Chemistry: Chemical Bonding and Structure

Remedial Special Class UNIT 2: Chemical Bonding and Structure

INTRODUCTION

  • Octet Rule: The octet rule states that during the formation of a chemical compound, each atom achieves an octet (8) of electrons in its highest occupied energy level by gaining, losing, or sharing electrons.
  • Chemical bonds are foundational to nearly everything a person sees or touches, including air, food, and clothing.

NOBLE GASES

  • Noble gases possess very stable electron arrangements:
    • Electron Configurations:
    • Helium (He): 2 electrons
    • Neon (Ne): 2, 8 electrons
    • Argon (Ar): 2, 8, 8 electrons
    • They have high ionization energies, low electron affinity, and a general lack of reactivity.
    • Valence Electrons: All noble gases (except He) have eight valence electrons.
    • The common electron configuration for noble gases (except Helium) is denoted as n s^2 n p^6, termed as an octet.

BONDING AND EXCEPTIONS TO THE OCTET RULE

  • When forming bonds, atoms strive to achieve the electron configuration of the nearest noble gas ( extit{ns$^2$np$^6$}).
  • Examples of Compounds:
    • Compounds obeying the octet rule: CH₄ (methane), NH₃ (ammonia)
    • Exceptions to the octet rule include: BeCl₂, BF₃, SF₆, PCl₃, PCl₅, which, although stable, do not adhere strictly to the octet rule, featuring either a shortfall or excess of electrons.

TYPES OF CHEMICAL BONDING

  1. Chemical Bonds: Forces of attraction holding atoms together can be categorized into intramolecular forces, which influence the chemical properties of species.
    • Main types of chemical bonds: covalent, ionic, and metallic bonds.
    • The bonding type depends on how atoms combine.

IONIC BONDING

  • Formation of Ionic Compounds: Typically occurs between metal cations and non-metal anions. An exception is the ammonium ion, which is not a metal but forms ionic compounds.
  • Example: The reaction of sodium metal with chlorine gas:
    • Reaction: ext{Na (s)} + ext{Cl}_2(g)
      ightarrow ext{NaCl (s)}
    • Sodium ion (Na⁺) forms by losing an electron, achieving the noble gas configuration of neon.
    • Chloride ion (Cl⁻) forms by gaining an electron, achieving the configuration of argon.

Electrostatic Force of Attraction

  • Exists between oppositely charged particles, such as in sodium chloride, where Na⁺ and Cl⁻ are attracted to each other, forming an ionic bond.

Lewis Electron-Dot Symbols

  • In Lewis symbols, the element symbol represents the nucleus and inner electrons, while dots represent valence electrons.
    • Example for Phosphorus: Three single dots around the element's symbol.

Formation of Ionic Bonding

  • Depends on:
    • Low ionization energy of metals.
    • High electron affinity of non-metals.
  • This process is exothermic, and the stability of ionic compounds is attributed to the packing of oppositely charged ions.
  • Lattice Energy (U):
    • Defined as the energy change when gaseous ions form one mole of a solid ionic compound or the energy required to separate one mole of the ionic substance into gaseous ions.
    • Indicative of the strength of ionic interactions and affects properties such as melting point, hardness, and solubility.
    • Determined using Hess's law of heat summation.

HESS'S LAW OF HEAT SUMMATION

  • States that the enthalpy change of an overall reaction equals the sum of the enthalpy changes for individual reactions.
  • Example Reaction Sequence for Sodium Fluoride (NaF):
    • ext{Na(g) + F(g)}
    • ext{Na(s) + F}2(g) ightarrow ext{ΔH}{total} = ext{ΔH}_1 + ext{ΔH}_2 + ext{ΔH}_3 + ext{…}
  • Figure illustrates the Born-Haber cycle for NaF(s).

FACTORS AFFECTING FORMATION OF IONIC BONDING

  1. Ionization Energy (IE):

    • Low ionization energy in metals favors ionic bond formation.
    • Example: Alkali metals form ionic compounds due to low IE.

  2. Electron Affinity (EA):

    • Higher electron affinities favor ionic bond formation.
    • Example: Halogens with high electron affinity tend to form ionic compounds with metals.

  3. Lattice Energy (U):

    • Larger lattice energy promotes ionic bond formation.
    • It measures the attractive forces between ions; depends directly on product of ionic charges (q_1 imes q_2) and inversely on distance (r) between ions.
    • Small ions with higher charges yield larger lattice energies.
    • If energy released exceeds energy absorbed during formation, ionic compound formation is favored.

EXCEPTIONS TO THE OCTET RULE IN IONIC COMPOUNDS

  • Deficient Electron Configurations:
    • Elements close to Helium (H, Li, Be, B) may seek a duplet configuration (2 electrons) instead of an octet due to unique electron configurations.
    • Examples: Compounds such as LiH, BeCl₂, and BF₃ exhibit stability despite having fewer than 8 electrons around the central atom.
  • More than an Octet (18-Electron Rule):
    • Transition and post-transition elements often follow the 18-electron rule due to d orbital involvement.
    • Larger losses of electrons are needed for these atoms to achieve noble gas configurations. Positively charged atoms lose highly principal quantum number electrons first.

PROPERTIES OF IONIC COMPOUNDS

  • Typically exist as crystalline solids at room temperature.
  • Crystalline ionic solids are often brittle, non-conductors of electricity, but can conduct when molten.
  • Have high melting and boiling points.
  • Generally non-volatile and soluble in inorganic solvents like water but insoluble in organic solvents.

COVALENT BONDING AND MOLECULAR GEOMETRY

  • Formation of Covalent Bonding: A covalent bond involves sharing of a pair of electrons between two atoms.
  • Examples of covalent molecules include HCl, H₂S, C₂H₄, N₂, CCl₄, BCl₃, H₂O, NH₃, SO₂, PCl₅, O₃.
  • Substances with covalent bonds are referred to as molecules.
  • Lewis Structures: Represent covalent bonding and shared electron pairs.

DRAWING LEWIS STRUCTURES

  1. Valence Electrons:

    • Determine the total number of valence electrons for each atom.
    • For polyatomic ions, adjust for charge: add for negative charges, subtract for positive charges.

  2. Skeletal Structure:

    • Place the most electropositive atom in central position.
    • Connect bonded atoms with electron-pair bonds (dashes). Hydrogen remains a terminal atom.

  3. Octet Rule Application:

    • Distribute remaining electrons to achieve octets for terminal atoms (except for hydrogen).

  4. Multiple Bonds:

    • If central atom lacks 8 electrons, consider moving electrons from terminal atoms to form double or triple bonds.
    • Example: Write the three equivalent Lewis structures for nitrate ion, NO₃⁻, and describe its resonance hybrid structure.

COORDINATE-COVALENT BONDING

  • A coordinate-covalent bond occurs when one atom donates both electrons to the electron pair shared with another atom.
  • Examples: O₃, NH₃, BF₃, POCl₃.

RESONANCE STRUCTURES

  • Multiple Lewis structures that represent the molecule's electron distribution. Shows movement of electrons between different atoms, clarifying the actual distribution.
  • Example: Ozone (O₃) has two resonance structures, satisfying the octet rule.

EXCEPTIONS TO THE OCTET RULE IN COVALENT BONDING

  • Deficiencies: Compounds with central atoms from Group IIA and IIIA such as BeCl₂, BF₃, AlCl₃ are sometimes stable with fewer than eight electrons.
  • More than Eight: Central atoms from Periods 3, 4, 5, and 6 (e.g., PF₅, SF₆, XeF₄) can accommodate more than 8 electrons due to available d orbitals.
  • Molecules with Odd Electrons: Molecules like ClO₂, NO, and NO₂ possess an odd number of valence electrons (19, 11, and 17 respectively) and have Lewis structures that reflect this.

POLAR AND NON-POLAR COVALENT MOLECULES

  • Nonpolar Covalent Bonds: Where electrons are shared equally, e.g., H–H, Cl–Cl.
  • Polar Covalent Bonds: Result when electrons are shared unequally due to differing electronegativities, such as H–Cl, where electrons spend more time near chlorine, resulting in partial charges (delta+ and delta-).
  • Bond Moment and Dipole Moment:
    • Measures polarity in diatomic covalent bonds with the definition of dipole moment being the product of charge magnitude (δ) and distance (d) between charges.

PROPERTIES OF COVALENT COMPOUNDS

  • Many covalent compounds are in gas or liquid state at room temperature, exhibiting weak intermolecular forces.
  • Some, like iodine, are solids. Covalent compounds typically have low melting and boiling points and tend to be insoluble in water but soluble in nonpolar solvents.
  • Nonpolar covalent compounds do not conduct electricity effectively.

MOLECULAR GEOMETRY: VALENCE SHELL ELECTRON PAIR REPULSION (VSEPR) THEORY

  • Electron Pair Arrangement: Suggests that electron pairs around a central atom spread out to minimize repulsion.
  • Types of electron sets include single bonds, double bonds, triple bonds, and lone pairs.
  • Arrangement in three-dimensional space leads to specific molecular shapes.
  • Arrangement:
    • Linear for two electron sets.
    • Trigonal Planar for three sets.
    • Tetrahedral for four sets.
    • Trigonal Bipyramidal for five sets.
    • Octahedral for six sets.

GUIDELINES FOR APPLYING VSEPR MODEL

  1. Write the Lewis structure of the molecule considering only electron sets at the central atom.
  2. Count total electron sets (bonding pairs + lone pairs).
  3. Predict shape using VSEPR geometry rules.
  4. Predict bond angles based on geometry.

MOLECULAR POLARITY

  • Molecular shape affects properties such as polarity, melting points, boiling points, and reactivity.
  • Example: CO₂ is a nonpolar molecule despite having polar bonds, a consequence of its linear shape.
  • In contrast, H₂O is polar due to its bent shape, resulting in a net dipole moment.

PREDICTING SHAPES OF MOLECULES

Table of Electron Groups and Molecular Geometry

  • AX₂: Linear, bond angle 180°.
  • AX₃: Trigonal planar, bond angle 120°.
  • AX₂E: Trigonal planar (bent), bond angle <120°.
  • AX₄: Tetrahedral, bond angle 109.5°.
  • AX₃E: Tetrahedral (trigonal pyramidal), bond angle <109.5°.
  • AX₂E₂: Tetrahedral (bent), bond angle <109.5°.
  • AX₅: Trigonal bipyramidal, with significant angles of 120° (equatorial) and 90° (axial).

INTERMOLECULAR FORCES IN COVALENT COMPOUNDS

  1. Intramolecular forces: Chemical bonds within a molecule (ionic, covalent, metallic).
  2. Intermolecular forces: Attractions between molecules, weaker compared to intramolecular forces.
    • Types of Intermolecular Forces:
      • Dipole-Dipole Forces: Present between polar molecules, responsible for properties like boiling points.
      • Hydrogen Bonding: A strong type of dipole-dipole interaction involving hydrogen bonded to electronegative atoms (F, O, N).
      • London Dispersion Forces: Intermolecular forces that occur in nonpolar molecules by inducing dipoles.

METALLIC BONDING

  • Defined as the interaction between metal nuclei and delocalized electrons (also known as conduction electrons).
  • Metallic bonds resemble a 'sea' of electrons surrounding positively charged metal kernels.
  • The bond strength depends on:
    1. Number of delocalized electrons: More delocalized electrons enhance bonding strength and elevate melting point.
    2. Packing arrangement: Tighter atomic packing increases interatomic forces and bond strength.
  • Unlike ionic bonding, metal ions in metallic bonds are more flexible in position due to electron mobility.

PROPERTIES OF METALS

  • Metals conduct electricity and heat due to mobile electrons.
  • Typically feature strong, opaque structures, malleable and ductile with high melting and boiling points.
  • Exhibit a wide range of melting points, from -39°C (mercury) to 3410°C (tungsten).

CHEMICAL BONDING THEORIES

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