Study Guide — BC 212 Lecture 1: Building Blocks of Molecules

Part 1: Introduction & Definitions

This section introduces the foundational concepts and terminology used in chemistry to describe matter and its components.

• Matter: Anything with mass and volume.

• Mass: Amount of matter present in a substance.

• Pure substance: Material with uniform composition throughout (e.g., distilled water).

• Mixture: Material with variable composition (e.g., air, salt water).

• Atom: Smallest particle of an element that retains its properties.

• Molecule: Smallest unit of a substance that retains chemical properties and can exist independently.

• Element: Pure substance consisting of only one type of atom.

• Compound: Pure substance made of two or more different atom types chemically bonded.

• Main biological elements: C, H, O, N (plus smaller amounts of others).

Part 2: Subatomic Particles & Isotopes

This section covers the structure of atoms, including subatomic particles and isotopes.

• Proton: Positively charged particle (+1), located in the nucleus; defines the element.

• Neutron: Neutral particle, located in the nucleus; adds mass but does not affect charge.

• Atomic number (Z): Number of protons in the nucleus.

• Mass number (A): Total number of protons and neutrons.

• Isotopes: Atoms of the same element (same Z) with different numbers of neutrons.

• Atomic weight: Weighted average mass based on isotope abundance.

Example: Hydrogen isotopes: protium (H), deuterium (H), tritium (H).

Part 3: Electronic Structure & Periodic Table

This section explains how electrons are arranged in atoms and how the periodic table organizes elements.

• Rows = periods: Indicate the number of electron shells.

• Columns = groups: Indicate the number of valence electrons (for main group elements).

• Valence electrons: Electrons in the outermost shell; determine bonding and reactivity.

• s, p, d, f blocks: Refer to subshells where electrons are found.

Example:

• Group 1A: 1 valence electron.

• Group 7A: 7 valence electrons.

• Group 8A (noble gases): Full valence shell, stable.

Part 4: Periodic Properties

This section describes trends in the periodic table that affect element properties.

• Metallic character: Increases to the left and down the table.

• Ionization energy: Energy required to remove an electron; increases to the right and up.

• Electronegativity: Tendency to attract electrons; increases to the right and up.

• Reactivity: Most reactive metals are at the bottom left; most reactive nonmetals are at the top right.

Part 5: Octet Rule & Bonding

This section introduces the octet rule and types of chemical bonding.

• Octet rule: Atoms tend to have 8 valence electrons (except H, He — 2 electrons).

• Bonding types:

  • Ionic: Transfer of electrons (metal + nonmetal).

  • Covalent: Sharing of electrons (nonmetal + nonmetal).

• Noble gases: Already stable (full octet).

Example: NaCl forms by Na+ giving an electron to Cl-. F2 forms by sharing a pair of electrons.

Part 6: Ions & Ionic Bonding

This section covers the formation of ions and ionic compounds.

• Cation: Positively charged ion (lost electrons).

• Anion: Negatively charged ion (gained electrons).

• Binary ionic compounds: Formed by balancing charges between cations and anions.

• Common polyatomic ions: NH4+ (ammonium), SO42- (sulfate), HCO3- (bicarbonate), CO32- (carbonate), PO43- (phosphate), CN- (cyanide).

Study Guide — BC 212 Lecture 2: Molecular Structure and Interactions with Water

Part 1: Covalent Bonding & Lewis Structures

This section explains how atoms share electrons to form covalent bonds and how to represent molecules using Lewis structures.

• Bonding capacity: Determined by group (valence electrons).

  • Group 14: 4 bonds (C).

  • Group 15: 3 bonds + 1 lone pair (N).

  • Group 16: 2 bonds + 2 lone pairs (O, S).

  • Group 17: 1 bond + 3 lone pairs (F, Cl).

  • Group 18: 0 bonds (noble gases).

• Lewis structures: Show bonding pairs (lines) and lone pairs (dots).

• Rules:

  • Full octet (or duet for H).

  • Total valence = sum of all atoms' charges.

  • Negative charge = extra lone pair(s).

• Exceptions: Period 3 (e.g., S) can exceed octet (e.g., PO43-, SO42-).

Part 2: 3D Shapes of Molecules (VSEPR)

This section describes how electron groups determine the three-dimensional shapes of molecules.

• Electron groups repel: Lead to 3D geometry.

• 2 groups = Linear, 180° (CO2).

• 3 groups = Trigonal planar, 120° (H2CO).

• 4 groups = Tetrahedral, 109.5° (CH4).

• Lone pairs: Take up space and distort geometry (e.g., NH3 = trigonal pyramidal, H2O = bent).

Part 3: Polarity

This section explains how differences in electronegativity and molecular geometry affect polarity.

• Bond polarity: Depends on electronegativity difference ().

  • <0.5 = nonpolar covalent.

  • 0.5–1.9 = polar covalent.

  • >1.9 = ionic.

• Dipole moment: Charge separation in a bond.

• Molecular polarity: Depends on geometry.

  • Symmetrical molecules (CO2, CH4, CF4) = nonpolar.

Part 4: Intermolecular Forces (IMFs)

This section covers the types and effects of forces between molecules.

• Types of IMFs:

  • London dispersion: Weakest; present in all molecules, especially nonpolar.

  • Dipole–dipole: Present in polar molecules.

  • Hydrogen bonds: Strongest IMF; special dipole–dipole interaction.

• Phase changes:

  • Endothermic: Requires heat (melting, evaporation, sublimation).

  • Exothermic: Releases heat (freezing, condensation, deposition).

• Strength trend: ionic > H-bond > dipole–dipole > London.

Part 5: Hydrogen Bonding

This section details the nature and importance of hydrogen bonds.

• H-bond donor: H covalently bound to O, N, or F.

• H-bond acceptor: Lone pair on O, N, or F.

• Directionality: H-bonds are stronger and more directional than dipole–dipole interactions.

• Examples: H2O, NH3, alcohols, DNA base pairing.

Part 6: Solutions & Properties

This section explains solution chemistry and related properties.

• Solution: Homogeneous mixture of solute + solvent.

• Electrolytes: Substances that produce ions in solution.

  • Strong: NaCl, KBr.

  • Weak: Acetic acid.

  • Nonelectrolyte: Glucose, sucrose.

• Osmolarity: Total particle concentration in solution.

  • NaCl → 2 ions per unit.

  • Al2(SO4)3 → 5 ions per unit.

• Osmosis: Solvent moves through semipermeable membrane.

• Dialysis: Solvent + small solutes pass through.

• Tonicity:

  • Isotonic: No change in cells.

  • Hypotonic: Cells swell.

  • Hypertonic: Cells shrink.

Table: Common Polyatomic Ions

Key Equations

• Atomic number:

• Mass number:

• Electronegativity difference:

Practice Questions & Answers

Practice questions are provided for each topic, with answers for self-assessment.

• Example Question: Which element is NOT found in significant amounts in the body: C, N, O, F?

• Answer: Fluorine (F).

• Example Question: How many total electrons in neutral oxygen?

• Answer: 8 electrons.

• Example Question: Which process is exothermic: condensation, evaporation, melting?

• Answer: Condensation.

Additional info: These notes are structured to cover the essential topics for GOB Chemistry, including atomic structure, periodic trends, chemical bonding, molecular geometry, intermolecular forces, and solution chemistry, with examples and practice questions for exam preparation.