Chapter 1: Chemistry – Matter on the Atomic Scale

1.1 The Scale of Chemistry

Chemistry is the study of matter and the changes it undergoes, from the macroscopic world we see to the atomic and molecular scale. Understanding the scientific method and key scientific terms is foundational.

• Scientific Method: A systematic approach to research involving observation, hypothesis formation, experimentation, and theory development.

• Hypothesis: A tentative explanation for an observation, which can be tested by experiments.

• Law: A concise statement or mathematical equation that summarizes observed behavior (e.g., Law of Conservation of Mass).

• Theory: A well-substantiated explanation of some aspect of the natural world that can incorporate laws, hypotheses, and facts (e.g., Atomic Theory).

• Chemistry: The science that studies the composition, properties, and transformations of matter.

• Matter: Anything that has mass and occupies space.

• Macroscopic Scale: Observable with the naked eye (e.g., a glass of water).

• Atomic/Molecular Scale: The scale of atoms and molecules, not visible without special instruments.

1.2 Classification of Matter

Matter can be classified based on its composition and properties. Understanding these classifications is essential for studying chemical changes and reactions.

• Atom: The smallest unit of an element that retains its chemical properties.

• Molecule: Two or more atoms chemically bonded together.

• Physical Property: A characteristic that can be observed without changing the substance's identity (e.g., color, melting point).

• Chemical Property: Describes a substance's ability to undergo chemical changes (e.g., flammability).

• Physical Change: A change that does not alter the chemical composition (e.g., melting, boiling).

• Chemical Change: A process that alters the chemical composition (e.g., rusting of iron).

• States of Matter:

  • Solid: Definite shape and volume.

  • Liquid: Definite volume, takes shape of container.

  • Gas: No definite shape or volume.

• Melting Point: Temperature at which a solid becomes a liquid.

• Freezing Point: Temperature at which a liquid becomes a solid.

• Boiling Point: Temperature at which a liquid becomes a gas.

• Mixture: Physical combination of two or more substances.

  • Homogeneous Mixture (Solution): Uniform composition throughout (e.g., saltwater).

  • Heterogeneous Mixture: Non-uniform composition (e.g., salad).

• Pure Substance: Matter with a fixed composition.

  • Element: Cannot be broken down into simpler substances.

  • Compound: Composed of two or more elements chemically combined.

1.3 Units and Measurement

Accurate measurement is fundamental in chemistry. The SI system is used for standardization, and understanding units, prefixes, and significant figures is essential.

• Metric Prefixes: kilo (103), deci (10-1), centi (10-2), milli (10-3), micro (10-6), nano (10-9).

• Scientific Notation: Expresses numbers as a product of a coefficient and a power of ten (e.g., ).

• Definitions:

  • Length: Meter (m)

  • Mass: Kilogram (kg)

  • Temperature: Kelvin (K), Celsius (°C)

  • Volume: Liter (L)

  • Energy: Joule (J)

• Density:

• Temperature Conversion:

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

• Precision: How close repeated measurements are to each other.

• Accuracy: How close a measurement is to the true value.

1.4 Unit Conversions

Unit conversions are performed using dimensional analysis, ensuring that units cancel appropriately and significant figures are maintained.

• Common Conversions: (Refer to your course's conversion sheet for those marked "Memorize").

• Dimensional Analysis: A method using conversion factors to change units.

• Density, Mass, Volume Relationship:

  • Given any two, calculate the third using .

Chapter 2: Elements and Compounds

2.1 The Structure of the Atom

Atoms are composed of subatomic particles with distinct properties. Understanding atomic structure is key to explaining chemical behavior.

• Proton: Positive charge (+1), mass ≈ 1 amu, located in the nucleus.

• Neutron: No charge (0), mass ≈ 1 amu, located in the nucleus.

• Electron: Negative charge (-1), mass ≈ 1/1836 amu, located outside the nucleus.

• Atomic Number (Z): Number of protons in the nucleus.

• Mass Number (A): Total number of protons and neutrons.

• Atomic Mass: Weighted average mass of an element's isotopes.

• Isotope: Atoms of the same element with different numbers of neutrons.

• Nuclide Symbol: , where X is the element symbol.

• Calculating Subatomic Particles:

  • Protons = Atomic number (Z)

  • Neutrons = Mass number (A) – Atomic number (Z)

  • Electrons = Protons (for neutral atoms); adjust for ion charge

• Atomic Weight Calculation:

2.2 Elements and the Periodic Table

The periodic table organizes elements by atomic number and properties. Recognizing element groups and their characteristics is essential for predicting chemical behavior.

• Atomic Number: Number of protons; unique to each element.

• Atomic Mass: Average mass of an element's atoms.

• Element Names and Symbols: Know elements 1–56 by name and symbol.

• Periodic Table Structure:

  • Period: Horizontal row.

  • Group (Family): Vertical column.

  • Main Group Elements: Groups 1A–2A and 3A–8A.

  • Transition Metals: Groups 3B–12B.

  • Lanthanides and Actinides: Two rows below the main table.

  • Metals, Nonmetals, Metalloids: Classified by physical and chemical properties.

  • Alkali Metals: Group 1A.

  • Alkaline Earth Metals: Group 2A.

  • Halogens: Group 7A.

  • Noble Gases: Group 8A.

• Allotropes: Different forms of the same element in the same physical state (e.g., diamond and graphite for carbon).

• Diatomic Elements: Elements that exist as molecules of two atoms: H2, N2, O2, F2, Cl2, Br2, I2.

2.3 Covalent Compounds

Covalent (molecular) compounds are formed by the sharing of electrons between nonmetal atoms. They can be represented in various ways and are named using specific rules.

• Covalent vs. Ionic Compounds:

  • Covalent: Nonmetals sharing electrons.

  • Ionic: Transfer of electrons between metals and nonmetals.

• Molecular Representations:

  • Molecular Formula: Shows the number and type of atoms (e.g., H2O).

  • Condensed Formula: Groups atoms to show structure (e.g., CH3CH2OH).

  • Structural Formula: Shows how atoms are bonded.

  • Wedge-and-Dash: Indicates 3D structure.

  • Ball-and-Stick, Space-Filling Models: Physical models of molecular shape.

• Naming Covalent Compounds: Use prefixes (mono-, di-, tri-, etc.) to indicate the number of each atom (e.g., CO2 is carbon dioxide).

2.4 Ions and Ionic Compounds

Ionic compounds are formed from the electrostatic attraction between cations and anions. Understanding ion charges and naming conventions is crucial.

• Cation: Positively charged ion (loss of electrons).

• Anion: Negatively charged ion (gain of electrons).

• Polyatomic Ions: Ions composed of multiple atoms (e.g., SO42–, NO3–).

• Common Polyatomic Ions: (Refer to Table 2.4.1 in your text; memorize those required by your instructor.)

• Predicting Charges:

  • Group 1A (alkali metals): +1

  • Group 2A (alkaline earth metals): +2

  • Group 5A: –3

  • Group 6A: –2

  • Group 7A (halogens): –1

  • Aluminum (Al): +3

  • Silver (Ag): +1

  • Zinc (Zn), Cadmium (Cd): +2

• Formulas of Ionic Compounds: Combine cations and anions in ratios that yield a neutral compound (e.g., Na+ and Cl– form NaCl).

• Naming Ionic Compounds: Name the cation first, then the anion (e.g., sodium chloride).

• Identifying Compound Type: Compounds with only nonmetals are usually covalent; those with metals and nonmetals are usually ionic.

Example Table: Classification of Matter

Example Table: Subatomic Particles

Additional info: Some details, such as the full list of polyatomic ions and conversion factors, should be referenced from your course materials as required by your instructor.