Chapter 1: Matter and Its Classification

Definition and Classification of Matter

Matter is anything that has mass and occupies space. It can be classified based on its composition and properties.

• Pure Substances: Consist of only one type of particle. They have fixed composition and distinct properties.

• Mixtures: Composed of two or more substances physically combined. Their composition can vary.

Types of Pure Substances

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

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

Types of Mixtures

• Homogeneous Mixtures (Solutions): Uniform composition throughout. Example: Saltwater, air.

• Heterogeneous Mixtures: Non-uniform composition; different parts are visible. Example: Salad, sand in water.

Physical and Chemical Changes

Changes in matter can be classified as physical or chemical.

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

• Chemical Change: Produces new substances with different properties. Example: Rusting of iron, burning wood.

Physical and Chemical Properties

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

• Chemical Properties: Describe a substance's ability to undergo chemical changes. Examples: Flammability, reactivity with acid.

Intensive vs. Extensive Properties

• Intensive Properties: Do not depend on the amount of matter. Examples: Density, boiling point.

• Extensive Properties: Depend on the amount of matter. Examples: Mass, volume.

SI Units and Measurement

• Seven Basic SI Units:

  • 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

Metric conversions use conversion factors and the factor-label (dimensional analysis) method.

• Example: Convert 5.0 cm to meters:

Density

• Definition: Density is mass per unit volume.

• Formula:

• Units: g/cm3 or kg/m3

Significant Figures and Scientific Notation

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

• Rules: All nonzero digits are significant; zeros between nonzero digits are significant; leading zeros are not significant; trailing zeros are significant if after a decimal point.

• Scientific Notation: Expresses numbers as a product of a number between 1 and 10 and a power of 10. Example:

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

Periodic Table Organization

The periodic table arranges elements by increasing atomic number (number of protons). Historically, elements were first arranged by atomic mass.

• Rows: Periods

• Columns: Groups or families

Element Families and Their Properties

• Alkali Metals (Group 1): Highly reactive, soft, react with water to form basic solutions.

• Alkaline Earth Metals (Group 2): Reactive, but less so than alkali metals; form basic oxides.

• Transition Metals (Groups 3-12): Good conductors, variable oxidation states, often form colored compounds.

• Halogens (Group 17): Very reactive nonmetals, form salts with metals.

• Noble Gases (Group 18): Inert, very low reactivity due to full valence shells.

Metals, Nonmetals, and Metalloids

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

• Nonmetals: Right side; dull, brittle, poor conductors.

• Metalloids: Border between metals and nonmetals; properties intermediate between metals and nonmetals. Examples: Silicon, arsenic.

Writing and Naming Chemical Compounds

• Ionic Compounds: Formed from metals and nonmetals. Name: cation (metal) + anion (nonmetal with -ide ending). Example: NaCl = sodium chloride.

• Molecular (Covalent) Compounds: Formed from nonmetals. Use prefixes to indicate number of atoms. Example: CO2 = carbon dioxide.

Binary and Ternary Compounds

• Binary Compounds: Contain two elements. Example: HCl, NaBr.

• Ternary Compounds: Contain three elements, often including a polyatomic ion. Example: NaNO3.

Transition Metal Ionic Compounds

• Transition metals can have multiple charges. Indicate charge with Roman numerals. Example: FeCl2 = iron(II) chloride.

Polyatomic Ions: Formulas and Charges

Percent Composition of a Compound

• Definition: The percent by mass of each element in a compound.

• Formula:

• Example: For H2O:

Hydrates: Formulas and Naming

• Hydrate: A compound that contains water molecules within its crystal structure.

• Naming: Name the compound, then add a prefix for the number of water molecules and the word 'hydrate.' Example: CuSO4·5H2O = copper(II) sulfate pentahydrate.

• Math: To find the formula, determine the ratio of moles of water to moles of anhydrous compound.

Counting Atoms in a Formula

• Multiply the subscript of each element by the coefficient (if present).

• Example: In 2 H2O, there are 4 H atoms and 2 O atoms.

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

• Definition: The mass of one mole of a substance, in grams per mole (g/mol).

• Calculation: Add the atomic masses of all atoms in the formula.

• Example: For NaCl:

The Mole Concept

• Definition: A mole is 6.02 × 1023 particles (Avogadro's number).

• Conversions:

  • 1 mole = 6.02 × 1023 particles

  • 1 mole of gas at STP = 22.4 L

  • 1 mole = molar mass in grams

• Example: To find moles from grams:

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

• Use the subscripts in a chemical formula to determine the ratio of elements.

• Example: In H2O, the mole ratio of H:O is 2:1.

• Mass ratios can be found by multiplying the number of moles by the atomic mass.

Empirical and Molecular Formulas

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

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

• Calculation:

  1. Find moles of each element.

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

  3. For molecular formula: where

• Example: If empirical formula is CH2O and molar mass is 180 g/mol, empirical mass is 30 g/mol, so and molecular formula is C6H12O6.

Additional info: Some explanations and examples have been expanded for clarity and completeness, as is standard in academic study guides.