General Chemistry Module 1

Matter and Its Classification

• Matter can exist in four states: solid, liquid, gas, and plasma/ionized gas.
• Solids have a definite shape and volume, liquids assume the container shape, and gases have an indefinite shape and volume.
• Plasma is the most abundant state of matter, with positively charged ions and negatively charged free electrons.

Laws of Chemical Combinations

• Law of Definite Proportions: a chemical compound always contains the same proportion of elements by mass.
• Law of Multiple Proportions: when two elements form more than one compound, the masses of one element that combine with a fixed mass of the other element are in a simple whole-number ratio.

Atomic Structure

• Democritus: matter is composed of indivisible atoms.
• John Dalton: atoms are indivisible, indestructible, and have a specific mass and properties.
• J.J. Thomson: electrons are negatively charged particles in a positively charged framework.
• Ernest Rutherford: atoms have a nucleus with positively charged protons and neutrons, and negatively charged electrons orbiting the nucleus.
• Niels Bohr: electrons occupy specific energy levels (orbitals) around the nucleus.

Matter Classification

• Pure substances: elements (simplest form of a substance) and compounds (two or more elements chemically united).
• Mixtures: two or more substances where individual substances retain their identities.

Phase Changes

• Melting: solid to liquid (fusion or liquefaction).
• Freezing: liquid to solid.
• Evaporation: liquid to gas.
• Condensation: gas to liquid.
• Sublimation: solid to gas.
• Deposition: gas to solid.

Fundamental Chemistry Laws

• Law of Conservation of Mass: matter is neither created nor destroyed in a chemical reaction.
• Law of Constant Proportions: a chemical compound always contains the same proportion of elements by mass.

Atomic Properties

• Atomic number: the number of protons in an atom's nucleus.
• Atomic mass: the sum of protons and neutrons in an atom's nucleus.
• Electrons: negatively charged particles orbiting the nucleus.
• Neutrons: particles with no charge in the nucleus.

Chemical Bonds

• Molecule: an aggregate of two or more atoms in a definite arrangement held together by chemical bonds.
• Ions: particles with a net positive or negative charge.
• Empirical formula: the simplest whole number ratio of atoms in a molecule.

Intermolecular Forces

• Van der Waals forces: weak forces between molecules.
• Hydrogen bonding: a strong intermolecular force between molecules with hydrogen atoms bonded to electronegative atoms.
• Dipole-dipole forces: forces between molecules with a permanent dipole moment.

Chemical Reactions

• Synthesis: a reaction in which two or more substances combine to form a new compound.
• Decomposition: a reaction in which a compound breaks down into two or more simpler substances.
• Single Displacement: a reaction in which one element displaces another element from a compound.
• Double Displacement: a reaction in which two compounds exchange partners.

Mole Relationships

• Avogadro's number: 6.022 x 10^23 atoms or molecules per mole.
• Mole: a unit of measurement equal to 6.022 x 10^23 atoms or molecules.
• Molarity: the number of moles of solute per liter of solution.
• Molality: the number of moles of solute per kilogram of solvent.
• Normality: the number of equivalents of solute per liter of solution.### Ionic Compounds
• Pb(NO3)4 is an example of an ionic compound, also known as Plumbic nitrate or Lead(IV) nitrate.
• Monovalent ions:
  • Group 1 (H, Li, Na, K, Ag) have a +1 charge.
  • Group 2 (Be, Mg, Ca, Sr, Ba, Zn, Cd) have a +2 charge.
  • Group 6A (Oxide, Sulfide) have a -2 charge.
  • Group 7A (Fluoride, Chloride, Bromide, Iodide) have a -1 charge.

Aufbau Principle

• Atoms may be built by progressive filling of energy levels, with lower energy levels being occupied first.
• The order of filling is: s, p, d, f.

Quantum Theories

• Pauli's Exclusion Principle: No two electrons will have the same set of quantum numbers.
• Heisenberg's Uncertainty Principle: It is impossible to accurately determine both the position and momentum of a particle at the same time.
• Hund's Rule: Orbitals are filled up singly before pairing up, with the most stable arrangement being the one with the greatest number of parallel spins.

Gas Laws

• Boyle's Law: P₁V₁ = P₂V₂
• Charles' Law: V₁ / T₁ = V₂ / T₂
• Gay-Lussac's Law: P₁ / T₁ = P₂ / T₂
• Combined Gas Law: P₁V₁ / T₁ = P₂V₂ / T₂
• Ideal Gas Law: PV = nRT

Quantum Numbers

• Principal Quantum Number (n): main energy level, size of orbital, and distance from the nucleus.
• Azimuthal Quantum Number (ℓ): angular momentum, shape of orbital, and subshell.
• Magnetic Quantum Number (mℓ): orientation of orbital in space.
• Magnetic Spin (ms): magnetic moment, spin, and rotation.

Gas Laws Applications

• Avogadro's Law: equal volumes of different gases have the same number of moles at STP.
• Dalton's Law of Partial Pressures: total pressure in a mixture of non-interacting gases is equal to the sum of the partial pressures of each gas.
• Graham's Law of Effusion: rate of effusion is inversely proportional to the square root of the density of the gas.
• Henry's Law of Gas Solubility: the pressure of a gas is proportional to its solubility in a liquid.

Thermodynamics

• Temperature: expressed in Kelvin (K), Celsius (°C), or Fahrenheit (°F).
• Internal Energy (U): total energy of a system.
• Enthalpy (H): total energy of a system, including heat.
• Entropy (S): measure of disorder or randomness in a system.
• Gibb's Free Energy (G): a measure of the energy available to do work.

Solutions

• Colligative Properties: dependent on the amount of solute present in the solution.
• Vapor Pressure Lowering: Raoult's Law states that the vapor pressure of a solution is dependent on the amount of nonvolatile solute added.
• Boiling Point Elevation: the boiling point of a solution is increased by the addition of a nonvolatile solute.
• Freezing Point Depression: the freezing point of a solution is decreased by the addition of a nonvolatile solute.
• Osmotic Pressure: the pressure needed to stop osmosis.

Chemical Kinetics

• Reaction Rate: the change in concentration of a reactant or product over time.
• Rate Law: expresses the relationship between the rate of reaction and the concentration of reactants.
• Factors Affecting Reaction Rate: nature of reactants, concentration of reactants, surface area, temperature, and catalysts.
• Reaction Rate Theories: Collision Theory, Transition Theory, and Activation Energy.

Chemical Equilibrium

• Law of Mass Action: the reaction rate is proportional to the product of the concentrations of the reactants to the power of their coefficients.

• Equilibrium Constant (K): a measure of the ratio of the concentrations of products to reactants at equilibrium.### Equilibrium and Le Chatelier's Principle

• Equilibrium: Keq = a/b (at a given temperature, the ratio of concentration of products to reactants)

• Keq > 1: favors product formation (to the right or forward reaction)

• Keq < 1: favors reactant formation (to the left or backward or reverse reaction)

• Le Chatelier's Principle: if an external stress is applied to a system at equilibrium, the system adjusts to partially offset the stress as it reaches a new equilibrium

• External stressors: concentration, pressure, and volume (for gases only), temperature, and the presence of a catalyst

Acid and Base Theories

• Arrhenius Theory: acids liberate H+ ions, bases liberate OH- ions
• Bronsted-Lowry Theory: acids donate H+ ions, bases accept H+ ions
• Lewis Theory: acids are electron pair acceptors, bases are electron pair donors

Ionic Equilibria

• Ion product constant (Q) is computed based on initial concentration
• Q < Ksp: unsaturated, Q = Ksp: saturated, Q > Ksp: supersaturated

Solubility Product Constant (Ksp)

• Ksp is a measure of the maximum amount of solute that can dissolve in a solvent
• Solubility (g/L) is the number of grams of solute dissolved in 1L of saturated solution
• Molar solubility (mol/L) is the number of moles of solute dissolved in 1L of saturated solution

Electrochemistry

• Study of the production of electricity from energy released during spontaneous and nonspontaneous chemical reactions
• Spontaneous: voltaic cells/galvanic cells (e.g., REDOX reaction, electrons migrate from anode to cathode)
• Nonspontaneous: electrolytic cells (e.g., electric current is applied to remove e- and transfer to another cell)

Periodic Table

• Antoine Lavoisier: first extensive list of elements (~ 33)
• Metals vs Nonmetals: oxides, reducing/oxidizing agents, conductivity, malleability, ductility, metallic luster, and state at room temperature (RT)
• Periods/Horizontal rows: 7, Groups/Vertical columns: 18, Actinides and Lanthanides: inner transition elements (f block)

Periodic Trends

• Ionization energy: energy needed to remove outermost electron in a neutral atom (increases from left to right and bottom to top)
• Electron affinity: energy given off when a neutral atom gains an extra electron (increases from left to right and bottom to top)
• Electronegativity: ability of an atom to attract an electron pair to itself (increases from left to right and bottom to top)

Radioactivity

• Spontaneous emission of particles/ionizing radiation by unstable nuclei of heavier elements (p+-to-n0 ratio, atomic number 92 and above: transuranic elements)
• Units: Curie (Ci), Becquerel (Bq), R.E.M. (roentgen equivalent in man), Rad/gray
• Types of radioactive emissions: alpha (α), beta (β), and gamma (γ)