Chemical Bonding, Lewis Structures, and Molecular Geometry: Study Notes for CHM1032 (Ch. 10)

Chemical Bonding and Molecular Structure

Introduction to Chemical Bonding

Chemical bonding explains how atoms combine to form molecules and compounds. Understanding bonding theories is essential for predicting molecular structure and properties.

• Chemical Bond: The force that holds atoms together in a molecule or compound.

• Bonding Theories: Models such as Lewis theory, VSEPR, and molecular orbital theory help explain how and why atoms bond.

• Example: Water (H2O) forms via covalent bonds between hydrogen and oxygen atoms.

Lewis Structures

Drawing Lewis Structures

Lewis structures are diagrams that show the arrangement of valence electrons around atoms in a molecule. They help predict bonding and molecular geometry.

• Valence Electrons: Electrons in the outermost shell, involved in bonding.

• Steps to Draw Lewis Structures:

  1. Count total valence electrons for all atoms.

  2. Arrange atoms, usually with the least electronegative atom in the center.

  3. Connect atoms with single bonds (pairs of electrons).

  4. Distribute remaining electrons to complete octets (or duets for H).

  5. Use double or triple bonds if necessary to satisfy octet rule.

• Example: The Lewis structure for CO2 is: O=C=O

Counting Bonding and Lone Pair Electrons

• Bonding Electrons: Electrons shared between atoms (in bonds).

• Lone Pair Electrons: Electrons not involved in bonding, located on atoms.

• Example: In CO2, there are 8 bonding electrons (4 pairs in double bonds).

Resonance Structures

Some molecules cannot be represented by a single Lewis structure. Resonance structures are multiple valid Lewis structures for the same molecule, showing delocalized electrons.

• Resonance: The actual structure is a hybrid of all possible resonance forms.

• Symbol: Resonance structures are separated by a double-headed arrow (↔).

• Example: Ozone (O3) has two resonance structures.

Periodic Table and Electronegativity

Periodic Table Overview

The periodic table organizes elements by increasing atomic number and groups elements with similar chemical properties.

• Groups: Vertical columns, elements have similar valence electron configurations.

• Periods: Horizontal rows, elements have the same number of electron shells.

• Example: Alkali metals (Group 1) are highly reactive.

Electronegativity

Electronegativity is a measure of an atom's ability to attract electrons in a chemical bond.

• Trend: Electronegativity increases across a period (left to right) and decreases down a group.

• Highest Electronegativity: Found in the upper right of the periodic table (excluding noble gases), with fluorine being the most electronegative element.

• Definition: The ability of an atom to attract electrons within a covalent bond.

Molecular Geometry and VSEPR Theory

Electron and Molecular Geometries

The shape of a molecule is determined by the arrangement of electron groups around the central atom, as described by the Valence Shell Electron Pair Repulsion (VSEPR) theory.

• Electron Groups: Bonds and lone pairs around a central atom.

• VSEPR Theory: Electron groups repel each other, arranging themselves as far apart as possible.

• Common Geometries:

  • Linear: 180° bond angle

  • Trigonal planar: 120° bond angle

  • Tetrahedral: 109.5° bond angle

  • Trigonal bipyramidal: 90°, 120° bond angles

  • Octahedral: 90° bond angle

Table: Electron and Molecular Geometries

The following table summarizes electron group arrangements, molecular geometries, bond angles, and examples:

Electron Groups	Lone Pairs	Electron Geometry	Molecular Geometry	Bond Angle	Example
2	0	Linear	Linear	180°	CO2
3	0	Trigonal planar	Trigonal planar	120°	BF3
3	1	Trigonal planar	Bent	<120°	SO2
4	0	Tetrahedral	Tetrahedral	109.5°	CH4
4	1	Tetrahedral	Trigonal pyramidal	<109.5°	NH3
4	2	Tetrahedral	Bent	<109.5°	H2O
5	0	Trigonal bipyramidal	Trigonal bipyramidal	90°, 120°	PCl5
6	0	Octahedral	Octahedral	90°	SF6

Polarity and Molecular Shape

Polar Covalent Bonds and Molecules

A polar covalent bond occurs when electrons are shared unequally between atoms due to differences in electronegativity. Molecular polarity depends on both bond polarity and molecular geometry.

• Polar Bond: Bond with unequal sharing of electrons.

• Polar Molecule: Molecule with an uneven distribution of charge, often due to both polar bonds and asymmetrical shape.

• Example: H2O is a polar molecule with bent geometry.

Practice and Application

Lewis Structures for Ions and Compounds

• Lewis Structure for CN-: Shows 10 electrons (including the extra electron for the negative charge).

• Lewis Structure for N2: Triple bond between two nitrogen atoms, with 6 bonding electrons.

• Lewis Structure for RbI: Rb+ and I- ions, each with a full octet.

Compound Formation and Formulas

• Predicting Formulas: Use the charges of ions to determine the empirical formula of ionic compounds.

• Example: The formula for a compound formed between barium (Ba2+) and aluminum (Al3+) is Ba3Al2.

Additional info:

• Students should be familiar with using the periodic table to determine valence electrons and predict bonding.

• Understanding molecular geometry is essential for predicting physical and chemical properties.