Chapter 1: Matter and Its Classification

Definition and Classification of Matter

• Matter is anything that has mass and occupies space.

• Matter can be classified into two main types:

  • Pure Substances: Have a fixed composition and distinct properties.

  • Mixtures: Combinations of two or more substances where each retains its own identity.

Pure Substances: Elements and Compounds

• Element: A substance that cannot be broken down into simpler substances by chemical means. Example: Oxygen (O2), Gold (Au).

• Compound: A substance composed of two or more elements chemically combined in fixed proportions. Example: Water (H2O), Sodium chloride (NaCl).

Mixtures: Homogeneous and Heterogeneous

• Homogeneous Mixture: Uniform composition throughout. Example: Saltwater, air.

• Heterogeneous Mixture: Non-uniform composition; components are visibly distinguishable. Example: Salad, sand and iron filings.

Physical and Chemical Changes

• Physical Change: Alters the form or appearance but not the composition. Example: Melting ice, dissolving sugar in water.

• Chemical Change: Results in the formation of one or more new substances. Example: Rusting of iron, burning of wood.

Physical and Chemical Properties

• Physical Property: Can be observed without changing the substance's identity. Example: Color, melting point, density.

• Chemical Property: Describes a substance's ability to undergo chemical changes. Example: Flammability, reactivity with acid.

• Intensive Property: Independent of the amount of substance. Example: Density, boiling point.

• Extensive Property: Depends on the amount of substance. Example: Mass, volume.

SI Units

• The seven basic SI units are:

  • Length: meter (m)

  • Mass: kilogram (kg)

  • Time: second (s)

  • Temperature: kelvin (K)

  • Amount of substance: mole (mol)

  • Electric current: ampere (A)

  • Luminous intensity: candela (cd)

Temperature Conversions

• To convert between Celsius, Fahrenheit, and Kelvin:

Metric Conversions and Factor Label Method

• Use conversion factors to convert between units.

• Factor Label Method (Dimensional Analysis): Multiply by conversion factors so units cancel appropriately.

• Example: Convert 5.0 cm to meters:

Density

• Density is mass per unit volume.

• Formula:

• Example: If a block has a mass of 10 g and a volume of 2 cm3, its density is 5 g/cm3.

Significant Figures and Scientific Notation

• Significant Figures: Digits in a measurement that are known with certainty plus one estimated digit.

• Scientific Notation: Expresses numbers as a product of a coefficient and a power of ten. Example:

Chapter 2 & 3: Atoms, Elements, and Chemical Calculations

Periodic Table Organization

• Originally organized by atomic mass; now organized by atomic number (number of protons).

• Elements are arranged in periods (rows) and groups (columns) with similar properties.

Families of Elements: Names and Properties

• Alkali Metals (Group 1): Highly reactive, soft, react with water (e.g., Na, K).

• Alkaline Earth Metals (Group 2): Reactive, but less so than alkali metals (e.g., Mg, Ca).

• Halogens (Group 17): Very reactive nonmetals (e.g., Cl, F).

• Noble Gases (Group 18): Inert, very low reactivity (e.g., Ne, Ar).

• Transition Metals (Groups 3-12): Good conductors, variable oxidation states (e.g., Fe, Cu).

Metals, Nonmetals, and Metalloids

• Metals: Left and center of the periodic table; shiny, malleable, good conductors.

• Nonmetals: Upper right; dull, brittle, poor conductors.

• Metalloids: Border between metals and nonmetals; properties intermediate between metals and nonmetals (e.g., Si, B).

Formula Writing and Naming Compounds

• Ionic Compounds: Metal + nonmetal; name metal first, then nonmetal with '-ide' ending. Example: NaCl = sodium chloride.

• Molecular Compounds: Nonmetal + nonmetal; use prefixes (mono-, di-, tri-, etc.). Example: CO2 = carbon dioxide.

Binary, Tertiary, and Transition Metal Compounds

• Binary Compounds: Contain two elements. Example: H2O.

• Tertiary Compounds: Contain three different elements. Example: NaNO3.

• Transition Metal Ionic Compounds: Use Roman numerals to indicate charge. Example: FeCl3 = iron(III) chloride.

Polyatomic Ions: Formulas and Charges

• Common polyatomic ions include:

Percent Composition of a Formula

• Percent composition = (mass of element in 1 mol of compound / molar mass of compound) × 100%

• Example: For H2O: %H = (2 × 1.01) / 18.02 × 100% ≈ 11.2%

Hydrates: Formula and Naming

• Hydrate: Compound with water molecules attached.

• Naming: Use prefixes to indicate number of water molecules (e.g., CuSO4·5H2O = copper(II) sulfate pentahydrate).

• Math: Calculate mass of water and anhydrous compound separately.

Counting Number of Atoms

• Multiply subscripts by coefficients to find total atoms.

• Example: 2 H2O has 4 H and 2 O atoms.

Gram Formula Mass (GFM) or Molar Mass (MM)

• Molar Mass: Mass of one mole of a substance (g/mol).

• Sum atomic masses of all atoms in the formula.

• Example: Molar mass of CO2 = 12.01 + 2 × 16.00 = 44.01 g/mol.

Mole Concept and Conversions

• 1 mole = particles (Avogadro's number).

• At STP, 1 mole of gas = 22.4 L.

• Conversions:

  • Grams to moles:

  • Moles to particles:

  • Moles to volume (gas at STP):

Mole-to-Mole and Mass-to-Mass Ratios

• Use coefficients from balanced equations to relate moles of reactants and products.

• Mass-to-mass: Convert mass to moles, use mole ratio, then convert back to mass.

Empirical and Molecular Formula Calculations

• Empirical Formula: Simplest whole-number ratio of atoms in a compound.

• Molecular Formula: Actual number of atoms of each element in a molecule.

• Calculation Steps:

  1. Find moles of each element.

  2. Divide by smallest number of moles to get ratio.

  3. Multiply to get whole numbers if necessary.

  4. For molecular formula: , where

Additional info: Where the original notes were brief, standard textbook definitions and examples were added for clarity and completeness.