Module 1: Structure of Matter (Atoms and Molecules)

Basic Definitions of Science and Matter

• Observation: The act of obtaining qualitative information through the senses.

• Qualitative Information: Non-numerical data describing qualities, such as color, texture, or smell (e.g., "the cauliflower is purple," "the metal handle was cold").

• Quantitative Information: Numerical data obtained through measurement (e.g., "I am sixteen years old," "the salt water boils at $106\,\text{!"C}$").

• Data: A collection of observations and measurements recorded during an investigation.

• Hypothesis: A tentative, testable explanation for an observation (e.g., "Concrete can be dissolved in boiling water").

• Experiment: A controlled procedure carried out to test a hypothesis.

• Theory: A well-substantiated explanation of some aspect of the natural world based on a body of facts that have been repeatedly confirmed (e.g., "Gases are composed of tiny billiard-like particles in constant motion which are colliding and transferring energy").

• Scientific Law: A statement based on repeated experimental observations that describes some aspects of the universe; it always applies under the same conditions (e.g., "In any chemical reaction, matter is not created or destroyed").

Review of Matter and Classification

• Atom: The smallest unit of an element that retains the properties of that element.

• Molecule: Two or more atoms chemically bonded together.

• Pure Substance: A form of matter that has a constant composition and properties that are consistent throughout; it can be an element or a compound.

• Element: A pure substance that cannot be broken down into simpler substances by chemical means; made up of only one type of atom.

• Compounds: Pure substances composed of two or more different elements chemically combined in fixed proportions. The properties of a compound are generally very different from the elements that compose it.

• Ion: An atom or molecule with a net electric charge due to the loss or gain of one or more electrons.

• Cation: A positively charged ion ($+$).

• Anion: A negatively charged ion ($-−$).

• Mixtures: Physical combinations of two or more substances where each substance retains its own chemical identity.

• Heterogeneous Mixture: A mixture where the components are not uniformly distributed (e.g., cookie batter, fruit salad).

• Homogeneous Mixture: A mixture where the components are uniformly distributed; also called a Solution (e.g., vitamin C dissolved in water, milk, paint).

• Physical Property: A characteristic of a substance that can be observed or measured without changing the identity of the substance (e.g., melting point, density).

• Physical Changes: Changes that affect the form of a chemical substance but not its chemical composition (e.g., snow melting, building a sandcastle).

• Chemical Changes: Processes where one or more substances are altered into one or more new and different substances (e.g., logs burning, digestion, rust, fireworks exploding).

A Brief History of Atomic Theory and Bohr Model

• Subatomic Particles:

  • Proton: Located in the nucleus; charge of $+1$; mass of approximately $1\,\text{amu}$.

  • Neutron: Located in the nucleus; neutral charge ($0$); mass of approximately $1\,\text{amu}$.

  • Electron: Located in shells orbiting the nucleus; charge of $-1$; mass is approximately $\frac{1}{1836}$ of a proton (usually ignored in atomic mass calculations).

• Atomic Number: The number of protons in the nucleus; defines the element.

• Atomic Mass: The sum of the number of protons and neutrons in an atom.

• Atomic Weight: The weighted average mass of all naturally occurring isotopes of an element.

• Isotopes: Atoms of the same element (same atomic number) with different numbers of neutrons (different atomic mass).

• Isotope Calculation Example: Chlorine is 76% Chlorine-35 and 24% Chlorine-37.

  • $\text{Atomic Weight} = (35 \times 0.76) + (37 \times 0.24) = 35.48\,\text{amu}$.

• The Rutherford Model Problem: According to classical physics, an electron orbiting a nucleus would constantly accelerate, radiate energy, and eventually spiral into the nucleus.

• Bohr Model:

  • Electrons exist in specific electron shells (energy levels) at fixed distances from the nucleus.

  • Valence Electrons: Electrons in the outermost shell.

  • Period number on the periodic table indicates the number of electronic shells.

  • Group number (for main group elements) indicates the number of valence electrons.

Introduction to Quantum Theory

• Quantum: The minimum amount of any physical entity involved in an interaction.

• Photon: A particle representing a quantum of light or other electromagnetic radiation.

• Wave-Particle Duality: Light and matter (protons, electrons) exhibit both wave-like and particle-like properties.

• Electronic Transition: When an electron moves between energy levels.

  • Moving to a higher shell: Absorbs a photon (energy).

  • Moving to a lower shell: Emits a photon (energy).

• De Broglie Equation: $\lambda = \frac{h}{mv}$

  • $\lambda$ is wavelength, $h$ is Planck's constant, $m$ is mass, and $v$ is velocity.

  • Objects with smaller mass and lower velocity act more like waves (longer wavelength).

• Quantum Mechanical Structure:

  • Electron Shell: Designated by principal quantum number $n$.

  • Sub-shell: Found within shells (s, p, d, f).

  • Orbital: A region of space within a sub-shell where there is a high probability of finding an electron. One orbital can hold a maximum of 2 electrons.

  • Shapes: s-orbitals are spherical; p-orbitals are dumbbell-shaped.

Electron Configurations and Ionization

• Electron Configuration Rules:

  • Electrons fill the lowest energy orbitals first (Aufbau principle).

  • Core Notation: A shorthand using the noble gas from the previous period (e.g., Sn core is [Kr]).

• Exceptions: Certain elements (like Cr or Cu) fill d-subshells differently to achieve stability.

• Ions:

  • Isoelectronic: Species that have the same electron configuration (e.g., $Mg^{2+}$ and $Ne$).

  • Transition metal cations often lose s-electrons before d-electrons.

Periodic Trends

• Atomic Radii:

  • Increases down a Group (more shells).

  • Decreases across a Period (higher effective nuclear charge pulls electrons closer).

• Ionic Radii:

  • Cations are smaller than their parent atoms (loss of electron/shell).

  • Anions are larger than their parent atoms (electron-electron repulsion).

• Ionization Energy (IE): The energy required to remove an electron from an atom.

  • Large atoms have low IE (electrons are further from nucleus and easier to remove).

  • IE increases across a period and decreases down a group.

  • Metals generally have lower IE than non-metals.

• Electrostatic Forces: The forces of attraction/repulsion between charged particles.

Electronegativity and Chemical Bonding

• Electronegativity: The tendency of an atom to attract shared electrons in a chemical bond. Fluorine ($F$) is the most electronegative element.

• Combining Capacity (Valence): The number of bonds an atom can form.

• Types of Bonds:

  • Nonpolar Covalent: Electrons are shared equally (minimal difference in electronegativity).

  • Polar Covalent: Electrons are shared unequally; they move closer to the more electronegative atom.

  • Ionic Bond: Electrons are transferred from one atom to another (large difference in electronegativity).

Lewis Structures and VSEPR Theory

• Lewis Structure: A diagram showing the bonding between atoms and the lone pairs of electrons.

• Octet Rule: Atoms tend to form bonds to achieve eight valence electrons.

  • Exceptions: Hydrogen ($2$), Boron ($6$), and Expanded Octets (elements in period 3 or below using d-orbitals).

• VSEPR Theory (Valence Shell Electron Pair Repulsion): Electron pairs repel each other, dictating the 3D shape of molecules.

• Coordination Number (CN): The sum of bonded atoms and lone pairs on a central atom.

• Geometries:

  • $CN = 2$: Linear ($180^{\circ}$).

  • $CN = 3$: Trigonal Planar ($120^{\circ}$).

  • $CN = 4$: Tetrahedral ($109.5^{\circ}$).

  • $CN = 5$: Trigonal Bipyramidal ($90^{\circ}/120^{\circ}$).

  • $CN = 6$: Octahedral ($90^{\circ}$).

• Molecular Shapes: Refers only to the position of atoms (not lone pairs), e.g., Bent, Trigonal Pyramidal, See-saw, T-shape, Square Planar.

Polarity and Intermolecular Forces (IMF)

• Bond Dipole: Indicated by an arrow ($\rightarrow$) pointing toward the more electronegative atom. The length represents the magnitude of the charge difference.

• Molecular Polarity: Determined by the vector sum of bond dipoles. Symmetrical molecules may have polar bonds but be non-polar overall.

• Intermolecular Forces vs. Intramolecular Forces:

  • Intramolecular: Forces within a molecule (chemical bonds).

  • Intermolecular: Forces between molecules.

• Types of IMF (from weakest to strongest):

  • London Dispersion Forces: Present in all molecules; caused by temporary dipoles.

  • Dipole-Dipole Interactions: In polar molecules.

  • Hydrogen Bonding: A strong dipole-dipole force occurring when H is bonded to N, O, or F.

• Physical Properties: Stronger IMF leads to higher boiling points. "Like dissolves like" (polar substances dissolve in polar solvents).