A. Atomic Architecture
I. Conservation of Matter: Atoms exist and are indestructible and incredibly small. In their chemical reactions with each other, pairings might be changed, bonds broken and formed, but in the end, atoms can never be created or destroyed. Elements can be characterized by the similar masses of their atoms. One consequence of this is that all rearrangements of atoms conserve matter. The first attempts to categorize matter led to tables in which the elements were organized first by mass from lightest to heaviest and then by chemical behavior that was observed to occur in cyclic or periodic ways, thus the first periodic tables. These early ideas of mass being the most important property led to great efforts to devise a relative mass scale for atoms. Today, we can measure relative masses very accurately by mass spectroscopy. Later we learned that while all atoms in their elemental states are neutral, they are made up of charged particles: negatively charged electrons which appear to be distributed throughout the volume of the atom and can be easily knocked off or added to from an ion, and positively charged protons, discovered to be densely packed together into a nucleus in the center of the atom. With the discovery of neutrons, the modern periodic table can be understood as being organized by atomic number, or the number of protons in the nucleus, and not by mass. In any chemical reaction, it is still true that mass is conserved.
Tools for Understanding Chemical Formulas and Reactions: Mass Spectroscopy, Molecular Bookkeeping or Stoichiometry: percent composition, mole fractions, gravimetric precipitations, balancing equations.
II. Equations of State: Gases, the most simple collection of atoms: Gases are both remarkable and simple in that many of their properties, such as pressure (P), volume (V), and temperature (T), can be seen experimentally to be independent of the elemental identity of the gas. The relationships between P, V, and T are therefore called Equations of State (the state being the gaseous state in this case). The seeming simplicity of the gaseous state challenged scientists to see if they could develop a model from “first principles” that might predict the molecular origin of these properties and derive the relationships between the properties that we observe experimentally. One such model: the kinetic molecular theory of gases, starts with a microscopic picture of a moving gas particle with kinetic energy and generalizes to a system of particles. The model seeks to relate the pressure of the system of particles to the average velocity. A remarkable moment is when we realize that the average velocity can be determined by measuring the temperature of the gas. The energy of an idealized system of gaseous particles is found to depend only on the temperature of a gas, and not on its identity. These idealized gases (ideal gases) are made real through the introduction of scorrection factors to account for particle size and interactions.
Tools for Equations of State: Introduction of a new “absolute” temperature scale known as Kelvin; Charles’, Boyles’ and Raoult’s Laws, the ideal gas law, Dalton’s Law of Partial Pressures.
Tools for Kinetic Molecular Theory of Gases: System Properties and Statistics of large systems, Boltzmann velocity distributions, Collision rates, Laws of Effusion and Diffusion.
III. Quantum Mechanical Atom: To think quantum mechanically we first need to know about light. Light can be characterized as an electromagnetic wave with a wavelength, frequency, and velocity. Light exhibits both wave-like (diffraction) and particle-like (photoelectric effect) properties. Pure atomic gases emit light when they are heated, and the pattern of the emitted light is a fingerprint for the elemental identity of the gas. Neils Bohr offered one explanation for the origin of the emitted light. His planetary model of the atom posited the existence of energy levels for the electrons. While this explained a lot, it didn’t explain everything. Moving further and deeper, the ideas of de Broglie, Schrödinger, and Heisenberg were used to formulate a quantum mechanical model for atomic structure in with electrons are organized in probabilistic “orbitals.” Armed with this interpretation, we go back and explore orbital energies, ground and excited state electron configurations. Quantum mechanics shows us how electrons can be characterized by quantum numbers, and how these in turn dictate the chemical properties of the elements and the shape of the Periodic Table.
Tools for Understanding Atomic Structure: Absorption and Emission Spectroscopy of Atoms, Energy Level Diagrams, Wave Functions
Tools for Understanding Electron Configurations of the Elements: Pauli Exclusion Principle, Hund’s Rule, Aufbau principle
B. Molecular Architecture
I. Bonding: Atoms form molecules when the valence electrons on one atom interact with the valence electrons on another atom in such a way as to lower the total energy of the system. Depending on the energies, these molecular assemblies can range from stable, long-lived covalent species to transient van der Waals molecules. In this section, we explore guidelines for predicting the shape, size, and polarity of these molecules based on the periodic properties of the elements.
Ionic Bonding – is what we see in salts and other crystals, and has its origin in charge interactions known as coulombic interactions. Neutral atoms can either give up or lose electrons to become charged ions. The ionic bond is a consequence of a positive ion’s attraction for a negative ion. Salts and crystals are important materials in the environment and form the substance of the inorganic materials that surround us. Incredibly strong, these interactions can often be magically overcome by dissolving the salt crystal in water.
Covalent Bonding - is what we observe in most compounds in which electrons are not transferred from one atom to the other, but rather shared between the two. Covalent molecules predominate in organic materials such as those that make up living organisms, and so these are incredibly important in biology and medicine. These bonds are explored through properties such as bond length, bond strength, and dipoles.
Non-covalent Bonding – is of the lowest energy, and yet these interactions are incredibly important. They have as their origin the ability of one atom or molecule to transiently polarize a neighbor, thus creating weak transient interactions which when added together can often affect such things as a boiling point or the shape of a protein molecule in solution.
Coordinate Covalent Bonding - Metal ions interact with small molecules and ions so as to form beautiful complexes of a regular shape and geometry. The surrounding ions (ligands) produce an effect on the valence orbital energies of the metal in a predictable manner described in ligand field theory. The symmetries and properties such as magnetism and color of these complexes are explained by this theory.
Tools for Understanding Bonding: Lewis Structures, VSEPR, Hybridization, Resonance, Molecular Orbital Theory, Periodic Properties, Energy Level Diagrams, Spectroscopy, Redox potentials