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: Materials with a fixed composition and distinct properties.

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

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

• Mixtures: Physical combinations of two or more substances where each retains its own properties.

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

  • Heterogeneous Mixtures: Non-uniform composition. 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. Examples: Melting ice, dissolving sugar in water.

• Chemical Change: Produces new substances with different properties. Examples: Rusting 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 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:

Temperature Conversions

• To convert between Celsius, Fahrenheit, and Kelvin:

Metric Conversions and Factor-Label Method

• Use conversion factors to convert between units.

• Example: To convert 5.0 cm to meters:

Density

• Density is the mass per unit volume of a substance.

• Formula:

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

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 number between 1 and 10 and a power of 10. Example:

Chapter 2 & 3: Atomic Structure, Periodic Table, and Chemical Formulas

Periodic Table Organization

The periodic table arranges elements by increasing atomic number (number of protons). Historically, it was first organized by atomic mass.

• Groups (columns): Elements with similar chemical properties.

• Periods (rows): Elements with the same number of electron shells.

Families of Elements

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

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

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

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

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

Metals, Nonmetals, and Metalloids

• Metals: Good conductors, malleable, ductile, shiny; found on the left and center of the table.

• Nonmetals: Poor conductors, brittle, dull; found on the right side.

• Metalloids: Properties intermediate between metals and nonmetals; found along the staircase line (e.g., Si, B).

Writing and Naming Chemical Compounds

• Ionic Compounds: Formed from metals and nonmetals. Name: cation + anion (e.g., NaCl: sodium chloride).

• Molecular Compounds: Formed from nonmetals. Use prefixes (e.g., CO2: carbon dioxide).

• Binary Compounds: Contain two elements (e.g., HCl).

• Ternary Compounds: Contain three elements, often with polyatomic ions (e.g., NaNO3).

• Transition Metal Compounds: Indicate charge with Roman numerals (e.g., FeCl2: iron(II) chloride).

Polyatomic Ions

Common polyatomic ions and their charges:

Percent Composition

• Percent composition is the percent by mass of each element in a compound.

• Formula:

• Example: In H2O, %H =

Hydrates: Formulas and Naming

• Hydrates are compounds with water molecules attached.

• Named as: [compound name] + prefix-hydrate (e.g., CuSO4·5H2O: copper(II) sulfate pentahydrate).

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

Counting Atoms in a Formula

• Multiply the subscript of each element by the number of formula units or molecules.

• Example: In Al2(SO4)3, there are 2 Al, 3 × 1 S = 3 S, and 3 × 4 O = 12 O atoms.

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

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

• Calculated by summing the atomic masses of all atoms in the formula.

• Example: NaCl: Na (22.99) + Cl (35.45) = 58.44 g/mol.

The Mole Concept

• One mole contains Avogadro's number of particles: .

• 1 mole of a gas at STP occupies 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 the subscripts in a chemical formula to determine the ratio of moles of each element.

• Mass ratios can be found by multiplying the mole ratio by the molar masses.

Empirical and Molecular Formulas

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

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

• To determine empirical formula from percent composition:

  1. Convert percentages to grams (assume 100 g sample).

  2. Convert grams to moles for each element.

  3. Divide by the smallest number of moles to get the simplest ratio.

• To find molecular formula: , where

Additional info: Some explanations and examples have been expanded for clarity and completeness.