Chapter 4: Introduction to Covalent Bonding, Molecular Geometry, and Bond Polarity

Course Announcements and Logistics

• Date: September 25, 2026.
• Assignments Due Sunday, September 27:
  • Week 5 Adaptive Learning Assignment.
  • Week 5 Homework.
  • Week 5 Reflection.
• Exam 1 Reflection:
  • Open for extra credit: provides an additional $+4\%$ added to the Exam 1 grade.
  • Due date: Monday, September 28.
• Exam 2 Details:
  • Date: Monday, October 5 during regular class time.
  • Coverage: Chapter 3 and Chapter 4.
• Upcoming Class Cancellation:
  • Friday, October 2: No class held and no office hours scheduled due to a faculty conference.
• Supplemental Instruction (SI) Sessions:
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  • Location: 60 W Charl 273.
• MASS Center Study Tables:
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Chapter 4 Learning Objectives and Simulation Tools

• Core Learning Objectives:
  • Predict the molecular shape around any given central atom in a molecule.
  • Utilize electronegativity values to determine whether a chemical bond is nonpolar covalent, polar covalent, or ionic.
  • Determine whether an entire molecule is polar or nonpolar.
• Interactive Simulation Resource:
  • 3D Molecular Shape Simulation (PhET Interactive Simulations): https://phet.colorado.edu/en/simulations/molecule-shapes

Valence-Shell Electron-Pair Repulsion (VSEPR) Theory

• Fundamental VSEPR Postulate:
  • Electron groups around a central atom arrange themselves as far apart in three-dimensional space as possible to minimize repulsive electrostatic interactions between negative electron clouds.
• Definition of an Electron Group:
  • A "group" defined under VSEPR theory can be any of the following attached to a central atom:
    • A single covalent bond ($1$ group).
    • A double covalent bond ($1$ group).
    • A triple covalent bond ($1$ group).
    • An unshared lone pair of electrons ($1$ group).
• Definition of Bond Angle:
  • The angle formed by the nuclei of two surrounding atoms with the nucleus of the central atom at the vertex.

Detailed Molecular Geometries and Bond Angles

• Linear Molecular Geometry:
  • Group Criteria: Any central atom surrounded by exactly $2$ electron groups.
  • Group Composition: May consist of bonds or lone pairs (a double bond or triple bond counts as $1$ group).
  • Bond Angle: Exactly $180^\circ$.
  • Examples:
    • Carbon dioxide ($\text{CO}_2$): Central Carbon atom surrounded by $2$ double bonds ($2$ groups, $0$ lone pairs).
    • Hydrogen cyanide ($\text{HCN}$): Central Carbon atom surrounded by $1$ single bond and $1$ triple bond ($2$ groups, $0$ lone pairs).

• Trigonal Planar Molecular Geometry:
  • Group Criteria: Any central atom surrounded by exactly $3$ electron groups with $0$ lone pairs.
  • Bond Angle: Exactly $120^\circ$.
  • Examples:
    • Boron trifluoride ($\text{BF}_3$): Central Boron atom bonded to $3$ Fluorine atoms ($3$ single bonds).
    • Formaldehyde ($\text{H}_2\text{CO}$): Central Carbon atom bonded to $2$ Hydrogen atoms via single bonds and $1$ Oxygen atom via a double bond ($3$ groups total).

• Tetrahedral Molecular Geometry:
  • Group Criteria: Any central atom surrounded by exactly $4$ electron groups with $0$ lone pairs.
  • Bond Angle: Exactly $109.5^\circ$.
  • Structural Representation Rules:
    • Solid lines represent bonds located within the two-dimensional plane of the paper.
    • Dashed wedges represent bonds pointing backward behind the plane.
    • Solid wedges represent bonds pointing forward in front of the plane.
  • Example: Methane ($\text{CH}_4$), featuring $1$ central Carbon bonded to $4$ Hydrogen atoms.

• Trigonal Pyramidal Molecular Geometry:
  • Group Criteria: Any central atom surrounded by $4$ electron groups, consisting specifically of $3$ bonded atoms and $1$ unshared lone pair.
  • Bond Angle: Slightly less than $109.5^\circ$ (approximately $107^\circ$), due to nonbonding lone pairs occupying more spatial volume and exerting stronger repulsion on adjacent bonding pairs.
  • Example: Ammonia ($\text{NH}_3$).

• Bent Molecular Geometry:
  • Group Criteria: Any central atom surrounded by $4$ electron groups, consisting specifically of $2$ bonded atoms and $2$ unshared lone pairs.
  • Bond Angle: Slightly less than $109.5^\circ$ (approximately $105^\circ$), heavily compressed by double lone pair repulsions.
  • Example: Water ($\text{H}_2\text{O}$).

Reference Table for Common Molecular Shapes

• Table 4.2: Common Molecular Shapes Around Central Atoms:
  • Linear: Total Groups: $2$ | Bonded Atoms: $2$ | Lone Pairs: $0$ | Bond Angle: $180^\circ$ | Examples: $\text{CO}_2$, $\text{HCN}$
  • Trigonal Planar: Total Groups: $3$ | Bonded Atoms: $3$ | Lone Pairs: $0$ | Bond Angle: $120^\circ$ | Examples: $\text{BF}_3$, $\text{H}_2\text{CO}$
  • Tetrahedral: Total Groups: $4$ | Bonded Atoms: $4$ | Lone Pairs: $0$ | Bond Angle: $109.5^\circ$ | Example: $\text{CH}_4$
  • Trigonal Pyramidal: Total Groups: $4$ | Bonded Atoms: $3$ | Lone Pairs: $1$ | Bond Angle: $\sim 109.5^\circ$ (approx. $107^\circ$) | Example: $\text{NH}_3$
  • Bent: Total Groups: $4$ | Bonded Atoms: $2$ | Lone Pairs: $2$ | Bond Angle: $\sim 109.5^\circ$ (approx. $105^\circ$) | Example: $\text{H}_2\text{O}$

Molecular Geometry Analysis of Complex Molecules

• Phosphine ($\text{PH}_3$) Shape Determination:
  • Structure: Phosphorus has $5$ valence electrons; it shares $3$ electrons to form single covalent bonds with $3$ Hydrogen atoms, retaining $1$ lone pair.
  • VSEPR Count: $4$ total groups ($3$ atoms + $1$ lone pair).
  • Resulting Geometry: Trigonal pyramidal with a bond angle of approximately $107^\circ$.

• Glycolic Acid Local Geometries:
  • Glycolic acid is an alpha-hydroxy acid (AHA) commonly used in skin care formulations to clear dead skin cells.
  • Hydroxyl Oxygen Atom: Bound to $2$ atoms ($\text{C}$ and $\text{H}$) with $2$ lone pairs ($4$ total groups) $\rightarrow$ Bent geometry.
  • $\text{CH}_2$ Carbon Atom: Bound to $4$ surrounding atoms ($2$ Hydrogen atoms, $1$ Oxygen atom, $1$ Carbon atom) with $0$ lone pairs ($4$ groups) $\rightarrow$ Tetrahedral geometry.
  • Carbonyl Carbon Atom: Bound to $3$ surrounding atoms ($1$ Carbon atom, $1$ carbonyl Oxygen atom, $1$ hydroxyl Oxygen atom) with $0$ lone pairs ($3$ groups) $\rightarrow$ Trigonal planar geometry.

• Vitamin B6 Geometry Analysis:
  • Identification of Bent Geometry: Atom [1] (Oxygen atom attached to Hydrogen and ring Carbon, possessing $2$ lone pairs) exhibits Bent geometry.
  • Bond Angle Determination: Carbon [2] is bonded to $4$ separate atoms with $4$ single bonds ($4$ groups), making its local geometry Tetrahedral with bond angles of $109.5^\circ$.

Electronegativity and Chemical Bond Polarity

• Continuum of Chemical Bonding:
  • Pure Covalent (Nonpolar Covalent): Electrons are shared equally between atoms due to identical or nearly identical electronegativities.
  • Polar Covalent: Shared electrons are shifted toward the atom with higher electronegativity, creating a partial positive charge ($\delta+$) on the less electronegative atom and a partial negative charge ($\delta-$) on the more electronegative atom. A dipole moment arrow points toward the negative end with a cross on the positive end.
  • Pure Ionic: Electron density is completely transferred, creating distinct positive cations and negative anions.
  • Prevalence: Polar covalent bonds represent the predominant bond type in chemical structures, far outnumbering pure ionic or pure covalent bonds.

• Electronegativity Definition and Periodic Trends:

  • Definition: Electronegativity is a measure of an atom's attraction for electrons involved in a chemical bond.
  • Value Scale: Higher numeric values indicate a stronger pull for shared bonding electrons.
  • Periodic Trends:
    • Electronegativity increases moving from left to right across any period of the periodic table.
    • Electronegativity increases moving from bottom to top within any main group.
    • Fluorine ($\text{F}$) possesses the highest electronegativity value of 4.0$.\n\n\n\n* **Electronegativity Values of Elements**:\n * Group 1A: \text{H} = 2.1$,$\text{Li} = 1.0$,$\text{Na} = 0.9$,$\text{K} = 0.8$,$\text{Rb} = 0.8\n * Group 2A: \text{Be} = 1.5$,$\text{Mg} = 1.2$,$\text{Ca} = 1.0$,$\text{Sr} = 1.0\n * Group 3A: \text{B} = 2.0$,$\text{Al} = 1.5$,$\text{Ga} = 1.6$,$\text{In} = 1.7\n * Group 4A: \text{C} = 2.5$,$\text{Si} = 1.8$,$\text{Ge} = 1.8$,$\text{Sn} = 1.8\n * Group 5A: \text{N} = 3.0$,$\text{P} = 2.1$,$\text{As} = 2.0$,$\text{Sb} = 1.9\n * Group 6A: \text{O} = 3.5$,$\text{S} = 2.5$,$\text{Se} = 2.4$,$\text{Te} = 2.1\n * Group 7A: \text{F} = 4.0$,$\text{Cl} = 3.0$,$\text{Br} = 2.8$,$\text{I} = 2.5\n\n# Determining Bond Classification via Electronegativity Difference\n\n* **Table 4.3: Electronegativity Difference (\Delta\text{EN}) Rules**:\n * \Delta\text{EN} < 0.5\,\text{units}: **Nonpolar Covalent Bond** (Electrons are shared equally).\n * \Delta\text{EN} = 0.5\text{--}1.9\,\text{units}: **Polar Covalent Bond** (Electrons are shared unequally and pulled toward the more electronegative element).\n * \Delta\text{EN} > 1.9\,\text{units}: **Ionic Bond** (Electrons are transferred completely from the less electronegative atom to the more electronegative atom).\n\n\n\n* **Specific Bond Examples**:\n * \text{C-C}$Bond:$\Delta\text{EN} = |2.5 - 2.5| = 0.0 \rightarrow **Nonpolar bond**.\n * \text{C-H}$Bond:$\Delta\text{EN} = |2.5 - 2.1| = 0.4 \rightarrow$**Nonpolar bond** (The small electronegativity difference of$0.4$between Carbon and Hydrogen is ignored; all$\text{C-H} bonds are treated as nonpolar).\n * \text{C-O}$Bond:$\Delta\text{EN} = |3.5 - 2.5| = 1.0 \rightarrow$**Polar covalent bond** (Dipole pointing toward Oxygen:$\text{C}^{\delta+} - \text{O}^{\delta-}).\n\n# Bond Polarity Practice Problems\n\n* **Comparing Polarity of \text{Br-Cl}$vs.$\text{F-Br}**:\n * Given values: \text{EN}(\text{Br}) = 2.8$,$\text{EN}(\text{Cl}) = 3.0$,$\text{EN}(\text{F}) = 4.0$.
  • $\Delta\text{EN}(\text{Br-Cl}) = |3.0 - 2.8| = 0.2$.
  • $\Delta\text{EN}(\text{F-Br}) = |4.0 - 2.8| = 1.2$.
  • Conclusion: $\text{F-Br}$ is significantly more polar than $\text{Br-Cl}$.

• Ordering Bonds by Decreasing Polarity ($\text{B-F}$, $\text{N-F}$, $\text{C-F}$):

  • Given values: $\text{EN}(\text{B}) = 2.0$, $\text{EN}(\text{C}) = 2.5$, $\text{EN}(\text{N}) = 3.0$, \text{EN}(\text{F}) = 4.0$.\n * \Delta\text{EN}(\text{B-F}) = |4.0 - 2.0| = 2.0$.
  • \Delta\text{EN}(\text{C-F}) = |4.0 - 2.5| = 1.5$.\n * \Delta\text{EN}(\text{N-F}) = |4.0 - 3.0| = 1.0$.
  • Correct order of decreasing polarity: $\text{B-F} > \text{C-F} > \text{N-F}$.

• Identifying Nonpolar Bonds in Vitamin B6 Structure:
  • Label A ($\text{C-O}$ bond): $\Delta\text{EN} = |3.5 - 2.5| = 1.0 \rightarrow$ Polar bond.
  • Label B ($\text{O-H}$ bond): $\Delta\text{EN} = |3.5 - 2.1| = 1.4 \rightarrow$ Polar bond.
  • Label C ($\text{C=N}$ bond): $\Delta\text{EN} = |3.0 - 2.5| = 0.5 \rightarrow$ Polar bond.
  • Label D ($\text{C-H}$ bond): $\Delta\text{EN} = |2.5 - 2.1| = 0.4 \rightarrow$ Nonpolar bond.
  • Conclusion: Label D represents the nonpolar bond.