Bonding and Polarity

Bonding and Polarity Study Notes

Unit 5: Chapter 8 Overview

  • Key Concepts: Understanding bonding, electron configurations, electronegativity, molecular shapes, and intermolecular forces.

Recap: Octet Rule

  • Octet Rule Definition: Atoms tend to form compounds in such a way that each atom achieves an electron configuration similar to that of noble gases, typically having 8 electrons in their valence shell.

    • Valence Shell Electrons: Atoms prefer to have 8 electrons in their outermost shell to achieve stability.

Exceptions to the Octet Rule

  • Atoms with Less than Octet:

    1. Beryllium (Be): Often has 4 valence electrons in compounds.

    2. Boron (B): May form stable compounds with 6 valence electrons.

    3. Aluminum (Al): Sometimes has less than 8 electrons when covalently bonded.

    • Reason: These atoms have low electronegativity (EN) and do not strongly attract shared electrons. They never form multiple bonds.

  • Atoms with an Odd Number of Electrons:

    • Example: Nitrogen (N) or Chlorine (Cl) can have 7 electrons in valence shell. This is uncommon.

  • Atoms with More than an Octet:

    • Atoms in Period 3 and above can have more than 8 electrons by utilizing empty d orbitals.

    • Example: Sulfur (S), Phosphorus (P).

    • Explanation: C, N, O, and F always form stable configurations with 8 electrons.

Bonding Theory for Covalent Compounds

  1. Lewis Models: A representation of covalent bonding using dots for valence electrons.

    • Example: Lewis structure for Ethane (H-C-H).

  2. VSEPR Theory (Valence Shell Electron Pair Repulsion):

    • Definition: Electron pairs will arrange themselves to minimize repulsion, resulting in a specific molecular geometry.

    • Electronic Shape vs Molecular Shape: Electronic shapes account for lone pairs while molecular shapes consider bonding pairs and lone pair repulsion.

Electron Arrangements and Shapes

  • General Principles: What this means for molecular shape:

    • Electrons push each other away and arrange to maximize distance.

    • Definitions of Groups:

    • A lone pair counts as one group.

    • A shared pair (single, double, or triple bond) counts as one group.

Arrangements of Electron Groups

  • Two Groups of Electrons:

    • Linear Shape: Groups are 180° apart.

    • Examples: Beryllium hydride (BeH2), Carbon Dioxide (CO2).

  • Three Groups of Electrons:

    • Trigonal Planar Shape: Groups are 120° apart. If lone pairs exist, molecular shape is altered.

    • Examples: Boron trifluoride (BF3), Sulfur dioxide (SO2) has bent shape due to lone pair.

  • Four Groups of Electrons:

    • Tetrahedral Shape: Groups are 109.5° apart. Molecular shape differs with lone pairs.

    • Examples: Methane (CH4) – tetrahedral; Ammonia (NH3) – trigonal pyramidal; Water (H2O) – bent.

  • Five Groups of Electrons:

    • Trigonal Bipyramidal Shape: 120° (equatorial) and 90° (axial) arrangement.

    • Examples: Phosphorus pentafluoride (PF5) – trigonal bipyramidal; Seesaw, T-shaped, Linear shapes arise from lone pairs.

  • Six Groups of Electrons:

    • Octahedral Shape: All angles are 90°.

    • Examples: Sulfur hexafluoride (SF6) – octahedral; Square pyramidal (5 bonding, 1 lone pair); Square planar (4 bonding, 2 lone pairs).

Types of Molecular Geometries Summary

Table of Electron and Molecular Geometries based on number of electron groups and the presence of lone pairs is essential for predicting shapes.

Covalent Bonding

  • Definition: Covalent bonds are formed by the sharing of valence electrons between atoms.

  • Sample Query: In H2 versus HF, electrons are shared evenly because of equal electronegativity, while unequal in HF due to differing EN values.

Electronegativity Values and Bond Types

  • Difference in electronegativity indicates bond type:

    • 0 - 0.4: Nonpolar covalent.

    • 0.4 - 2.0: Polar covalent.

    • > 2.0: Ionic bond.

Unequal Sharing Leads to Polarity

  • Dipole Moment: Indicates the charge separation in a polar bond; greater EN difference leads to a stronger polarity.

  • Polarity Continuum:

    • Pure nonpolar (equal sharing) → Polar covalent (unequal sharing) → Ionic (transfer of electrons).

Molecule Polarity vs Bond Polarity

  • Molecules can contain polar bonds but be nonpolar overall. Check Lewis structure:

    • Asymmetrical Structure: Polar molecule.

    • Symmetrical Structure: Nonpolar molecule.

  • Example Inquiry: Assessing the polarity of H2O versus CO2 despite H-O being polar, comparing molecular symmetry.

Polarity Importance

  • Impact on Intermolecular Forces: Polarity heavily influences how particles interact via intermolecular forces (IMFs) versus intramolecular forces.

Types of Intermolecular Forces (IMFs)

  1. Dispersion Forces (London Dispersion Forces):

    • Arise from temporary dipoles; exist in all molecules; strength increases with size.

  2. Dipole-Dipole Interactions:

    • Exist in polar molecules due to the alignment of permanent dipoles; strength is size-independent.

  3. Hydrogen Bonding:

    • A strong form of dipole-dipole interaction occurring between hydrogen and highly electronegative atoms (O, N, F).

    • Considered a weak chemical bond; dominates in small molecules.

Strength of Intermolecular Forces

  • Order of strength for small molecules:

    • Dispersion Forces < Dipole-Dipole < Hydrogen Bonds.

    • For large molecules, LDFs may surpass both other forces. IMFs significantly affect physical properties like melting and boiling points.

Summary

  • Instantaneous vs Permanent Dipoles:

    • Instantaneous Dipole: Temporary, occurs due to momentary electron distribution.

    • Permanent Dipole: Always present due to electronegativity differences.

  • Examples:

    • Only Instantaneous Dipoles: Nonpolar molecules like H2, BH3, CH4.

    • Permanent Dipoles: Polar molecules like HCl, H2O, NH3.

Application Exercises

Evaluate the polarity and types of intermolecular forces in various compounds, comparing structures for examples like CO2, NH3, SH2. Draw Lewis structures to support assessments.