Chapter 4: Chemical Bonding and Molecular Geometry

Chapter 4 Chemical Bonding and Molecular Geometry

4.1 Ionic Bonding

  • Objectives:

    • Explain the formation of cations, anions, and ionic compounds

    • Predict the charge of common metallic and nonmetallic elements, and write their electron configurations

  • Ions:

    • Definition: Atoms or molecules bearing an electrical charge.

    • Cations: Positive ions formed when neutral atoms lose electrons (e.g., Na+).

    • Anions: Negative ions formed when neutral atoms gain electrons (e.g., Cl−).

  • Ionic Compounds:

    • Compounds composed of ions, held together by ionic bonds.

    • Nature of Ionic Bonds: Electrostatic forces of attraction between oppositely charged ions.

    • Properties of Ionic Compounds:

    • Crystalline solids, rigid, brittle, high melting and boiling points.

    • Poor conductors of electricity in solid state, but good conductors when dissolved in water or melted.

    • Example: Sodium chloride (NaCl) is formed from sodium (Na) and chlorine (Cl).

      • Sodium (a soft metal) and chlorine (a poisonous gas) react vigorously to form NaCl, a non-toxic, essential compound for life.

      • Sodium must be stored in oil due to its reactivity; chlorine is corrosive and dangerous.

  • Formation of Ionic Compounds:

    • Binary Ionic Compounds: Consisting of two elements, one metal (cations) and one nonmetal (anions).

    • Periodic Properties:

    • Metals lose electrons easily (low ionization potentials, left side of periodic table).

    • Nonmetals gain electrons to fill valence shells (high electron affinities, upper-right corner of periodic table).

  • Charge Neutrality:

    • Total positive charges must equal total negative charges in an ionic compound.

    • Compound Formula: Represents the simplest ratio of ions (e.g., aluminum oxide, Al2O3, contains two Al3+ for every three O2−).

    • Example: Ionic bond strength in sodium chloride requires 769 kJ to dissociate its ions:

      NaCl(s) \rightarrow Na^+(g) + Cl^-(g) \; ; \Delta H = 769 \text{ kJ}

  • Electronic Structures of Cations:

    • Main group elements lose valence electrons to resemble the electron configuration of preceding noble gases.

    • For groups 1 and 2, cation charge equals group number.

    • Example: Calcium (Ca) forms Ca2+ with the configuration 1s²2s²2p⁶ (isoelectronic with Ar).

    • Exceptions: Elements like Tl, Sn, Pb, and Bi can form ions with different charges due to the inert pair effect.

  • Electronic Structures of Anions:

    • Formed when nonmetals gain electrons to fill their s and p orbitals.

    • Example: Oxygen configuration 1s²2s²2p⁴ forms O2− (oxide ion) with 1s²2s²2p⁶.

4.2 Covalent Bonding

  • Objectives:

    • Describe the formation of covalent bonds

    • Define electronegativity and assess the polarity of covalent bonds

  • Covalent Bonds:

    • Formed by mutual attraction of atoms for shared electrons.

    • Example: Hydrogen (H2) molecule involves two hydrogen atoms sharing a pair of electrons.

  • Properties of Covalent Compounds:

    • Generally lower melting and boiling points compared to ionic compounds.

    • Many exist as liquids or gases at room temperature, typically softer than ionic solids.

    • Poor conductors of heat and electricity due to neutrality of covalent bonds.

  • Formation of Covalent Bonds:

    • Atoms overlap and share electrons to decrease potential energy.

    • Energy is released when chemical bonds are formed and must be added to break them.

    • Example: Energy required to break H2 bond is 436 kJ.

  • Pure versus Polar Covalent Bonds:

    • Pure covalent bonds occur in identical atoms (e.g., H2, Cl2).

    • Polar covalent bonds form between different atoms leading to unequal sharing of electrons (e.g., HCl).

  • Electronegativity:

    • Definition: Measure of the atom's ability to attract electrons in a bond.

    • Higher electronegativity leads to partial negative charge in polar covalent bonds.

    • The electronegativity values increase across a period and decrease down a group. Fluorine is the most electronegative (4.0).

  • Bond Polarity Based on Electronegativity:

    • Difference in electronegativity (ΔEN) determines bond type:

    • Nonpolar covalent (ΔEN = 0), polar covalent (ΔEN > 0), ionic (ΔEN >> 1).

4.3 Chemical Nomenclature

  • Objectives:

    • Derive names for common types of inorganic compounds using a systematic approach

  • Chemical Nomenclature:

    • Ionic compounds are named by combining the cation name with the anion name (e.g., NaCl: sodium chloride).

    • Use Roman numerals for metals with variable charges (e.g., iron(II) chloride).

  • Examples of Naming Ionic Compounds:

    • Monatomic ions: NaCl (sodium chloride), Na2O (sodium oxide).

    • Polyatomic ions: KNO3 (potassium nitrate), CaSO4 (calcium sulfate).

  • Everyday Ionic Compounds:

    • Commonly found in products:

    • NaCl → table salt; NaHCO3 → baking soda.

4.4 Lewis Symbols and Structures

  • Objectives:

    • Write Lewis symbols for neutral atoms and ions

    • Draw Lewis structures depicting the bonding in simple molecules

  • Lewis Symbols:

    • Consist of an elemental symbol surrounded by dots representing valence electrons.

    • Used for describing valence electron configurations and depicting ionic and covalent bonds.

  • Lewis Structures:

    • Illustrates bonding between atoms (shared pairs) and lone pairs.

    • Rule: Main group elements form enough bonds to achieve eight electrons (octet rule).

  • Multiple Bonds:

    • Single, double, and triple bonds correspond to shared pairs:

    • Example: C=O in CO, where carbon forms double bonds.

  • Step-by-Step Method for Writing Lewis Structures:

    1. Count total valence electrons.

    2. Draw a skeleton structure.

    3. Complete octets for terminal atoms.

    4. Place remaining electrons on central atom.

    5. Rearrange electrons to form multiple bonds if necessary.

  • Special Cases:

    • Molecules with odd-electron counts, electron-deficient, and hypervalent molecules do not adhere to octet rule.

4.5 Formal Charges and Resonance

  • Objectives:

    • Compute formal charges for atoms in any Lewis structure

    • Explain resonance

  • Formal Charge:

    • Calculation method to predict most reasonable Lewis structure:

    • Use:
      \text{Formal Charge} = \text{Valence Electrons} - \text{Lone Pair Electrons} - \frac{1}{2}\text{Bonding Electrons}

  • Resonance:

    • Occurs when two or more Lewis structures with the same arrangement can be drawn.

    • The actual structure is the average of these resonance forms.

    • Example: The nitrate ion (NO3−) can be represented by different resonance forms.

4.6 Molecular Structure and Polarity

  • Objectives:

    • Predict the structures of small molecules using VSEPR theory

    • Explain molecular polarity and assess polarity based on bonding and structure

  • VSEPR Theory:

    • Predicts molecular geometry based on electron pairs (bonding and lone) repulsing one another.

    • Electron-pair geometries: linear, trigonal planar, tetrahedral, trigonal bipyramidal, and octahedral.

  • Molecular Polarity:

    • Determined by the bond dipoles and molecular geometry.

    • Polar molecules have an uneven distribution of charge.

    • Condition for Polarity: 1. Contain at least one polar covalent bond.

    1. geometry must not allow dipole cancels.

Conclusion

  • Ionic and covalent bonds play crucial roles in the structure and properties of compounds. Understanding these concepts is pivotal in chemistry and related applications.

  • Recognizing the distinctions between ionic and covalent compounds facilitates a better grasp of their behaviors and applications in various contexts.