Bonding, Molecular Geometry, and Chemical Quantities: GOB Chemistry Study Notes

Bonding and Molecular Structure

Vocabulary and Key Concepts

This section introduces essential terms and concepts related to chemical bonding, molecular geometry, and chemical quantities. Understanding these terms is foundational for predicting molecular properties and performing chemical calculations.

• Valence Electrons: Outer electrons responsible for bonding between atoms.

• Core Electrons: Inner electrons not involved in bonding.

• Lewis Dot Structures: 2D representations of compounds showing valence electrons as dots around element symbols.

• Octet Rule: Atoms tend to have 8 electrons in their valence shell (except H, He, B).

• VSEPR Theory: 3D representation of compounds; predicts molecular shapes based on electron pair repulsion.

• Electron Geometry (EG): Indicates how many regions of electron density (bonds and lone pairs) are attached to a central atom.

• Molecular Geometry (MG): Describes the shape of the molecule, considering only atoms (not lone pairs).

• Bond Polarity: Indicates difference in electronegativity between atoms in a bond.

• Molecular Polarity: Describes how equally electrons are shared in a molecule; determines if a molecule is polar or non-polar.

• Intermolecular Forces: Forces between molecules, including dipole-dipole, hydrogen bonding, and London dispersion forces.

• Dipole-Dipole: Unequal sharing of electrons between polar molecules.

• Hydrogen Bonding: Strong dipole-dipole interaction occurring only in molecules with N-H, O-H, or F-H bonds.

• London Dispersion Forces: Weak forces present in all compounds due to temporary shifts in electron density.

• Mole: Specific unit representing $6.022 \times 10^{23}$ particles (Avogadro's number).

• Avogadro's Number: $6.022 \times 10^{23}$

• Molar Mass: Sum of atomic masses from the periodic table for a compound, expressed in g/mol.

• Dimensional Analysis: Tool used to convert units in chemical calculations.

• Scientific Notation: Allows for small/large numbers to be written concisely (e.g., $6.022 \times 10^{23}$).

• Stoichiometry: Conversion between element/compound quantities utilizing subscripts and molar ratios.

Lewis Dot Structures and Molecular Geometry

Drawing Lewis Dot Structures

Lewis dot structures are used to represent the arrangement of valence electrons in molecules. They help predict molecular geometry and polarity.

• NH3 (Ammonia): N has 5 valence electrons, H has 1 each. Total: 8 electrons. EG: Tetrahedral (4 regions of electron density). MG: Trigonal pyramidal (3 atoms, 1 lone pair). Polarity: Polar.

• CO2 (Carbon Dioxide): C has 4, O has 6 each. Total: 16 electrons. EG: Linear. MG: Linear. Polarity: Non-polar.

• PCl3 (Phosphorus Trichloride): P has 5, Cl has 7 each. Total: 26 electrons.

• H2S (Hydrogen Sulfide): H has 1 each, S has 6. Total: 8 electrons. EG: Tetrahedral. MG: Bent. Polarity: Polar.

• CCl4 (Carbon Tetrachloride): C has 4, Cl has 7 each. Total: 32 electrons. EG: Tetrahedral. MG: Tetrahedral.

• BF3 (Boron Trifluoride): B has 3, F has 7 each. Total: 24 electrons. Polarity: Non-polar, all electrons are shared equally.

• PBr3 (Phosphorus Tribromide): P has 5, Br has 7 each. Total: 26 electrons. Polarity: Polar, lone pair on P makes electrons unequally shared.

• CF3Br (Bromotrifluoromethane): C has 4, F has 7 each, Br has 7. Total: 32 electrons. EG: Tetrahedral. MG: Tetrahedral.

Electron Geometry vs. Molecular Geometry

Electron geometry considers all regions of electron density (bonds and lone pairs), while molecular geometry considers only the arrangement of atoms.

• Electron Geometry (EG): Tells how many things (atoms and lone pairs) are attached to the central atom. Example: Tetrahedral (4 regions).

• Molecular Geometry (MG): Differentiates between electron lone pairs and elements bonded to the central atom. Example: Trigonal pyramidal (3 atoms, 1 lone pair).

Intermolecular Forces

Types of Intermolecular Forces

Intermolecular forces determine physical properties such as boiling point and solubility.

• Hydrogen Bonding: Occurs in molecules with N-H, O-H, or F-H bonds. Example: H2O, NH3.

• Dipole-Dipole: Occurs in polar molecules. Example: HBr, NaCl.

• London Dispersion: Present in all molecules, especially non-polar ones. Example: CH4.

Example Table: Intermolecular Forces in Selected Compounds

Compound	Intermolecular Force
H2O	Hydrogen bonding
HBr	Dipole-dipole
CH4	London dispersion
NH3	Hydrogen bonding
NaCl	Dipole-dipole

Chemical Quantities and Calculations

Mole Calculations and Avogadro's Number

The mole is a fundamental unit in chemistry for counting particles. Avogadro's number ($6.022 \times 10^{23}$) is used to convert between moles and number of particles.

• Number of Particles: $\text{Particles} = \text{Moles} \times 6.022 \times 10^{23}$

• Molar Mass: $\text{Molar Mass} = \sum \text{Atomic Masses}$ (from periodic table)

• Dimensional Analysis: Used to convert between grams, moles, and number of particles.

Example Calculations:

• How many particles in 2.5 moles of RbCl? $2.5 \text{ moles RbCl} \times 6.022 \times 10^{23} = 1.6 \times 10^{24}$ particles

• Convert 45°C to Kelvin: $T_K = T_C + 273$ $T_K = 45 + 273 = 318 \text{ K}$

• How many particles of Cu in 0.246 g CuCl3? $0.246 \text{ g CuCl}_3 \times \frac{1 \text{ mole}}{169.805 \text{ g}} \times \frac{6.022 \times 10^{23}}{1 \text{ mole}} = 8.72 \times 10^{20}$ particles

• How many grams are in 0.75 moles of H2O? $0.75 \text{ moles} \times 18.015 \text{ g/mole} = 13.511$ g

• How many molecules in 10.0 g H2O? $10.0 \text{ g} \times \frac{1 \text{ mole}}{18.015 \text{ g}} \times 6.022 \times 10^{23} = 3.34 \times 10^{23}$ molecules

• How many grams in 3.61 × 1023 molecules of CO2? $3.61 \times 10^{23} \text{ molecules} \times \frac{1 \text{ mole}}{6.022 \times 10^{23}} \times 44.009 \text{ g/mole} = 22.0$ g

• How many sodium ions in 25.0 g NaCl? $25.0 \text{ g} \times \frac{1 \text{ mole}}{58.443 \text{ g}} \times 6.022 \times 10^{23} = 2.58 \times 10^{23}$ ions

• How many grams in 5.00 × 1024 molecules of C2H6? $5.00 \times 10^{24} \text{ molecules} \times \frac{1 \text{ mole}}{6.022 \times 10^{23}} \times 30.07 \text{ g/mole} = 250$ g

Additional Info

• Shapes Affect Properties: The shape of a molecule affects how molecules interact (e.g., gases vs. solids) and influences physical properties such as boiling point and solubility. For example, H2O's bent shape leads to strong hydrogen bonding, while CO2's linear shape results in weak intermolecular forces.

• Polarity and Interactions: Polar molecules have stronger intermolecular forces (hydrogen bonding, dipole-dipole) compared to non-polar molecules (London dispersion).