Chapter 1: Matter, Measurement & Problem Solving

Nomenclature of Chemistry and Science

Understanding the language of chemistry is essential for clear communication and comprehension of scientific concepts.

• Scientific Law: A concise statement that summarizes observed phenomena; it describes what happens but not why.

• Scientific Theory: A well-substantiated explanation of some aspect of the natural world that can incorporate laws, hypotheses, and facts.

• Hypothesis: A tentative explanation or prediction that can be tested by experiments.

• Development: Scientists use observations, experiments, and logical reasoning to develop hypotheses, which may become theories or laws after extensive validation.

• Example: The Law of Conservation of Mass states that mass is neither created nor destroyed in a chemical reaction.

Types of Matter

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

• Elements: Pure substances consisting of only one type of atom (e.g., O2).

• Compounds: Substances composed of two or more elements chemically combined in fixed ratios (e.g., H2O).

• Molecules: Two or more atoms bonded together; can be elements (O2) or compounds (CO2).

• Mixtures: Physical combinations of two or more substances.

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

  • Heterogeneous Mixtures: Non-uniform composition (e.g., salad, sand in water).

States of Matter

Matter exists in different physical forms called states or phases.

• Solids: Definite shape and volume; particles are closely packed.

  • Crystalline Solids: Particles arranged in an orderly pattern (e.g., salt, diamond).

  • Amorphous Solids: Particles lack a long-range order (e.g., glass, plastic).

• Liquids: Definite volume but no definite shape; particles are close but can move past each other.

• Gases: No definite shape or volume; particles are far apart and move freely.

Drawing Particles in Different Phases

• Solids: Densely packed, orderly arrangement.

• Liquids: Close together but disordered.

• Gases: Widely spaced, random arrangement.

Energy Changes in Phase Transitions

Phase changes involve energy transfer, typically as heat.

• Melting (solid to liquid): Absorbs energy.

• Vaporization (liquid to gas): Absorbs energy.

• Freezing (liquid to solid): Releases energy.

• Condensation (gas to liquid): Releases energy.

• Sublimation (solid to gas): Absorbs energy.

• Deposition (gas to solid): Releases energy.

Physical vs. Chemical Changes

• Physical Change: Alters the form or appearance but not the composition (e.g., melting ice).

• Chemical Change: Alters the composition, producing new substances (e.g., rusting iron).

• Example: Dissolving sugar in water is a physical change; burning wood is a chemical change.

Conservation Laws

• Law of Conservation of Mass: Mass is conserved in a chemical reaction.

• Law of Conservation of Energy: Energy cannot be created or destroyed, only transformed.

• Potential Energy: Stored energy due to position or composition (e.g., a ball at the top of a hill).

• Example: A magnet attached to a fridge has lower potential energy than one held away from the fridge.

Units of Measurement

• SI Units: Standard units used in science (meter, kilogram, second, etc.).

• Imperial Units: Common in the US (inch, pound, gallon, etc.).

• Unit Conversion: Use conversion factors to switch between units.

  • Example:

• Metric Prefixes: Indicate scale (e.g., kilo-, centi-, milli-).

• Volume Conversion: For cubic units, cube the conversion factor.

  • Example:

Density

• Definition: Density is mass per unit volume.

• Formula:

• Applications: Use density to convert between mass and volume.

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

Significant Figures and Mathematical Operations

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

• Rules:

  • Nonzero digits are always significant.

  • Zeros between nonzero digits are significant.

  • Leading zeros are not significant.

  • Trailing zeros are significant if there is a decimal point.

• Rounding: Round to the correct number of significant figures based on the operation performed.

• Precision vs. Accuracy:

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

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

Chapter 2: Atoms & Elements

Atomic Theory and Models

• Dalton’s Atomic Theory: Matter is composed of small, indivisible particles called atoms. Atoms of the same element are identical; atoms combine in simple ratios to form compounds. (Note: Some points have been revised with modern discoveries.)

• Rigid Sphere Model: Atoms are solid, indestructible spheres.

Discovery of Subatomic Particles

• Thomson’s Cathode Ray Experiment: Discovered the electron by observing that cathode rays were deflected by electric and magnetic fields, indicating the presence of negatively charged particles.

• Plum Pudding Model: Atoms are spheres of positive charge with embedded electrons.

• Rutherford’s Gold Foil Experiment: Alpha particles were directed at thin gold foil; most passed through, but some were deflected, leading to the discovery of the nucleus.

• Planetary Model: Atoms have a small, dense, positively charged nucleus surrounded by electrons.

Subatomic Particles

• Proton: Positive charge (+1), mass ≈ 1 amu.

• Neutron: No charge (0), mass ≈ 1 amu.

• Electron: Negative charge (–1), mass ≈ 0.0005 amu.

The Periodic Table

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

• Atomic Symbol: One- or two-letter abbreviation for an element (e.g., H for hydrogen).

• Atomic Mass: Weighted average mass of an element’s isotopes (in amu).

Isotopes, Ions, and Periodic Trends

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

• Cations: Positively charged ions (loss of electrons).

• Anions: Negatively charged ions (gain of electrons).

• Metals vs. Nonmetals: Metals are typically on the left and center of the periodic table; nonmetals are on the right.

• Common Ion Charges:

  • Alkali metals (Group 1): +1

  • Alkaline earth metals (Group 2): +2

  • Halogens (Group 17): –1

Compounds and Chemical Bonds

• Ionic Bonds: Formed by transfer of electrons from metals to nonmetals; result in a lattice of ions.

• Covalent Bonds: Formed by sharing electrons between nonmetals; result in discrete molecules.

• Properties: Ionic compounds are usually solid, have high melting points, and conduct electricity when molten or dissolved. Covalent molecules can be gases, liquids, or solids and generally do not conduct electricity.

Chemical Formulas and Molecular Representations

• Chemical Formula: Indicates the types and numbers of atoms in a compound (e.g., NaCl, H2O).

• Molecular Formula: Shows the actual number of each type of atom in a molecule.

• Structural Formula: Shows how atoms are connected.

• Ball-and-Stick Model: 3D representation showing atoms as balls and bonds as sticks.

• Space-Filling Model: Shows the relative sizes of atoms and how they overlap.

Diatomic Elements

• Some elements exist naturally as molecules composed of two atoms. The seven diatomic elements are:

  • H2, N2, O2, F2, Cl2, Br2, I2

Chapter 3: Molecules and Compounds

The Mole and Avogadro’s Number

The mole is a counting unit used in chemistry to relate the mass of substances to the number of particles.

• Definition: One mole contains entities (Avogadro’s number).

• Application: Used to count atoms, molecules, ions, or other particles.

• Example: 1 mole of H2O contains molecules of water.

Molar Mass

• Atomic Mass: Mass of one atom (in amu).

• Molecular Mass: Sum of atomic masses in a molecule.

• Formula Mass: Sum of atomic masses in an ionic compound’s formula unit.

• Calculating Molar Mass: Add the atomic masses of all atoms in the formula.

  • Example: Molar mass of H2O = (2 × 1.01) + (1 × 16.00) = 18.02 g/mol

Dimensional Analysis with Moles

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

• Common Conversions:

  • Grams ↔ Moles:

  • Moles ↔ Particles:

• Example: How many molecules are in 36.0 g of H2O?

  • Step 1:

  • Step 2: molecules