formation:

  1. One atom must have at least one lone pair of electrons.

  2. The other atom/ion must be short of at least one pair of electrons.

• Formation of Ammonium Ion ($NH_4^+$):

  • Ammonia ($NH_3$) has a lone pair on the Nitrogen atom.

  • It combines with a Hydrogen ion ($H^+$), which has no electrons.

  • The shared pair creates a coordinate bond ($→$).

  • Once formed, all four $N-H$ bonds are identical.

  • Note: $NH_4Cl$ contains all three bond types: covalent, coordinate, and ionic.

• Formation of Hydronium Ion ($H_3O^+$):

  • Water ($H_2O$) has two lone pairs on the Oxygen atom.

  • When an acid is added to water, $H^+$ is released and attaches to a lone pair on oxygen: $H^{+} + H_2O \rightarrow H_3O^{+}$.

• Formation of Hydroxyl Ion ($OH^-$):

  • Formed when one $H^+$ is removed from a water molecule: $H_2O \rightarrow H^{+} + OH^{-}$.

  • The shared pair remains with the more electronegative oxygen, giving the ion a negative charge.

Physical Observations and Periodic Trends

• Electricity Experiment:

  • Solutions of $MgCl_2$, $NaCl$, $NaOH$, and $CuSO_4$ (ionic) allow current to pass, making a bulb glow.

  • Solutions of distilled water, sugar, alcohol, chloroform, benzene, and petrol (covalent) do not allow current to pass.

• Periodic Table Trends (Period 3 Chlorides):

  • Group I ($NaCl$): Ionic Solid.

  • Group II ($MgCl_2$): Ionic Solid.

  • Group III ($AlCl_3$): Partially ionic-partially covalent solid.

  • Group IV ($SiCl_4$): Covalent liquid.

  • Group V ($PCl_3/PCl_5$): Covalent liquid/solid.

  • Group VI ($S_2Cl_2$): Covalent liquid.

• Periodic Table Trends (Period 3 Oxides):

  • $Na_2O$, $MgO$, $Al_2O_3$ are Ionic solids.

  • $SiO_2$, $P_2O_5$, $SO_2/SO_3$, $Cl_2O_7$ are Covalent solids/gases.