Chapter 4: Molecules and Compounds

Ionic and Covalent Bonds

Atoms combine to form compounds through two primary types of chemical bonds: ionic bonds and covalent bonds. Understanding the nature of these bonds is fundamental to predicting compound properties.

• Ionic Bonds: Formed when electrons are transferred from a metal atom to a nonmetal atom, resulting in oppositely charged ions that attract each other.

• Covalent Bonds: Formed when two nonmetal atoms share one or more pairs of electrons.

• Formation: Ionic bonds typically occur between elements with large differences in electronegativity, while covalent bonds occur between elements with similar electronegativities.

• Example: Sodium chloride (NaCl) forms via ionic bonding; water (H2O) forms via covalent bonding.

Lewis Dot Structures

Lewis dot structures are diagrams that represent the valence electrons of atoms within a molecule. They are essential for visualizing bonding and lone pairs.

• Octet Rule: Most atoms tend to have eight electrons in their valence shell (exceptions include hydrogen, beryllium, boron, etc.).

• Drawing Steps:

  1. Count total valence electrons.

  2. Arrange atoms (central atom is usually least electronegative).

  3. Connect atoms with single bonds.

  4. Distribute remaining electrons to satisfy octet rule.

• Example: The Lewis structure for methane (CH4) shows carbon at the center with four single bonds to hydrogen atoms.

Naming and Writing Formulas for Compounds

Correctly naming and writing formulas for ionic and covalent compounds is crucial for communication in chemistry.

• Ionic Compounds: Name the cation (metal) first, then the anion (nonmetal with -ide ending). For transition metals, indicate charge with Roman numerals.

• Covalent Compounds: Use prefixes (mono-, di-, tri-, etc.) to indicate the number of each atom. The more electronegative element is named last with an -ide ending.

• Example: CO2 is carbon dioxide; Na2SO4 is sodium sulfate.

Polyatomic Ions

Polyatomic ions are ions composed of two or more atoms covalently bonded, carrying a net charge.

• Common Examples: Nitrate (NO3-), sulfate (SO42-), ammonium (NH4+), carbonate (CO32-).

• Recognition: Memorize names, formulas, and charges for common polyatomic ions.

Mass Percent of an Element in a Compound

The mass percent of an element in a compound expresses the fraction of the compound's mass contributed by that element.

• Formula:

• Example: In H2O, the mass percent of hydrogen is calculated using the molar masses of H and O.

Empirical and Molecular Formulas from Analysis

Elemental and combustion analysis can be used to determine a compound's empirical formula (simplest whole-number ratio) and molecular formula (actual number of atoms).

• Steps:

  1. Convert mass or percent composition to moles for each element.

  2. Divide by the smallest number of moles to get the simplest ratio.

  3. Empirical formula is written from these ratios.

  4. Molecular formula is determined by comparing empirical formula mass to molar mass.

• Example: A compound with 40% C, 6.7% H, and 53.3% O by mass has an empirical formula of CH2O.

Chapter 5: Molecular Shapes

Polarity and Electronegativity

Electronegativity is the ability of an atom to attract shared electrons. The difference in electronegativity between atoms determines bond polarity.

• Nonpolar Covalent Bonds: Electrons are shared equally (electronegativity difference < 0.5).

• Polar Covalent Bonds: Electrons are shared unequally (electronegativity difference between 0.5 and 2.0).

• Example: The O-H bond in water is polar; the H-H bond in H2 is nonpolar.

Lewis Structures with Octet Rule Exceptions

Some molecules have atoms that do not follow the octet rule, such as those with incomplete octets (e.g., BF3), expanded octets (e.g., SF6), or odd numbers of electrons (radicals).

• Incomplete Octet: Elements like B and Be may have fewer than 8 electrons.

• Expanded Octet: Elements in period 3 or higher can have more than 8 electrons.

• Example: The Lewis structure for PF5 shows phosphorus with 10 valence electrons.

Resonance Structures and Formal Charge

Resonance structures are alternative Lewis structures for a molecule or ion that differ only in the placement of electrons. Formal charge helps identify the most stable resonance structure.

• Formal Charge Formula:

• Most Stable Structure: The structure with formal charges closest to zero and negative charges on the most electronegative atoms is preferred.

• Example: The carbonate ion (CO32-) has three resonance structures with delocalized negative charge.

VSEPR Theory: Molecular and Electron Geometry

Valence Shell Electron Pair Repulsion (VSEPR) theory predicts the 3D shape of molecules based on repulsion between electron groups around a central atom.

• Electron Geometry: Arrangement of all electron groups (bonding and lone pairs).

• Molecular Geometry: Arrangement of only the atoms (ignoring lone pairs).

• Bond Angles: Determined by the number of electron groups and presence of lone pairs.

• Example: Methane (CH4) is tetrahedral; ammonia (NH3) is trigonal pyramidal due to one lone pair.

Chapter 6: Valence Bond Theory

Hybridization of Central Atoms

Valence bond theory explains molecular bonding using the concept of hybridization, where atomic orbitals mix to form new, equivalent hybrid orbitals.

• Electron Groups and Hybridization:

  • 2 electron groups: sp hybridization

  • 3 electron groups: sp2 hybridization

  • 4 electron groups: sp3 hybridization

  • 5 electron groups: sp3d hybridization

  • 6 electron groups: sp3d2 hybridization

• Example: In methane (CH4), carbon is sp3 hybridized; in ethene (C2H4), each carbon is sp2 hybridized.