Study Notes on Models of Chemical Bonding

Chapter 9: Models of Chemical Bonding

9.1 Atomic Properties and Chemical Bonds

  • Overview of chemical bonds, their types, and significance in chemistry.

9.2 Types of Chemical Bonding

  • Ionic bonding:

    • Involves the transfer of electrons.

    • Typically occurs between metals and nonmetals.

  • Covalent bonding:

    • Involves the sharing of electrons.

    • Typically occurs between nonmetals bonded to nonmetals.

  • Metallic bonding:

    • Involves electron pooling.

    • Occurs when a metal bonds with another metal.

9.3 Bond Energy and Chemical Change

  • Bond energy: Defined as the energy required to break a bond.

    • Relates to the stability of the bond; higher energy indicates a stronger bond.

9.4 Between the Extremes: Electronegativity and Bond Polarity

  • Electronegativity: The ability of an atom to attract shared electrons.

    • Polar covalent bonds occur when electrons are shared unequally.

    • The more electronegative atom gains a partial negative charge, while the other generates a partial positive charge.

9.5 An Introduction to Metallic Bonding

  • Electron sea model:

    • Metal atoms contribute valence electrons to form a delocalized electron “sea”.

    • Metal ions are arranged orderly within this mobile electron sea, providing metals their conductive properties.

9.6 Summary of Key Concepts

  • Comparison of metals and nonmetals based on bonding behavior.

Comparisons of Metals and Nonmetals

  • Metals tend to be hard, rigid, and brittle with high melting points.

  • Ionic compounds do not conduct electricity in their solid state; however, they conduct when melted or dissolved in water due to free-moving ions.

9.7 Lewis Electron-Dot Symbols

  • Lewis dot structures represent valence electrons for main-group elements:

    • The group number indicates the number of valence electrons. Example:

    • Nitrogen (N) in Group 15 has 5 valence electrons.

Drawing Lewis Symbols:

  1. Place one dot on each of the four sides of the element symbol until all valence electrons are represented.

  2. Pair them as necessary.

  • For metals, the total number of dots in the Lewis symbol equals the number of electrons lost (cation formation).

  • For nonmetals, unpaired dots signify the number of electrons gained to form an anion or shared for covalent bonds.

  • The octet rule states that atoms will lose, gain, or share electrons to achieve eight valence electrons (or two for hydrogen and lithium).

9.8 The Ionic Bonding Model

  • An ionic bond forms when a metal transfers electrons to a nonmetal, creating ions that attract each other to form a solid compound.

    • Total electrons lost by metals equal total electrons gained by nonmetals.

  • Example Problem 9.1:

    • Problem: Depict ion formation using Lewis symbols.

    • Solution: Sodium loses one electron; oxygen gains two; thus, two sodium ions are needed for each oxide ion, forming Na₂O.

9.9 Lattice Energy

  • Lattice energy is the energy required to separate 1 mole of an ionic solid into gaseous ions.

    • It measures the strength of the ionic bond.

  • Coulomb’s Law: determines the strength of interactions based on charges and ionic size.

Factors influencing Lattice Energy:

  • Ionic Size:

    • Smaller ions interact more effectively, leading to higher lattice energy.

  • Ionic Charge:

    • Higher charges lead to stronger attractions and higher lattice energy.

9.10 Properties of Ionic Compounds

  • Ionic compounds are usually hard and brittle, with high melting points.

  • Conduct electricity when molten or dissolved because ions are mobile.

  • They crack under pressure due to repulsion between like charges in the crystal lattice.

9.11 Conductivity and Melting Points of Ionic Compounds

  • Melting point (°C) and boiling point (°C) of various ionic compounds:

    • CsBr: mp 636, bp 1300

    • NaCl: mp 801, bp 1413

    • MgO: mp 2852, bp 3600

9.12 Covalent Bonding

  • Covalent bonds involve sharing of electron pairs between atoms.

  • Different types of bonds exhibit different properties and bond orders.

Bond Orders:

  • Single bond: One pair of electrons shared (bond order = 1).

  • Double bond: Two pairs of electrons shared (bond order = 2).

  • Triple bond: Three pairs of electrons shared (bond order = 3).

9.13 Bond Length and Energy Trends

  • Higher bond orders correspond to shorter bond lengths and greater bond energies.

  • Bond lengths increase down a group; bond energies tend to decrease.

9.14 Sample Problems on Bonding Concepts

  • Sample Problem 9.3: Rank bonds based on length and strength.

    • Plan to determine order using atomic size and bond order relations.

Conclusion

  • Understanding bonding types, properties, and energies is crucial for predicting the behavior of materials and understanding chemical reactions.

Connections to Real-World Applications

  • Knowledge of ionic and covalent bonds is fundamental in fields such as materials science, organic chemistry, and biochemistry as it dictates reactivity, properties, and applications of compounds and materials.