General Chemistry- 2

Pauli’s Exclusion Principle

• States that no two electrons in an atom can have the same set of four identical quantum numbers.

• In the first shell (n = 1), there is only one orbital (1s) which can accommodate a maximum of two electrons, provided they have opposite spins.

Electron Configurations in Shells

Second Shell Configurations

• The second shell (n = 2) has four orbitals: one s orbital (l = 0) and three p orbitals (l = 1).

• This shell can hold a total of eight electrons, with configurations spread across the different orbitals according to quantum numbers.

Higher Shell Configurations

• d sublevels can accommodate up to ten electrons.

• f sublevels may hold up to fourteen electrons.

• Understanding these configurations is crucial for calculating the maximum number of electrons that can be accommodated in any shell according to Pauli's exclusion principle.

Hund’s Rule of Maximum Multiplicity

• States that electrons are distributed among the orbitals of a subshell to maximize the number of unpaired electrons, ensuring they all have the same direction of spin.

• Electrons fill orbitals singly before pairing begins, which is demonstrated through filling diagrams of n = 1 and n = 2 shells.

Ground-State Electron Configuration of Elements

Key Rules for Electron Configuration

• Rule 1: Each shell can hold a maximum of 2n² electrons, where n is the shell number.

• Rule 2: Electrons occupy s, p, d, and f orbitals according to their capacity: s = 2, p = 6, d = 10, f = 14.

• Rule 3: Electrons fill available orbitals in increasing order of energy (Aufbau Principle).

Increasing Energy of Orbitals

• The general sequence of increasing energy levels is: 1s < 2s < 2p < 3s < 3p < 4s < 3d < 4p < 5s < 4d < 5p < 6s < 4f < 5d.

• This energy order illustrates how subshells fill based on their energy levels.

Orbital Occupancy Rules

• Rule 4: Each orbital may hold a maximum of two electrons with opposite spins.

• Rule 5: In degenerate orbitals, electrons prefer to occupy separate orbitals before they pair, following Hund’s Rule.

Electron Configuration Table

Classification of Elements in the Periodic Table

• Elements are grouped in increasing order of atomic number into horizontal periods (7) and vertical groups (18).

• Elements are classified into four blocks based on the subshell in which their valence electrons reside:

  • s-block: Groups 1 and 2 (alkali and alkaline earth metals)

  • p-block: Groups 13 to 18 (includes metals, metalloids, and nonmetals)

  • d-block: Transition metals (elements in groups 3 to 12)

  • f-block: Inner transition metals (lanthanides and actinides).

Modern Periodic Law

• The chemical and physical properties of elements are periodic functions of their atomic numbers, proposed by Henry Moseley in 1913.

Electronic Configuration in Periods

• The principal quantum number (n) determines the period of an element. Elements in a period have the same number of atomic orbitals.

s-Block Elements

• Valence electrons enter the s-orbital.

• Elements in Group 1 (alkali metals) and Group 2 (alkaline earth metals) follow the ns¹ and ns² configuration, respectively.

• These are soft metals, reactive, with low ionization enthalpies and are generally not found in pure forms due to high reactivity.

p-Block Elements

• Valence electrons enter the p-orbital, comprising Groups 13 to 18, termed Representative Elements or Main Group Elements.

• The outer electronic configuration typically is ns² np², and non-metals and metalloids are exclusive to the p-block. Non-metallic character decreases down the group, while each period concludes with a noble gas configuration (ns² np⁶).

d-Block Elements

• Called transition elements, where valence electrons enter the d-orbital.

• The electronic configuration for these elements is represented as (n-1)d¹-10 ns¹-2 and spans groups 3 to 12 making up three series: 3d, 4d, and 5d.

f-Block Elements

• Consists of lanthanides and actinides.

• The electronic configuration of lanthanides follows 6s² 4f¹-14, while actinide electronic configurations can be irregular. Inner transition elements comprise the f-block.

Ionic Bonding

• Formed through the transfer of electrons from one atom to another, creating cations and anions which are then held together through electrostatic attraction.

Conditions for Ionic Bond Formation

1. Valence Electrons: One atom should possess 1, 2, or 3 valence electrons while the other should have 5, 6, or 7.

2. Electronegativity Difference: A large difference in electronegativities (≥2) between atoms is essential to form ionic bonds. For example, sodium (Na) has an electronegativity of 0.9 and chlorine (Cl) has 3.0, with a difference of 2.1 facilitating bond formation.

Net Energy Change in Ionic Bond Formation

• For stable ionic compounds, energy release must exceed the energy required for electron transfer:

  • (a) Ionization energy must be low for atom A.

  • (b) Electrons need high electron affinity for atom B.

  • (c) High electrostatic attraction energy leads to bond formation when net energy change is positive.

Characteristics of Ionic Compounds

1. Solids at room temperature.

2. High melting and boiling points.

3. Hard and brittle.

4. Generally soluble in water.

5. Poor electrical conductors in solid form.

6. No isomerism.

7. Ionic reactions tend to occur rapidly.

Covalent Bond

• Proposed by G.N. Lewis, a covalent bond forms when two atoms share electron pairs to attain stability.

• Compounds containing covalent bonds are termed covalent compounds.

Conditions for Covalent Bond Formation

1. Valence Electrons: Both atoms should have 5, 6, or 7 valence electrons for stable configurations.

2. Equal Electronegativity: Ideally should be equal or nearly equal to facilitate electron sharing without transferring.

3. Equal Sharing: Best when atoms' affinities for bonding electrons are similar.

Examples of Covalent Compounds

• Examples include:

• Water (H₂O): Formed by sharing pairs of electrons with covalent bonds.

• Methane (CH₄): Each hydrogen atom shares electrons with carbon to form strong covalent bonds.

Characteristics of Covalent Compounds

1. Can exist as gases, liquids, or solids at room temperature.

2. Generally exhibit low melting and boiling points.

3. Not hard or brittle.

4. Soluble in organic solvents.

5. Do not conduct electricity well.

6. Show isomerism in several cases.

7. Involve molecular reactions.

Coordinate Covalent Bond

• Defined as a covalent bond where both shared electrons originate from one atom, forming coordinate compounds. The atom contributing both electrons is called the ligand.

Examples of Coordinate Covalent Bonds

• Ammonium ion (NH₄⁺), hydronium ion (H₃O⁺), and others.

Polar Covalent Bonds

• Occur when electrons are shared unequally, leading to partial positive and negative charge distribution across the molecule.

• Example includes water (H₂O) where oxygen is more electronegative than hydrogen, resulting in a polar molecule.

Nonpolar and Polar Covalent Bonds

• Nonpolar bonds result from equal sharing; polar bonds arise when electron sharing is unequal, influenced by the atoms' electronegativities.

Hydrogen Bonding

• Involves attraction between a hydrogen atom bonded to a highly electronegative atom and lone pairs of electrons on another electronegative atom.

Characteristics of Hydrogen Bonds

1. Formed mainly by O, N, and F due to their high electronegativities.

2. Weaker than covalent bonds (bond energies < 10 kcal/mole).

3. Can produce chains or clusters of molecules linked through hydrogen bonding.

4. Directional nature due to hybridization and orbital orientation.

Types of Hydrogen Bonds

• Intermolecular: Between different molecules of the same compound (e.g., in water).

• Intramolecular: Occurs within a single molecule (e.g., in 2-nitrophenol).

Examples of Hydrogen Bonds

• Found in water (H₂O) and ammonia (NH₃) demonstrating unique structural properties attributed to hydrogen bonding.

Characteristics of Hydrogen-Bonded Compounds

1. Exhibit abnormally high boiling and melting points compared to other compounds of similar molecular weights.

2. Often show high solubility in polar solvents like water.

3. Form three-dimensional crystalline lattices, exemplified by ice, where molecules organize into tetrahedral networks.