CHM165 Exam 1 Study Guide: Chapters 1-3 (Matter, Atoms, Molecules, and Compounds)

Chapter 1: Matter, Measurement & Problem Solving

Scientific Method in Chemistry

The scientific method is a systematic approach used in chemistry to investigate phenomena, acquire new knowledge, or correct and integrate previous knowledge. It involves observation, hypothesis formation, experimentation, and analysis.

• Observation: Gathering data through senses or instruments.

• Hypothesis: A tentative explanation for an observation.

• Experiment: Testing the hypothesis under controlled conditions.

• Theory: A well-substantiated explanation based on repeated experiments.

• Law: A statement describing consistent natural phenomena (e.g., Law of Conservation of Mass).

Correlation vs Causation: Correlation is a relationship between two variables, while causation indicates that one variable directly affects another. In chemistry, experiments are designed to distinguish between these.

Matter, Atoms, and Molecules

Matter is anything that has mass and occupies space. It is composed of atoms (the smallest unit of an element) and molecules (two or more atoms bonded together).

• States of Matter: Solid, liquid, gas.

• Physical Change: Changes affecting form but not composition (e.g., melting ice).

• Chemical Change: Changes resulting in new substances (e.g., rusting iron).

Energy: Kinetic and Potential

Energy is the capacity to do work. In chemistry:

• Kinetic Energy: Energy due to motion.

• Potential Energy: Stored energy due to position or composition.

Measurement and SI Units

Chemists use the SI (International System of Units) for measurements:

• Length: meter (m)

• Mass: kilogram (kg)

• Time: second (s)

• Temperature: kelvin (K)

• Amount: mole (mol)

Temperature conversions:

• $C = \frac{5}{9}(F - 32)$

• $K = C + 273.15$

Conversions and Dimensional Analysis

Dimensional analysis is used to convert between units using conversion factors. For example, converting meters to milliliters requires knowledge of metric prefixes.

Density: Measures mass per unit volume.

• $\text{Density} = \frac{\text{Mass}}{\text{Volume}}$

Properties of Matter

• Extensive Properties: Depend on the amount of matter (e.g., mass, volume).

• Intensive Properties: Independent of the amount (e.g., density, temperature).

Significant Figures and Exponential Notation

Significant figures reflect the precision of measurements. Exponential notation is used for very large or small numbers (e.g., $1.0 \times 10^3$).

Accuracy, Precision, and Error

• Accuracy: Closeness to the true value.

• Precision: Repeatability of measurements.

• % Error Calculation: $\text{Percent Error} = \frac{\text{Measured Value} - \text{True Value}}{\text{True Value}} \times 100\%$

• Random Error: Unpredictable variations.

• Systematic Error: Consistent, repeatable error.

Chapter 2: Atoms & Elements

Historical Development of Atomic Theory

• Aristotle vs Democritus: Democritus proposed atoms; Aristotle rejected the idea.

• Antoine Lavoisier: Law of Conservation of Mass: Mass is neither created nor destroyed.

• Joseph Proust: Law of Definite Proportions: A compound always contains the same proportion of elements.

• John Dalton: Law of Multiple Proportions and Atomic Theory.

Dalton's Atomic Theory:

• Elements are composed of atoms.

• Atoms of the same element are identical.

• Atoms combine in simple ratios to form compounds.

• Atoms are indivisible in chemical reactions.

Atomic Models and Discoveries

• Dalton’s Model: Solid, indivisible spheres.

• Plum Pudding Model (Thomson): Electrons embedded in a positive sphere.

• Nuclear Model (Rutherford): Dense nucleus with electrons orbiting.

• JJ Thomson: Discovered electrons via cathode ray experiments.

• Millikan: Oil drop experiment determined electron charge.

• Maria Curie: Studied radioactivity; discovered alpha, beta, gamma particles.

• Rutherford: Gold foil experiment revealed nucleus.

• Chadwick: Discovered the neutron.

Subatomic Particles and Atomic Structure

• Proton: Positive charge, mass ≈ 1 amu.

• Neutron: Neutral, mass ≈ 1 amu.

• Electron: Negative charge, mass ≈ 0.0005 amu.

Atomic number ($Z$): Number of protons. Mass number ($A$): Number of protons + neutrons.

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

Atomic mass calculation:

• $\text{Atomic Mass} = \sum (\text{Isotope Mass} \times \text{Fractional Abundance})$

Ions and Periodic Table Organization

• Ions: Atoms that have gained or lost electrons.

• Mendeleev: Organized elements by increasing atomic mass; predicted undiscovered elements.

• Periodic Table: Main group, transition metals, rare earth metals, metals, non-metals, metalloids.

• Families: Alkali metals, alkaline earth metals, halogens, noble gases.

The Mole and Chemical Calculations

• Mole: $6.022 \times 10^{23}$ particles (Avogadro's number).

• Use the periodic table and dimensional analysis to determine atoms, electrons, protons, neutrons, moles, and grams.

Chapter 3: Molecules and Compounds

Chemical Bonds and Formulas

Chemical bonds hold atoms together in compounds. Types include:

• Ionic Bonds: Transfer of electrons between metal and nonmetal.

• Covalent Bonds: Sharing of electrons between nonmetals.

• Metallic Bonds: Delocalized electrons among metals.

Chemical formulas represent the composition of compounds.

Empirical and Molecular Formulas

• Empirical Formula: Simplest whole-number ratio of elements.

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

Classifying Elements and Compounds

• Element: Pure substance of one type of atom.

• Compound: Substance composed of two or more elements.

Writing and Naming Chemical Formulas

• Metal + Nonmetal: Name metal, then nonmetal with -ide.

• Transition Metal + Nonmetal: Use Roman numerals for charge.

• Polyatomic Ions: Use standard names (e.g., sulfate, nitrate).

• Nonmetal + Nonmetal: Use prefixes (mono-, di-, tri-).

• Acids: Binary (hydro- + root + -ic acid), Polyatomic (root + -ic/-ous acid).

• Hydrates: Compound + "n" water molecules (e.g., CuSO4·5H2O).

Molar Mass and Mass Percent

• Molar Mass: Sum of atomic masses in a compound (g/mol).

• Mass Percent: $\text{Mass \%} = \frac{\text{Mass of Element}}{\text{Total Mass of Compound}} \times 100\%$

Conservation of Matter and States of Matter Symbols

• Conservation of Matter: Matter is not created or destroyed in chemical reactions.

• State Symbols: (s) solid, (l) liquid, (g) gas, (aq) aqueous.

Organic Compounds and Bond Types

• Single Bonds: Alkane (C–C)

• Double Bonds: Alkene (C=C)

• Triple Bonds: Alkyne (C≡C)

• Organic compounds are typically made of carbon and hydrogen.

Example: Ethane (C2H6) is an alkane with single bonds.

Additional info:

• For Table 1.2 conversions, refer to standard metric prefixes (e.g., 1 m = 1000 mm).

• For organic compound naming, use IUPAC rules for alkanes, alkenes, and alkynes.