Chemistry Basics: Atoms and Chemical Formulas

Atom

• An atom is indivisible and the smallest unit of matter.
• Subatomic particles include:
  • Proton (+) discovered by Ernest Rutherford.
  • Electron (−) discovered by Joseph John Thomson.
  • Neutron discovered by James Chadwick.
• The nucleus contains most of the atom's mass, with atomic mass calculated as the sum of protons and neutrons.
• The atomic number represents the number of protons, which equals the number of electrons in a neutral atom.
• Isotopes are variants of the same element with differing masses.

Chemical Formula

Types

• Empirical Formula represents the simplest ratio of elements.
  • Examples:
    • H₂O₂ simplifies to HO.
    • C₄H₁₀ simplifies to C₂H₅.
    • C₆H₁₂O₆ simplifies to CH₂O.
• Molecular Formula shows the exact number of atoms.
  • Examples:
    • Mg₃N₂
    • Ca(OH)₂

Branches of Chemistry

• Analytical Chemistry: Focused on the analysis of substances.
• Biochemistry: Studies chemical processes in living organisms.
• Organic Chemistry: Deals primarily with carbon-containing compounds.
• Inorganic Chemistry: Covers compounds that are not organic.
• Physical Chemistry: Combines principles of physics with chemistry.

Acids

• Binary Acids contain hydrogen and examples include:
  • HBr (hydrobromic acid)
  • HF (hydrofluoric acid)
  • HI (hydroiodic acid)
• Ternary Acids contain hydrogen and oxygen; examples include:
  • H₃PO₄ (phosphoric acid)
  • H₃PO₃ (phosphorous acid)

Molar Mass

• Molar mass is computed as the sum of atomic weights multiplied by their respective subscripts.
• Example calculations:
  • For HBr:
    • H: 1 × 1 = 1
    • Br: 1 × 80 = 80
    • Total = 81 g/mol
  • For H₂PO₄:
    • H: 2 × 1 = 2
    • P: 1 × 31 = 31
    • O: 4 × 16 = 64
    • Total = 97 g/mol

Percentage Composition

• Percentage of each element can be determined using:
  • % of element = (number of atoms × atomic mass) / molar mass × 100

Structural Formula

• Represents the arrangement of atoms within molecules; examples include:
  • H₂O depicted as H-O-H
  • C₄H₁₀ depicted as H-C-C-C-C-H

Atoms and Electrons

• Atoms may gain or lose electrons, leading to oxidation states.
  • Cation: Positively charged (loses electrons).
  • Anion: Negatively charged (gains electrons).

Example Compounds

Chemical Applications

• To determine empirical formulas from percentage composition, convert percentages to moles, find mole ratios, and derive the simplest formula.
  • Example compound with 28.82% Mg, 14.24% C, and 56.93% O results in empirical formula MgCO₃.

Molecular Formula Example

• For diborane:
  • Calculate grams of boron and hydrogen.
  • Find empirical formula BH₃.
  • Molar mass calculation leads to molecular formula B₂H₆.

Chemical Equations

• Basic format: Reactants → Products.
• Example of word equation: Aluminum + Sulfuric acid → Hydrogen + Aluminum sulfate.
• Example of skeletal equation: ( Al_{(s)} + H_2 SO_4 \longrightarrow H_2 + Al_2 (SO_4)_3 ).

Types of Reactions

• Synthesis: A + B → AB.
• Decomposition: AB → A + B.
• Single Replacement: A + BC → AC + B.
• Double Replacement: AB + CD → AD + BC.

Naming Compounds

• Binary Ionic Compounds end in -ide (e.g., NaF - sodium fluoride).
• Compounds with metals that have varying oxidation states use Stock or Traditional nomenclature (e.g., Iron (II) chloride, ferrous chloride).

Balancing Chemical Equations

• Ensure the number of each type of atom is the same on both sides of the equation.
• Example:
  • ( BaCl_2 + 2Na(OH){2} → Ba(OH){2} + 2NaCl ) balances Ba, Cl, Na, O, and H.
  • ( C_2H_4 + 3O_2 → 2CO_2 + 2H_2O ) balances C, H, and O.