General Chemistry Study Guide: Molecular Structure, Nomenclature, and Calculations

Molecular Structure

Lewis Structures and Bonding

Understanding molecular structure is fundamental in general chemistry. Lewis structures are diagrams that represent the bonding between atoms and the lone pairs of electrons in a molecule.

• Lewis Structures: Draw accurate Lewis structures to show single, double, and triple bonds between atoms. Each line represents a pair of shared electrons.

• Electron Transfer Diagrams: For ionic compounds, use electron transfer diagrams to show the movement of electrons from one atom to another, resulting in the formation of ions.

• Example: The Lewis structure for water (H2O) shows two single bonds between oxygen and hydrogen, with two lone pairs on oxygen.

VSEPR Theory and Molecular Geometry

The Valence Shell Electron Pair Repulsion (VSEPR) Theory is used to predict the geometry of molecules based on the repulsion between electron pairs around a central atom.

• VSEPR Theory: Electron pairs (bonding and lone pairs) arrange themselves to minimize repulsion, determining the molecular shape.

• Common Geometries: Linear, trigonal planar, tetrahedral, trigonal bipyramidal, octahedral.

• Example: Methane (CH4) has a tetrahedral geometry.

Sigma and Pi Bonds

Covalent bonds can be classified as sigma (σ) or pi (π) bonds. Sigma bonds are the first bonds formed between two atoms, while pi bonds are additional bonds in double or triple bonds.

• Number of Sigma and Pi Bonds: Single bond = 1 sigma; double bond = 1 sigma + 1 pi; triple bond = 1 sigma + 2 pi.

• Example: Ethylene (C2H4) has a double bond between carbons: 1 sigma and 1 pi bond.

Hybridization

Hybridization describes the mixing of atomic orbitals to form new hybrid orbitals suitable for bonding.

• Types of Hybridization: sp (linear), sp2 (trigonal planar), sp3 (tetrahedral).

• Example: In methane (CH4), carbon is sp3 hybridized.

Chemical Nomenclature

Naming Compounds

Chemical nomenclature is the system for naming chemical substances. It varies for ionic compounds, covalent compounds, and acids.

• Ionic Compounds: Name the cation first, then the anion. For example, NaCl is sodium chloride.

• Covalent Compounds: Use prefixes to indicate the number of atoms (mono-, di-, tri-, etc.). For example, CO2 is carbon dioxide.

• Acids: If the anion ends in -ide, the acid name begins with 'hydro-' and ends with '-ic acid' (e.g., HCl: hydrochloric acid). If the anion ends in -ate or -ite, use '-ic acid' or '-ous acid' respectively (e.g., H2SO4: sulfuric acid).

Calculations

Molarity

Molarity (M) is a measure of concentration, defined as moles of solute per liter of solution.

• Formula:

$M = \frac{\text{moles of solute}}{\text{liters of solution}}$

• Example: Dissolving 0.5 moles of NaCl in 1 liter of water gives a 0.5 M solution.

Dilution Calculations

To dilute a solution, use the relationship between initial and final concentrations and volumes.

• Formula:

$M_1V_1 = M_2V_2$

• Example: To make 250 mL of 0.1 M solution from a 1.0 M stock, use $V_1 = \frac{M_2V_2}{M_1}$.

Mass and Volume from Molarity

Calculate the mass of solute needed for a solution of given molarity and volume.

• Formula:

$\text{mass} = M \times V \times \text{molar mass}$

• Example: To prepare 500 mL of 0.2 M NaCl, calculate moles: $0.2 \times 0.5 = 0.1$ mol, then mass: $0.1 \times 58.44 = 5.844$ g.

Empirical Formulas

The empirical formula represents the simplest whole-number ratio of atoms in a compound.

• Steps:

  1. Convert mass of each element to moles.

  2. Divide by the smallest number of moles.

  3. Round to nearest whole number to get the ratio.

• Example: A compound with 40 g C and 6.7 g H: $\frac{40}{12} = 3.33$ mol C, $\frac{6.7}{1} = 6.7$ mol H. Ratio: 1 C : 2 H, empirical formula CH2.

Reference Materials and Exam Preparation

Materials to Bring

• Number 2 pencil

• Scientific calculator

Materials Provided

• Periodic table

*Additional info: Academic context and examples have been added to expand on the brief points in the original notes.*