Basic Concepts of Chemistry

Study Notes

Chemistry involves the study of matter, its properties, and the changes it undergoes.
Matter is divided into pure substances and mixtures based on chemical composition.

Pure substances: have a fixed composition and consistent properties; examples include distilled water and pure metals.
Mixtures: do not have fixed compositions; examples include paint and concrete.

Organic chemistry: focuses on carbon compounds.
Inorganic chemistry: deals with substances not covered by organic chemistry.
Physical chemistry: studies the principles and physical properties of matter.
Biological chemistry: examines the chemistry of living organisms.
Analytical chemistry: involves the identification and quantification of chemical substances.

Pure substances are categorized into:
- Elements: Cannot be broken down into simpler substances; classified as metals, nonmetals, and metalloids.
- Compounds: Combinations of two or more elements.

Metals: Lustrous, conductive, ductile, malleable; examples include gold and copper.
Non-metals: Dull, poor conductors, brittle; examples include nitrogen and iodine.
Metalloids: Have mixed characteristics; examples include silicon and arsenic.

Physical properties: Can be observed without changing composition (e.g., color, odor, melting point).
Chemical properties: Only observable during chemical reactions (e.g., flammability, reactivity).

Mass and Weight: Mass is constant and does not change, while weight varies due to gravity. SI unit for mass is the kilogram (kg).
Length: Measured in meters (m); small lengths can be expressed in nanometers (nm) and picometers (pm).
Volume: Amount of space occupied, measured in cubic meters (m³) or liters (L).

Law of Conservation of Mass: Mass cannot be created or destroyed (Lavoisier).
Law of Definite Proportions: Compounds have a consistent ratio of elements by mass (Proust).
Law of Multiple Proportions: Different compounds formed from the same elements will have mass ratios of small whole numbers (Dalton).
Gay-Lussac’s Law: Gases react in simple volume ratios at constant temperature and pressure.

Dalton's Atomic Theory: Matter consists of indivisible atoms, and atoms combine in fixed ratios to form compounds.
Atomic mass: Measured in atomic mass units (amu); 1 amu is one-twelfth the mass of a carbon-12 atom.

Average atomic mass is the weighted average based on isotopes and their abundances.
Example for carbon: average atomic mass is 12.011 u based on isotopes (^{12}C), (^{13}C), and (^{14}C).

Molecular mass: Total average atomic mass of all atoms in a molecule.
Formula mass: Sum of atomic masses in ionic compounds like NaCl, where discrete molecules do not exist.

Molecular mass of water (H₂O) = 2(1.0 u) + 16.0 u = 18.0 u.
Formula mass of NaCl = 23.0 u (Na) + 35.5 u (Cl) = 58.5 u.### Atomic and Molar Mass
Atomic mass is expressed in atomic mass units (u), while molar mass is the mass of one mole of a substance in grams per mole (g mol-1).
Examples of elemental atomic and molar masses include:
- Hydrogen (H): 1.0 u, 1.0 g mol-1
- Carbon (C): 12.0 u, 12.0 g mol-1
- Oxygen (O): 16.0 u, 16.0 g mol-1
For polyatomic substances, molar mass numerically equals molecular/formula mass in u.

Common polyatomic substances and their respective molar and molecular masses:
- Oxygen (O2): 32.0 u, 32.0 g mol-1
- Water (H2O): 18.0 u, 18.0 g mol-1
- Sodium Chloride (NaCl): 58.5 u, 58.5 g mol-1

Urea's molecular formula: NH2CONH2
Calculation of molecular mass:
- 2 Nitrogen (N): 2 x 14 u
- 1 Carbon (C): 1 x 12 u
- 4 Hydrogen (H): 4 x 1 u
- 1 Oxygen (O): 1 x 16 u
Total molecular mass = 60 u
Molar mass of urea = 60 g mol-1
Number of moles in 5.6 g of urea:
- Moles = mass / molar mass = 5.6 g / 60 g mol-1 = 0.0933 mol
Number of molecules using Avogadro's constant:
- Molecules = 0.0933 mol x 6.022 x 10^23 molecules/mol ≈ 5.618 x 10^22 molecules

For 52 moles of Argon (Ar):
- Atoms = 52 moles x 6.022 x 10^23 atoms/mol ≈ 313.144 x 10^23 atoms of Ar
For 52 u of Helium (He):
- Atomic mass of He = 4.0 u, thus 52 u corresponds to:
- Atoms = 52 / 4.0 ≈ 13 atoms of He
For 52 g of Helium:
- Molar mass of He = 4.0 g mol-1
- Moles = 52 g / 4.0 g mol-1 = 13 mol
- Atoms = 13 mol x 6.022 x 10^23 atoms/mol ≈ 78.286 x 10^23 atoms of He

Volume of ammonia (NH3) at STP = 67.2 dm3
Molar volume of a gas at STP = 22.4 dm3/mol
Number of moles of NH3:
- Moles = Volume / Molar volume = 67.2 dm3 / 22.4 dm3/mol = 3.0 mol
Number of molecules in 3.0 moles:
- Molecules = 3.0 mol x 6.022 x 10^23 molecules/mol ≈ 18.066 x 10^23 molecules

Moles of gases can be calculated using volume rather than mass.
Avogadro's law indicates that one mole of any gas occupies 22.4 dm3 at standard temperature (0°C) and pressure (1 atm), known as the molar volume of a gas.