Chapter 5: Chemical Bonding I and Molecular Geometry

Chapter 5: Chemical Bonding and Molecular Geometry

Introduction

Chemical bonding is fundamental to understanding chemistry, explaining why atoms bond to form molecules and predicting molecular properties. This chapter covers the nature of chemical bonds, how to represent molecules, and how molecular shape affects chemical behavior.

Classifying Bonds

Types of Chemical Bonds

• Pure Covalent Bonds: Electrons are shared equally between atoms.

• Polar Covalent Bonds: Electrons are shared unequally, creating a dipole moment.

• Ionic Bonds: Electrons are transferred from one atom to another, forming ions.

Example: H2 (pure covalent), HCl (polar covalent), NaCl (ionic).

Electronegativity and Bond Polarity

Electronegativity

• Definition: The ability of an atom in a chemical bond to attract electrons to itself.

• Periodic Trend: Electronegativity increases from left to right across a row and decreases down a column in the periodic table.

• Values: Fluorine is the most electronegative element ($ext{EN} = 4.0$), while cesium and francium are among the least.

Equation: No universal equation, but values are tabulated (Pauling scale).

Bond Polarity

• Polar Bonds: Formed when atoms of different electronegativities share electrons unequally.

• Dipole Moment ($oldsymbol{ ext{μ}}$): Measures the separation of charge in a bond.

Equation: $\mu = q \times r$

Percent Ionic Character

• Calculated as: $\text{Percent ionic character} = \frac{\text{measured dipole moment}}{\text{dipole moment if electrons were completely transferred}} \times 100\%$

Writing Lewis Structures

Steps for Molecular Compounds

1. Write the correct skeletal structure for the molecule.

2. Count the total number of valence electrons.

3. Distribute electrons as bonding pairs between atoms.

4. Distribute remaining electrons as lone pairs to terminal atoms, then to the central atom.

5. If necessary, form double or triple bonds to satisfy the octet rule.

6. Check that the total number of electrons matches the number calculated in step 2.

Example: For CO2, the structure should have 16 valence electrons, with double bonds between C and each O.

Polyatomic Ions

• Adjust the total number of electrons based on the ion's charge.

• Enclose the Lewis structure in brackets and indicate the charge.

Example: For NH4+, add one less electron and indicate the positive charge.

Resonance and Formal Charge

Resonance Structures

• Some molecules have multiple valid Lewis structures; the actual structure is a hybrid.

• Resonance structures have the same skeletal structure but different electron arrangements.

Example: SO2 has two resonance structures with different placements of double bonds.

Formal Charge

• Definition: A fictitious charge assigned to each atom in a Lewis structure to help determine the most stable arrangement.

• Calculation: $\text{Formal charge} = \text{valence electrons} - (\text{nonbonding electrons} + \frac{1}{2} \times \text{bonding electrons})$

• The best Lewis structure minimizes formal charges and places negative charges on the most electronegative atoms.

Exceptions to the Octet Rule

Odd-Electron Species

• Molecules or ions with an odd number of electrons (e.g., NO).

Incomplete Octets

• Some elements (e.g., B, Be) form stable compounds with fewer than eight electrons.

Expanded Octets

• Elements in the third period and beyond can have more than eight electrons (e.g., SF6).

Bond Energies and Bond Lengths

Bond Energy

• Definition: The energy required to break 1 mole of a bond in the gas phase.

• Trend: Multiple bonds are stronger than single bonds.

Example: C–H bond: 414 kJ/mol; C≡C bond: 839 kJ/mol.

Bond Length

• Definition: The average distance between the nuclei of two bonded atoms.

• Trend: Bond length decreases as bond order increases (single > double > triple).

Bond Type	Bond Length (pm)
H–H	74
H–C	110
C–C (single)	154
C=C (double)	134
C≡C (triple)	120
O–H	96

VSEPR Theory: Molecular Shapes

Valence Shell Electron Pair Repulsion (VSEPR) Theory

• Electron groups (lone pairs, single, double, triple bonds) arrange themselves to minimize repulsion.

• The geometry of a molecule is determined by the number of electron groups around the central atom.

Basic Molecular Geometries

Electron Groups	Geometry	Bond Angles
2	Linear	180°
3	Trigonal Planar	120°
4	Tetrahedral	109.5°
5	Trigonal Bipyramidal	90°, 120°
6	Octahedral	90°

Effect of Lone Pairs

• Lone pairs occupy more space than bonding pairs, reducing bond angles.

• Shapes such as bent, trigonal pyramidal, and seesaw result from lone pairs.

Predicting Molecular Geometry

Step-by-Step Procedure

1. Draw the Lewis structure.

2. Count the number of electron groups around the central atom.

3. Assign electron group geometry and molecular geometry.

4. Adjust for lone pairs to determine the final shape.

Example: NH3 has four electron groups (three bonds, one lone pair) → tetrahedral electron geometry, trigonal pyramidal molecular geometry.

Molecular Polarity

Determining Polarity

• Polarity depends on both bond polarity and molecular geometry.

• Symmetrical molecules (e.g., CO2) are nonpolar even if bonds are polar.

• Asymmetrical molecules (e.g., H2O) are polar.

Vector Addition and Dipole Moments

• Dipole moments are vector quantities; the overall molecular dipole is the vector sum of individual bond dipoles.

• For molecules with more than two polar bonds, use vector addition to determine net polarity.

Key Terms and Concepts

• Electronegativity: Ability of an atom to attract electrons in a bond.

• Dipole Moment ($\mu$): Separation of charge in a molecule.

• Formal Charge: Fictitious charge assigned to atoms in Lewis structures.

• Resonance: Multiple valid Lewis structures for a molecule.

• Octet Rule: Atoms tend to have eight electrons in their valence shell.

• VSEPR Theory: Electron groups arrange to minimize repulsion.

• Molecular Geometry: The arrangement of atoms in space.

Equations and Relationships

• Dipole Moment: $\mu = q \times r$

• Percent Ionic Character: $\text{Percent ionic character} = \frac{\text{measured dipole moment}}{\text{dipole moment if electrons were completely transferred}} \times 100\%$

• Formal Charge: $\text{Formal charge} = \text{valence electrons} - (\text{nonbonding electrons} + \frac{1}{2} \times \text{bonding electrons})$

Summary Table: Molecular Geometries

Electron Groups	Lone Pairs	Geometry	Bond Angles
2	0	Linear	180°
3	0	Trigonal Planar	120°
3	1	Bent	~120°
4	0	Tetrahedral	109.5°
4	1	Trigonal Pyramidal	~107°
4	2	Bent	~104.5°
5	0	Trigonal Bipyramidal	90°, 120°
6	0	Octahedral	90°

Additional info:

• These notes expand on the original file by providing full definitions, equations, and context for each topic, ensuring a self-contained study guide suitable for exam preparation.