GenChem1 SHS - Quick Reference Notes

Analytical Chemistry: identifying and measuring components; qualitative and quantitative methods
Inorganic Chemistry: compounds without carbon (minerals, metals, salts)
Physical Chemistry: physical properties, energy transformations, states, reaction rates
Biochemistry
Organic Chemistry: carbon-containing compounds; basis of life
Geochemistry
Pharmaceutical Chemistry
Electrochemistry
Industrial Chemistry
Metallurgical Chemistry
Nuclear Chemistry
Photochemistry
Food Chemistry
Agricultural Chemistry
Polymer Chemistry
Colloidal Chemistry

Solid: definite shape and volume; particles tightly packed; vibrate in fixed positions
Liquid: indefinite shape; definite volume; particles close and can flow
Gas: no definite shape or volume; particles far apart and move randomly

Increasing temperature/pressure increases energy; overcomes attractive forces; phase changes (e.g., solid → liquid → gas)
Decreasing temperature/pressure reduces energy; phase changes reverse (gas → liquid → solid)

Property: distinguishing characteristic used to identify/describe a substance
Physical Properties: observed or measured without changing chemical identity
- Intensive: do not depend on amount (e.g., color, density, melting point, boiling point)
- Extensive: depend on amount (e.g., volume, mass)
Chemical Properties: describe how a substance changes to form a new substance
Changes in matter:
- Physical Changes: no chemical composition change (e.g., melting, dissolving, freezing)
- Chemical Changes: new substances formed (e.g., rusting, burning)

Mixture: physical combination of two or more pure substances; each retains identity
- Homogeneous Mixture (Solution): single visible phase; uniform properties (e.g., air, salt in water)
- Heterogeneous Mixture: visibly different phases (e.g., oil and water)
- Colloids: particles larger than solutions but smaller than suspensions (e.g., milk, fog)
- Suspensions: particles settle out over time (e.g., muddy water)
Pure Substance: single kind of matter; cannot be separated into other substances by physical means
- Elements: simplest form; cannot be chemically broken down
- Metals, Nonmetals, Metalloids
- Compounds: two or more elements chemically combined; can be decomposed by chemical means
- Acids, Bases, Salts
- Ionic Compounds vs Covalent (Molecular) Compounds

SI Units: standardized units for scientific measurements
Mass: amount of matter; measured with a balance
Volume: space that an object occupies; measured with a graduated cylinder or pipette
Reading a graduated cylinder: read at the bottom of the meniscus

Density: mass per unit volume; an intensive property
ρ = m/v ; m = ρ v (mass, density, volume)
Common density units: g/cm^3 or g/mL
Unit Conversions (Dimensional Analysis / Factor-Label Method): multiply by conversion factors to cancel units
Examples:
- 2000 g × (1 kg / 1000 g) = 2.0 kg
- 5 g × (1000 mg / 1 g) = 5000 mg
- 2 cm × (1 km / 100000 cm) = 0.00002 km
Ladder method: decimal movement between units; an alternative to factor-label

Precision: closeness of multiple measurements to each other
Accuracy: closeness of a measurement to the true value
Percent Error: $\%\text{Error} = \left|\frac{E - A}{A}\right| \times 100\%$
Scientific Notation: a \times 10^{b},\quad 1 \le a < 10,\quad b \in \mathbb{Z}
Exponent interpretation: positive exponent increases magnitude (rightward in standard form), negative exponent decreases magnitude (leftward)

Early Philosophical Ideas
- Democritus (~400 BC): matter consists of indivisible atoms (atomos); atoms are in constant motion
- Aristotle (~350 BC): matter is continuous; four elements (earth, air, fire, water)
Modern Atomic Theory (Models)
- Dalton (Early 1800s) – Billiard Ball Model
- Postulates: matter is composed of atoms; atoms of same element identical; atoms differ between elements; conservation of mass; atoms combine in simple whole-number ratios; chemical reactions rearrange atoms
- Law basis: Conservation of Mass; Definite Proportions; Multiple Proportions
- Thomson (1897) – Plum Pudding Model: electrons in a positively charged sphere
- Rutherford (1911) – Nuclear Model: gold foil; nucleus; mostly empty space; electrons orbit
- Bohr (1913) – Bohr Model (Shell Model): electrons in fixed energy levels; quantized energies; transitions via quanta
- Chadwick (1932) – Discovery of the neutron

Atomic Number (Z): number of protons; defines element
Mass Number (A): protons + neutrons
Isotopes: same Z, different A (different number of neutrons)
Molecules: two or more atoms chemically bonded; smallest unit of a compound
Ions: atoms/molecules with net charge
- Cation: positive (lost electrons)
- Anion: negative (gained electrons)

History: Dmitri Mendeleev (1869) organized elements by increasing atomic mass and grouped by similar properties; predicted undiscovered elements
Modern Periodic Table: arranged by increasing atomic number
- Periods: horizontal rows; same number of electron shells
- Groups/Families: vertical columns; same valence electrons; similar properties