Basic Concepts of Chemistry

Introduction to Chemistry

• 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 vs Mixtures

• 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.

Branches of Chemistry

• 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.

Classification of Pure 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.

Properties of Matter

• 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.

States of Matter

• Solid: Particles tightly packed in a structured arrangement.
• Liquid: Particles close together but can move.
• Gas: Particles are widely spaced.

Properties of Matter

• 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).

Measurement of Properties

• 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).

SI Units

• The International System of Units (SI) defines seven base units:
  • Length: metre (m)
  • Mass: kilogram (kg)
  • Time: second (s)
  • Electric Current: ampere (A)
  • Temperature: kelvin (K)
  • Amount of Substance: mole (mol)
  • Luminous Intensity: candela (cd)

Laws of Chemical Combination

• 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.

Atomic Theory and Atomic Mass

• 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

• 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).

Moles and Molar Mass

• A mole contains Avogadro's number (6.022 \times 10^{23}) particles.
• Molar mass: Mass in grams per mole corresponds to the atomic mass in amu.

Molecular Mass and Formula Mass

• 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.

Calculations Examples

• 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.

Polyatomic Substances

• 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 Example Calculation

• 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

Atom Calculation Examples

• 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

Ammonia Gas Calculation

• 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

Understanding Moles and Gases

• 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.