Bonding and Lewis Dot Structures Notes

Bonding and Lewis Dot Structures

• Context: This discussion focuses on key concepts in chemical bonding, including electronegativity, the different types of chemical bonds (nonpolar covalent, polar covalent, ionic), and the application of Lewis dot structures to represent atoms, ions, and compounds.

1. Electronegativity

• Definition: Electronegativity (EN) is a measure of the ability of an atom in a chemical compound to attract electrons towards itself. It is a fundamental property that dictates the nature of chemical bonds.
• Periodic Trends: Electronegativity generally:
  • Increases across a period (left to right): As you move from left to right across a row of the periodic table, the nuclear charge increases, and electrons are added to the same principal energy level. This stronger attraction pulls bonding electrons closer to the nucleus, increasing electronegativity.
  • Decreases down a group (top to bottom): As you move down a column, the atomic radius increases due to the addition of more electron shells. The increased distance between the nucleus and the valence electrons, along with increased shielding by inner electrons, reduces the nuclear attraction for bonding electrons, thus decreasing electronegativity.
• Impact on Bond Character: The difference in electronegativity ($\Delta EN$) between two bonded atoms is crucial for determining the type of chemical bond formed:
  • Small $\Delta EN$: Leads to covalent bonds.
  • Larger $\Delta EN$: Leads to polar covalent bonds.
  • Very large $\Delta EN$: Leads to ionic bonds.

2. Types of Chemical Bonds

• Chemical bonds are forces that hold atoms together in molecules or compounds, primarily involving valence electrons.
• Nonpolar Covalent Bonds:
  • Formation: Occur when electrons are shared equally between two atoms.
  • Electronegativity Difference: Characterized by a very small or zero difference in electronegativity ($\Delta EN \approx 0$).
  • Examples: Typically found between identical nonmetal atoms (e.g., $H<em>2$, $O</em>2$, $Cl_2$).
• Polar Covalent Bonds:
  • Formation: Occur when electrons are shared unequally between two atoms due to a significant difference in their electronegativities.
  • Electronegativity Difference: Characterized by an intermediate difference in electronegativity (0 < \Delta EN < 1.7, though this range can vary slightly depending on the source).
  • Result: The more electronegative atom attracts the shared electron pair more strongly, developing a partial negative charge ($\delta^-$), while the less electronegative atom develops a partial positive charge ($\delta^+$).
  • Examples: $H<em>2O$, $HCl$, $NH</em>3$.
• Ionic Bonds:
  • Formation: Occur when there is a complete transfer of one or more valence electrons from one atom to another, resulting in the formation of charged ions (cations and anions).
  • Electronegativity Difference: Characterized by a very large difference in electronegativity ($\Delta EN \ge 1.7$).
  • Result: Electrostatic attraction between oppositely charged ions holds the compound together.
  • Examples: Typically formed between a metal (which loses electrons to form a cation) and a nonmetal (which gains electrons to form an anion) (e.g., $NaCl$, $MgO$).

3. Lewis Dot Structures

• Purpose: Lewis dot structures (also known as Lewis structures or electron-dot structures) are diagrams that show the bonding between atoms of a molecule and the lone pairs of electrons that may exist in the molecule. They provide a simple visual representation of valence electrons and their arrangement.
• Representation: Dots are used to represent valence electrons, and lines are used to represent shared electron pairs (covalent bonds).
• Drawing Steps (General):
  1. Count Total Valence Electrons: Sum the valence electrons from all atoms in the molecule or ion. For anions, add electrons equal to the negative charge; for cations, subtract electrons equal to the positive charge.
  2. Determine Central Atom: Usually the least electronegative atom (never hydrogen).
  3. Draw Skeletal Structure: Connect the central atom to terminal atoms with single bonds (each single bond uses 2 electrons).
  4. Distribute Remaining Electrons: Place lone pairs on terminal atoms first to satisfy their octets (or duets for hydrogen).
  5. Place Remaining Electrons on Central Atom: Any leftover electrons are placed on the central atom as lone pairs.
  6. Form Multiple Bonds (if necessary): If the central atom does not have an octet, move lone pairs from terminal atoms to form double or triple bonds until the central atom achieves an octet (or expanded octet for period 3 and below).
• Examples:
  • Atom: Carbon (C) has 4 valence electrons, so its Lewis structure is $C \cdot$ (with 4 dots around it).
  • Ion: Oxide ion ($O^{2-}$) has 6 (from oxygen) + 2 (from charge) = 8 valence electrons. Its Lewis structure is $[\cdot \cdot O \cdot \cdot]^{2-}$ (with 8 dots representing 4 lone pairs around O within brackets, and a $2-$ charge).
  • Compound: Methane ($CH_4$).
  1. Valence electrons: C (4) + 4H (4x1) = 8 electrons.
  2. C is central.
  3. Single bonds to 4 H atoms: 4 bonds x 2 e-/bond = 8 electrons used.
  4. All electrons used, and C has an octet, H has a duet.
  5. Lewis structure: H-C-H (with two H's above and below C, each bond being a line).