Chapter 1: Matter and Measurement

Matter: Definition and Classification

Matter is anything that has mass and occupies space. It can be classified into two main types: pure substances and mixtures.

• Pure Substances: Have a fixed composition and distinct properties. They are further divided into:

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

  • 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. They can be separated by physical means and are classified as:

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

  • Heterogeneous Mixtures: Non-uniform composition. Example: Sand and iron filings.

Physical and Chemical Changes

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

• Chemical Change: Results in the formation of new substances. Examples: Rusting of iron, burning of 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: Independent of the amount of substance. Examples: Density, boiling point.

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

SI Units and Measurement

• The seven basic SI units are:

Temperature Conversions

• To convert between Celsius, Fahrenheit, and Kelvin:

Metric Conversions and Factor-Label Method

• Use conversion factors to change units. The factor-label method (dimensional analysis) involves multiplying by fractions that represent the relationship 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 coefficient and a power of ten. Example:

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

The Periodic Table: Organization and Families

• The periodic table was originally organized by atomic mass, but is now arranged by atomic number (number of protons).

• Families/Groups:

  • Alkali Metals (Group 1): Highly reactive, soft metals (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, colorless gases (e.g., Ne, Ar).

  • Transition Metals (Groups 3-12): Metals with variable charges and colored compounds (e.g., Fe, Cu).

Metals, Nonmetals, and Metalloids

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

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

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

Writing and Naming Chemical Compounds

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

• Molecular Compounds: Formed from nonmetals. Use prefixes to indicate number of atoms (e.g., CO2: carbon dioxide).

• Binary Compounds: Contain two elements (e.g., H2O).

• Ternary Compounds: Contain three different elements (e.g., NaNO3).

• Transition Metal Ionic Compounds: Use Roman numerals to indicate the metal's charge (e.g., FeCl2: iron(II) chloride).

Polyatomic Ions: Formulas and Charges

Percent Composition of a Compound

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

• Formula:

• Example: In H2O, %H =

Hydrates: Formulas and Naming

• Hydrate: An ionic compound with water molecules attached.

• Naming: Name the compound, then add a prefix for the number of waters (e.g., CuSO4·5H2O: copper(II) sulfate pentahydrate).

• Math: Calculate the percent water or determine the formula from mass data.

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)

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

• Sum the atomic masses of all atoms in the formula.

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

The Mole Concept

• One mole contains Avogadro's number () of particles (atoms, molecules, or ions).

• At STP (Standard Temperature and Pressure), one mole of any gas occupies 22.4 L.

• Conversions:

  • Grams to moles:

  • Moles to particles:

  • Moles to volume (gas at STP):

Mole Ratios and Mass Ratios

• In a chemical formula, the ratio of moles of each element is given by the subscripts.

• Mass ratios can be found by multiplying the mole ratio by the molar mass of each element.

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

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: Divide the molecular mass by the empirical formula mass to find the multiple.

• Example: If the empirical formula is CH2O (mass = 30 g/mol) and the molecular mass is 180 g/mol, the molecular formula is C6H12O6.

Additional info: Some explanations and examples have been expanded for clarity and completeness, following standard general chemistry curriculum.