Chemical Bonding, Molecular Structure, and Compounds: Study Guide (Chapters 4–6)

Chapter 4: Molecules and Compounds

Diatomic Elements

Some elements exist naturally as diatomic molecules, meaning they are composed of two atoms bonded together.

• Diatomic elements: Hydrogen (H2), Nitrogen (N2), Oxygen (O2), Fluorine (F2), Chlorine (Cl2), Bromine (Br2), Iodine (I2).

• Mnemonic: "Have No Fear Of Ice Cold Beer" helps remember the diatomic elements.

Octet Rule and Element Families

The octet rule states that atoms tend to gain, lose, or share electrons to achieve a full set of eight valence electrons, similar to noble gases.

• Alkali metals: Lose one electron to form cations.

• Halogens: Gain one electron to form anions.

• Noble gases: Already have a full octet; generally unreactive.

• Other families: React to achieve octet via sharing or transfer of electrons.

Noble Gas Configurations in Ions

When forming ions, atoms achieve electron configurations similar to noble gases.

• Cations: Lose electrons to match the previous noble gas configuration.

• Anions: Gain electrons to match the next noble gas configuration.

• Example: Na+ has the same electron configuration as Ne.

Ionic vs. Molecular (Covalent) Substances

Compounds can be classified as ionic or molecular based on the types of elements involved and the nature of their bonding.

• Ionic compounds: Formed from metals and nonmetals; involve transfer of electrons.

• Molecular (covalent) compounds: Formed from nonmetals; involve sharing of electrons.

• Example: NaCl (ionic), H2O (molecular).

Covalent vs. Ionic Bonds

Covalent bonds involve sharing of electrons between atoms, while ionic bonds involve transfer of electrons from one atom to another.

• Covalent: Usually between nonmetals.

• Ionic: Usually between a metal and a nonmetal.

Atoms vs. Ions; Molecules vs. Polyatomic Ions

• Atoms: Neutral species with equal numbers of protons and electrons.

• Ions: Charged species formed by gaining or losing electrons.

• Molecules: Neutral group of atoms bonded covalently.

• Polyatomic ions: Group of covalently bonded atoms with a net charge.

Formulas and Names

Compounds and ions are named and represented by formulas according to specific rules.

• Common cations: Na+, K+, Ca2+

• Common anions: Cl-, O2-, SO42-

• Polyatomic ions: NH4+ (ammonium), NO3- (nitrate)

• Ionic compounds: NaCl, CaSO4

• Binary molecular compounds: CO2 (carbon dioxide), SF6 (sulfur hexafluoride)

• Ternary compounds: Compounds with three different elements, e.g., NaNO3

Percent Composition and Empirical Formula

Percent composition is the mass percent of each element in a compound. The empirical formula is the simplest whole-number ratio of elements.

• Percent composition formula:

$\text{Percent composition} = \frac{\text{mass of element in 1 mol compound}}{\text{molar mass of compound}} \times 100\%$

• Empirical formula: Determined from percent composition or combustion analysis.

Molecular Formula from Empirical Formula and Molar Mass

The molecular formula is a whole-number multiple of the empirical formula.

• Formula:

$\text{Molecular formula} = (\text{Empirical formula})_n$

$n = \frac{\text{Molar mass}}{\text{Empirical formula mass}}$

Combustion Analysis

Combustion analysis is used to determine the empirical formula of compounds containing C, H, and another element.

• Process: Compound is burned; masses of CO2 and H2O produced are measured.

• Calculation: Use masses to determine moles of C and H; subtract from total mass to find other element.

Chapter 5: Chemical Bonding I

Bond Ionicity and Molecular Character

Bonds can be classified based on their ionic or molecular (covalent) character.

• More ionic: Large difference in electronegativity between atoms.

• More molecular: Small or no difference in electronegativity.

Bond Types: Ionic, Polar Covalent, Nonpolar Covalent

Bonds are classified by the difference in electronegativity between the atoms.

• Ionic: Electronegativity difference > 2.0

• Polar covalent: Electronegativity difference between 0.5 and 2.0

• Nonpolar covalent: Electronegativity difference < 0.5

• Periodic table: Use to estimate electronegativity differences.

Degree of Polarity

The degree of polarity depends on the difference in electronegativity; greater difference means more polar bond.

• Example: H–F is more polar than H–Cl.

Lewis Diagrams

Lewis diagrams show the arrangement of electrons in molecules and ions.

• Multiple bonding: Double or triple bonds may be needed to satisfy the octet rule.

• Example: O2 has a double bond.

Formal Charge

Formal charge helps determine the most stable Lewis structure.

• Formula:

$\text{Formal charge} = \text{Valence electrons} - (\text{Nonbonding electrons} + \frac{1}{2} \text{Bonding electrons})$

Multiple Lewis Diagrams and Stability

Some molecules can have more than one valid Lewis structure; the most stable has the lowest formal charges.

• Resonance: When multiple structures are possible, the actual structure is a hybrid.

Resonance Structures and Bond Order

Resonance occurs when electrons are delocalized over multiple atoms.

• Bond order formula:

$\text{Bond order} = \frac{\text{Number of bonds}}{\text{Number of resonance structures}}$

• Example: In NO3-, bond order is 1.33.

Octet Rule Exceptions

Some species violate the octet rule:

• Expanded octet: Atoms with more than eight electrons (e.g., SF6).

• Radical: Odd number of electrons (e.g., NO).

• Deficient octet: Less than eight electrons (e.g., BF3).

VSEPR Theory and Molecular Geometry

Valence Shell Electron Pair Repulsion (VSEPR) theory predicts the geometry and bond angles in molecules.

• Linear: 180°

• Trigonal planar: 120°

• Tetrahedral: 109.5°

• Example: CH4 is tetrahedral.

Bond Length and Bond Dissociation Energy

Bond length is the distance between nuclei; bond dissociation energy is the energy required to break a bond.

• Shorter bonds: Higher bond energy.

• Multiple bonds: Shorter and stronger than single bonds.

Molecular Polarity

Molecular polarity depends on bond polarity and molecular geometry.

• Polar molecules: Have a net dipole moment.

• Nonpolar molecules: Symmetrical geometry cancels dipoles.

Chapter 6: Chemical Bonding II

Atomic Orbitals in Bond Formation

Bonds form by overlap of atomic orbitals; each atom uses specific orbitals to form bonds.

• s, p, d orbitals: Used in bonding depending on the atom.

• Example: Carbon uses 2s and 2p orbitals.

Sigma and Pi Bonds

Sigma (σ) bonds are formed by head-on overlap; pi (π) bonds are formed by side-on overlap.

• Single bond: One sigma bond.

• Double bond: One sigma and one pi bond.

• Triple bond: One sigma and two pi bonds.

Hybrid Orbitals and Spatial Orientation

Hybridization explains the observed geometry of molecules.

• sp: Linear, 180°

• sp2: Trigonal planar, 120°

• sp3: Tetrahedral, 109.5°

Bonding in Polyatomic Molecules

Polyatomic molecules use hybrid orbitals and various atomic orbitals to form bonds.

• Example: In CH4, carbon uses sp3 hybrids.

Orbital Overlap Types

Each bond in a molecule is formed by a specific type of orbital overlap.

• Sigma bonds: Direct overlap.

• Pi bonds: Side-on overlap.

• Number of each: Count based on bond types.

Valence Bond Theory vs. Molecular Orbital Theory

Two main theories explain chemical bonding:

• Valence Bond Theory: Bonds form by overlap of atomic orbitals; explains hybridization.

• Molecular Orbital Theory: Atomic orbitals combine to form molecular orbitals delocalized over the molecule.

• Main differences: Valence Bond Theory focuses on localized bonds; Molecular Orbital Theory explains delocalization and magnetic properties.

Theory	Bonding Description	Key Features
Valence Bond Theory	Overlap of atomic orbitals	Localized bonds, hybridization
Molecular Orbital Theory	Combination of atomic orbitals into molecular orbitals	Delocalized electrons, explains magnetism