General Chemistry 1 — Review Notes: Properties, Atomic Models, Periodic Table, Nomenclature, Reactions

Properties of Matter, Classification of Matter, and Simple Separation Techniques

  • 1.1 Properties of Matter

    • Matter ➔ Has mass and occupies space.
    • Chemical Change ➔ Change in the substance’s composition (examples: lighting a candle; digesting food).
    • Physical Change ➔ Change of physical state, form, or appearance (examples: ice melting; grinding coffee beans).
    • Chemical Properties ➔ Ability of a substance to change into one or more substances (examples: corrosiveness of acid; iron rusting).
    • Physical Properties ➔ A physical characteristic of a substance that can be observed without changing its composition (examples: state of matter; malleability of aluminum).

  • 1.2 Classification of Matter

    • Matter is classified into pure substances (elements and compounds) and mixtures (homogeneous and heterogeneous).
    • A. Pure substances
    • An element or compound made up of one type of particle.
    • Have a fixed composition and definite properties.
    • a. Elements ➔ Pure substances that cannot be broken down (examples: Al = Aluminum; Cl = Chlorine).
    • b. Compounds ➔ Composed of two or more elements chemically combined in fixed ratios (examples: CH₄ = Methane; NaCl = Sodium Chloride).
    • B. Mixtures
    • Composed of two or more substances physically combined, no fixed ratio.
    • a. Homogeneous mixtures are uniform (examples: Sugar in tea; Saltwater).
    • b. Heterogeneous mixtures are not uniform (examples: Oil and water; Salad).

  • 1.3 Simple Separation Techniques

    • 1. Filtration ➔ Uses a porous barrier to separate solids from liquids based on particle size.
    • 2. Distillation ➔ Uses heat to separate mixtures based on differences in boiling points.
    • 3. Crystallization ➔ Formation of solids from an impure mixture.
    • 4. Sublimation ➔ Solids turn into vapor, bypassing the liquid phase.
    • 5. Chromatography ➔ Separates substances based on solubility and affinity to the mobile and stationary phases.
    • 6. Evaporation ➔ Removes solvent by vaporization to separate solute.
    • 7. Mechanical Picking ➔ Manual separation based on physical properties.
    • 8. Magnetic Separation ➔ Uses magnets to separate ferromagnetic substances.
    • 9. Decantation ➔ Separates liquid from solid by gently pouring off the top layer; based on density difference.

The Development of Atomic Model, the Periodic Table, and Nomenclature

  • 2.1 The Development of Atomic Model
    • A. Development of the Atomic Model
    • Leucippus and Democritus proposed that matter is composed of tiny indivisible particles named atomos (Greek for indivisible).
    • 1. John Dalton ➔ Proposed the particle nature of matter; father of atomic theory.
    • a. Dalton’s Atomic Theory
      • 1. All matter is composed of tiny, indivisible particles called atoms.
      • 2. Each element has unique atoms, different from atoms of other elements.
      • 3. Atoms are neither created nor destroyed.
      • 4. Atoms of different elements combine with each other in certain whole-number proportions to form compounds.
    • b. Revisions of Dalton’s Atomic Theory
      • 1. Atoms are not indestructible; they contain smaller particles.
      • 2. Atoms of one element may differ in mass; they are identical in some basic respects.
    • c. Dalton’s Model ➔ Universally accepted as our current view of matter.
      • Laws associated with atomic theory:
      • 1. Law of Conservation of Mass: mass is conserved in ordinary chemical reactions.
      • 2. Law of Constant Composition: compounds contain the same elements in the same proportion by mass.
      • 3. Law of Multiple Proportions: the ratio of masses of the second element with a fixed mass of the first is a simple whole-number ratio.
    • 2. Plum Pudding Model (Joseph John Thomson) ➔ Atom as a massive positively charged blob embedded with negatively charged electrons.
    • 3. Nuclear Model (Ernest Rutherford) ➔ Atom has a dense, positively charged nucleus surrounded by electrons.
    • 4. Planetary Model / Bohr’s Model (Niels Bohr) ➔ Electrons move around the nucleus in orbits like planets around the sun.
    • 5. Quantum (Mechanical) Model (Erwin Schrödinger) ➔ Electrons move in waves around the electron cloud.
    • Subatomic Particles
    • Electron ➔ negatively charged particle (discovered by J.J. Thomson, 1897).
    • Proton ➔ positively charged particle (discovered by Ernest Rutherford, 1917).
    • Neutron ➔ neutral particle (discovered by James Chadwick, 1932).
    • Isotopic Notation by Henry Moseley
    • Atomic Number (Z) ➔ number of protons.
    • Atomic Mass (A) ➔ number of protons plus neutrons.
    • Atomic Number = number of protons = number of electrons.
    • Atomic Mass = number of protons + number of neutrons.
    • Number of neutrons = A − Z.
    • 2.2 The Modern Periodic Table
    • A. Dmitri Mendeleev is known as the father of the modern periodic table.
    • B. The periodic law: Elements are arranged according to atomic number and organized into groups with similar properties.
    • 1. Chemical Formula ➔ A shorthand symbol for compounds; uses element symbols and subscripts to show proportions.
    • 2. Empirical Formula ➔ Presents the simplest ratio of atoms in a compound.
    • 3. Molecular Formula ➔ Presents the exact number of atoms in a single molecule.
      • Example: Glucose
      • Molecular Formula: C6H{12}O_6
      • Empirical Formula: CH_2O
      • Example: Benzene
      • Molecular Formula: C6H6
      • Empirical Formula: CH
    • 2.3 NOMENCLATURE: Chemical Naming & Formula Writing
    • A. Metal-Nonmetal Binary (Ionic) Compounds
      • Cation is listed before anion; the cation-anion ratio must give a net charge of zero.
      • Formula Writing ➔ Use Criss-Cross rule: Criss-cross the charge of each ion to determine subscripts of the opposite ion. Reduce to lowest terms if possible.
      • Example: Aluminum Oxide ➔ Al forms a 3+ cation; Oxygen forms a 2− anion ➔ 2 Al³⁺ with 3 O²⁻ → empirical formula: ext{Al}2 ext{O}3
      • Example: Beryllium Sulfide ➔ Be²⁺ and S²⁻ → BeS (simplest whole-number ratio).
      • Formula Naming
      • Metals with Fixed Charges: Metal + Nonmetal + ide (e.g., Aluminum Oxide → Aluminum Oxide).
      • Metals with Two or More Charges (Roman numeral system)
        • IUPAC Name: Metal (Roman Numeral Charge) Ion
        • Classical System: -ous for lower charge, -ic for higher charge
      • Examples (Table 2.3.1)
      • K⁺ and Br⁻ ➔ KBr (Potassium Bromide)
      • Zn²⁺ and O²⁻ ➔ ZnO (Zinc Oxide)
      • Binary Ionic Compounds with Multiple Charges (Table 2.3.2)
      • CuCl (Lower Charge Cl) ➔ Copper(I) Chloride; Cuprous Chloride
      • CuCl₂ (Higher Charge Cl) ➔ Copper(II) Chloride; Cupric Chloride
      • Sb₂S₃ (Lower Charge S) ➔ Antimony(III) Sulfide; Antimonous Sulfide
      • Sb₂S₅ (Higher Charge S) ➔ Antimony(V) Sulfide; Antimonic Sulfide
    • B. Compounds with Polyatomic Ions
      • Polyatomic ions ➔ Ions containing more than one atom covalently bonded.
      • Oxyanion ➔ Polyatomic ions containing at least one oxygen.
      • Formula Writing ➔ Similar to ionic compounds; if the oxyanion has a subscript, enclose the oxyanion in parentheses when needed (e.g., Mg²⁺ with ClO⁻ → Mg(ClO)₂).
      • Naming with oxyanions depends on the number of oxygens:
      • 1 O: hypo__ite
      • 2 O: ite
      • 3 O: ate
      • 4 O: per___ate
      • Examples: K⁺ with NO₂⁻ ➔ KNO₂ naming follows metal + polyatomic ion rules.
      • Oxyacids: formed by hydrogen with polyatomic ions.
      • Naming: acid name derives from the anion root; -ite becomes -ous, -ate becomes -ic, followed by 'acid'.
      • Example: H₃PO₄ → Phosphoric Acid; PO₄³⁻ → Phosphate.
      • Example: Mg²⁺ with ClO⁻ (chlorite) → Mg(ClO)₂.
    • C. Nonmetal-Nonmetal Binary (Molecular) Compounds
      • Rules:
      • Write the less electronegative element first; more electronegative element second.
      • Suffix -ide on the second element.
      • Use Greek prefixes (mono-, di-, tri-,) to indicate the number of atoms; omit mono- on the first element if only one.
      • Examples (Table 2.3.3): HF → Hydrogen Fluoride; N₂O₄ → Dinitrogen Tetroxide; CO₂ → Carbon Dioxide; Br₂F₅ → Dibromine Pentafluoride.
    • D. Acids
      • Pure state naming treats hydrogen like a metal with one charge; add hydro- prefix to the anion root; change -ide to -ic and add 'acid'.
      • Polyatomic ions ending in -ide follow the binary acid naming rule.
      • Examples (Table 2.3.4/2.3.5): HF (gas) → Hydrogen Fluoride; Hydrofluoric Acid (aqueous). HCl → Hydrochloric Acid; HClO → Hypochlorous Acid; HIO₃ → Iodic Acid; H₃PO₄ → Phosphoric Acid; PO₄³⁻ → Phosphate; H₃PO₃ → Phosphorous Acid; PO₃³⁻ → Phosphite; etc.
    • E. Bases
      • Bases contain hydroxide (OH⁻). Naming: state the cation first, then add hydroxide.
      • Examples (Table 2.3.6): Na⁺ OH⁻ → NaOH (Sodium hydroxide); Al³⁺ → Al(OH)₃ (Aluminum hydroxide); NH₄⁺ → NH₄OH (Ammonium hydroxide); K⁺ → KOH (Potassium hydroxide).

CHEMICAL REACTIONS AND EQUATIONS

  • 3.1 Chemical Reactions

    • A process that results in the production of at least one substance from a chemical change.
    • In a chemical reaction, a substance is converted into another substance (or substances).
    • The Law of Conservation of Mass states that atoms are neither created nor destroyed during a chemical reaction; they are merely rearranged (Lavoisier).

  • 3.2 Chemical Equations

    • Formulas and symbols to describe changes that have occurred in the reaction.
    • Example: When hydrogen gas reacts with oxygen gas to form water: 2H2(g) + O2(g)
      ightarrow 2H_2O(l)
    • Reactants ➔ left of the arrow; starting substances.
    • Products ➔ right of the arrow; produced substances.
    • Coefficients ➔ numbers beside formulas indicating relative numbers of molecules.
    • Symbols:
    • → or ⇌ ➔ separates reactants from products; reads as yields, produces, forms, liberates.
    • + ➔ separates reactants or products from each other.
    • ↑ ➔ formation of a gaseous product.
    • ↓ ➔ presence of a precipitate.
    • (s), (l), (g), (aq) ➔ state of matter; solid, liquid, gas, aqueous (in solution).
    • Δ ➔ heating; reaction is heated.

  • 3.3 Writing Chemical Equations

    • Criss-Cross Rule ➔ To determine subscripts by criss-crossing the charges on ions; ensures zero net charge.
    • Gases ➔ Those at room temperature that are diatomic: Have No Fear Of Ice Cold Beer mnemonic for H, F, O, Cl, Br, N, I; write as ext{H}2, ext{N}2, ext{F}2, ext{O}2, ext{I}2, ext{Cl}2, ext{Br}_2. Noble gases (He, Ne, Ar, Kr, Xe) are monoatomic.
    • Subscript notation ➔ solid (s), liquid (l), gas (g) denoted as subscripts.

  • 3.4 Balancing Chemical Equations

    • A balanced equation has the same number of atoms of each element on both sides of the arrow.
    • Example: ext{H}2(g) + ext{O}2(g)
      ightarrow ext{H}_2 ext{O}(l)
    • Steps:
    • Step 1: Write skeleton with reactants and products.
    • Step 2: Count atoms of each element on both sides; treat polyatomic ions as a single unit if they appear unchanged.
    • Step 3: Balance by adjusting coefficients (trial and error), starting with elements that appear once on each side.
    • Step 4: Verify atom counts on both sides.
    • Step 5: Write coefficients in the lowest whole-number ratio.
    • Example balanced form: ext{H}2(g) + ext{O}2(g)
      ightarrow 2 ext{H}_2 ext{O}(l)

  • 3.5 Evidences of a Chemical Change

    • Color change.
    • Temperature change (heat released or absorbed).
    • Evolution of gas (bubbles).
    • Formation of a solid (precipitate).
    • Production of light.

  • 3.6 TYPES OF CHEMICAL REACTIONS

    • 1. Combustion Reaction ➔ A substance reacts with oxygen to release energy as heat and light; general equation: Fuel + O₂ → CO₂ + H₂O.
    • Examples:
      • CH4(g) + 2O2(g)
        ightarrow CO2(g) + 2H2O(l)
      • C3H8(g) + 5O2(g) ightarrow 3CO2(g) + 4H_2O(l)
      • C2H5OH(l) + 3O2(g) ightarrow 2CO2(g) + 3H_2O(l)
    • 2. Combination (Synthesis) Reaction ➔ Two or more substances combine to form a single product; general: A + B → AB.
    • Examples:
      • 2H2(g) + O2(g)
        ightarrow 2H_2O(l)
      • 2Na(s) + Cl_2(g)
        ightarrow 2NaCl(l)
      • 2CO(g) + O2(g) ightarrow 2CO2(g)
    • 3. Decomposition Reaction ➔ A single compound breaks down into simpler substances; energy required in heat, light, or electricity.
    • Examples:
      • CaCO3(s) ightarrow CaO(s) + CO2(g)
      • 2H2O(l) ightarrow 2H2(g) + O_2(g)
      • 2AgCl(l)
        ightarrow 2Ag(s) + Cl_2(g)
    • 4. Single Displacement Reaction ➔ One element replaces another in a compound; occurs when a more reactive element replaces a less reactive one.
    • General: A + BC → AC + B.
    • Examples:
      • Zn(s) + 2HCl(aq)
        ightarrow ZnCl2(aq) + H2(g)
      • Fe(s) + CuSO4(aq) ightarrow FeSO4(aq) + Cu(s)
      • Cl2(g) + 2NaBr(aq) ightarrow 2NaCl(aq) + Br2(l)
    • 5. Double Displacement (Metathesis) Reaction ➔ Exchange of positive ions between two compounds; often forms a precipitate.
    • General: AB + CD → AD + CB.
    • Examples:
      • Pb(NO3)2(aq) + 2KI(aq)
        ightarrow PbI2(s) + 2KNO3(aq)
      • AgNO3(aq) + NaCl(aq) ightarrow AgCl(s) + NaNO3(aq)
      • BaCl2(aq) + Na2SO4(aq) ightarrow BaSO4(s) + 2NaCl(aq)

PRACTICE PROBLEMS

  • I. Multiple Choice (questions summarized)

    1. Recover alcohol and water from a mixture using distillation. Answer: C) Distillation.
    2. Mixed components of sand, saltwater, and rocks: sequence to obtain pure components. Answer: D) Filtration → Mechanical Picking → Distillation.
    3. Homogeneous mixture among options: Answer: D) Saltwater.
    4. Which are all compounds? Answer: D) Ethanol, Vinegar, Rust.
    5. True statement about subatomic particles? Answer: D) Subatomic particles have different masses and charges (not all the same).
    6. Coefficients for a balanced equation CO + NO → CO₂ + N₂ (from given options) → Answer: A) 1, 1, 1, 2.
    7. Atomic model with electrons embedded in a positively charged blob? Answer: C) Plum Pudding Model.
    8. How atoms form compounds? Answer: C) By combining in whole-number proportions.
    9. An ion with +3 charge and 15 electrons: protons and neutrons? Answer: D) 18 protons, 16 neutrons.
    10. Sodium-24 isotope: Atomic Number, Atomic Mass, Protons, Electrons, Neutrons?
      • Answer: B) Atomic Number = 11, Atomic Mass = 24, Protons = 11, Electrons = 11, Neutrons = 13.

  • II. Complete the table (Individuals Ions)

    • 1. Cu₂O: Copper(I) Oxide from Cu⁺ and O²⁻.
    • 2. H₂CO₃: Carbonic Acid from H⁺ and CO₃²⁻.
    • 3. Na₂SO₄: Sodium Sulfate from Na⁺ and SO₄²⁻.
    • 4. Cr₂(SO₄)₃: Chromium(III) Sulfate from Cr³⁺ and SO₄²⁻.

  • III. Balance the following equations and identify the type of reaction

    1. A balloon filled with hydrogen gas reacts with oxygen to produce water vapor.
    • Type: Combination/Formation? (Typically, 2H₂ + O₂ → 2H₂O is a synthesis; water formation.)
    1. Ethanol (C₂H₅OH) is burned in oxygen to produce carbon dioxide and water.
    • Type: Combustion.
    1. A solution of sodium sulfate with barium chloride yields barium sulfate precipitate and sodium chloride in solution.
    • Type: Double Displacement (precipitation).
    1. Iron(III) oxide reacts with sulfur to produce iron and sulfur dioxide gas.
    • Type: Single Displacement (redox) with oxide reacting with sulfur to form Fe and SO₂.

Answer Key (as provided in the transcript)

  • I. Multiple Choice

    • 1. C | 2. B | 3. D | 4. D | 5. D | 6. A | 7. C | 8. C | 9. D | 10. B

  • II. Individual Ions

    • 1. Cu⁺ + O²⁻ → Cu₂O → Copper(I) Oxide
    • 2. H⁺ + CO₃²⁻ → H₂CO₃ → Carbonic Acid
    • 3. Na⁺ + SO₄²⁻ → Na₂SO₄ → Sodium Sulfate
    • 4. Cr³⁺ + SO₄²⁻ → Cr₂(SO₄)₃ → Chromium(III) Sulfate

  • III. Balance the following equations and identify the reaction type

    • 1. 2H₂ + O₂ → 2H₂O (Combination/ synthesis)
    • 2. C₂H₅OH + 3O₂ → 2CO₂ + 3H₂O (Combustion)
    • 3. Na₂SO₄ + BaCl₂ → BaSO₄ ↓ + 2NaCl (Double Displacement/ precipitation)
    • 4. 2Fe₂O₃ + 3S → 4Fe + 3SO₂ (Single Displacement/redox)