Study Guide: Matter, Atoms, and Molecules (Chapters 1–3, CHEM 134)

Chapter 1: Matter, Measurement, and Problem Solving

Classification of Matter

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

• States of Matter: Solid, liquid, gas

• Composition: Elements, compounds, mixtures (homogeneous and heterogeneous)

• Example: Air (homogeneous mixture), sand and iron filings (heterogeneous mixture)

Separation Techniques

Physical methods are used to separate mixtures based on differences in physical properties.

• Filtration: Separates solids from liquids using a porous barrier.

• Distillation: Separates substances based on differences in boiling points.

• Example: Saltwater can be separated by distillation; sand from water by filtration.

Physical vs. Chemical Changes and Properties

• Physical Change: Alters state or appearance without changing composition (e.g., melting ice).

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

• Physical Property: Observed without changing composition (e.g., density, color).

• Chemical Property: Observed during a chemical change (e.g., flammability).

Energy in Chemistry

• Forms of Energy: Kinetic, potential, thermal, chemical, electrical, etc.

• Importance: Energy changes accompany chemical and physical processes.

• Example: Combustion releases thermal energy.

Measurement and Units

• SI Units: Standard units for scientific measurement (meter, kilogram, second, mole, etc.).

• Prefix Multipliers: kilo- (103), centi- (10-2), milli- (10-3), etc.

• Derived Units: Formed from base units (e.g., m/s, g/cm3).

Principles of Measurement

• Precision: Consistency of repeated measurements.

• Accuracy: Closeness to the true value.

• Systematic Error: Consistent, repeatable error (e.g., miscalibrated instrument).

• Random Error: Unpredictable variations.

• Statistical Calculations: Average, median, relative error (RE%).

• Equation for Relative Error (%): $\mathrm{RE\%} = \frac{\lvert \text{measured value} - \text{true value} \rvert}{\text{true value}} \times 100\%$

Unit Conversions and Dimensional Analysis

• Unit Conversion: Changing from one unit to another using conversion factors.

• Dimensional Analysis: Systematic approach to problem solving using units as a guide.

• Example: Converting 5.0 g to mg: $5.0\ \mathrm{g} \times \frac{1000\ \mathrm{mg}}{1\ \mathrm{g}} = 5000\ \mathrm{mg}$

Graph Interpretation

• Graphs visually represent data; understanding axes, trends, and relationships is essential for analysis.

Chapter 2: Atoms and Elements

Fundamental Laws

• Law of Conservation of Mass: Mass is neither created nor destroyed in a chemical reaction.

• Law of Definite Proportions: A compound always contains the same elements in the same proportion by mass.

• Law of Multiple Proportions: Elements can combine in different ratios to form different compounds.

Key Experiments

• Thomson: Discovered the electron using cathode ray tubes.

• Millikan: Measured the charge of the electron (oil drop experiment).

• Rutherford: Discovered the nucleus via gold foil experiment.

Dalton’s Atomic Theory

• All matter is composed of atoms.

• Atoms of a given element are identical.

• Atoms cannot be created or destroyed in chemical reactions.

• Atoms combine in simple whole-number ratios to form compounds.

Structure of the Atom

• Subatomic Particles: Protons (+1, ~1 amu), neutrons (0, ~1 amu), electrons (-1, ~0 amu)

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

• Nuclide Symbol: $^{A}_{Z}\mathrm{X}$, where A = mass number, Z = atomic number, X = element symbol

• Example: $^{14}_{6}\mathrm{C}$ is carbon-14.

Periodic Table Organization

• Main Group Elements: Groups 1, 2, 13–18

• Transition Elements: Groups 3–12

• Metals, Nonmetals, Metalloids: Classified by position and properties

• Family Names: Alkali metals, alkaline earth metals, halogens, noble gases, etc.

• Predicting Ion Charges: Based on group number (e.g., Group 1 forms +1 ions)

Mass Spectrometry

• Used to determine isotopic composition and atomic masses.

• Interpretation involves identifying isotopes and calculating percent/fractional abundances.

Average Atomic Mass Calculation

• Equation: $\text{Average atomic mass} = \sum (\text{fractional abundance} \times \text{isotope mass})$

The Mole and Molar Mass

• Mole: Amount of substance containing Avogadro’s number ($6.022 \times 10^{23}$) of entities.

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

Calculations Involving Atoms

• Conversions:

  • mol ↔ number of atoms: $\text{mol} \times 6.022 \times 10^{23} = \text{number of atoms}$

  • mass ↔ mol: $\text{mass (g)} \div \text{molar mass (g/mol)} = \text{mol}$

  • mass ↔ number of atoms: Combine above steps

Chapter 3: Molecules and Compounds

Types of Bonds and Formulas

• Ionic Compounds: Formed from metals and nonmetals; transfer of electrons.

• Molecular Compounds: Formed from nonmetals; sharing of electrons.

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

• Molecular Formula: Actual number of atoms of each element.

• Structural Formula: Shows connectivity of atoms.

• Example: Glucose: Empirical (CH2O), Molecular (C6H12O6), Structural (shows bonds).

Atoms vs. Molecules

• Atom: Smallest unit of an element.

• Molecule: Two or more atoms bonded together.

Classification of Elements and Compounds

• Elements: Single type of atom

• Compounds: Two or more elements chemically combined

Homonuclear Diatomic and Polyatomic Elements

• Diatomic Elements: H2, N2, O2, F2, Cl2, Br2, I2

• Polyatomic Elements: P4, S8

Chemical Nomenclature

• Rules for naming ionic and molecular compounds, acids, and hydrates.

• Ability to write names from formulas and vice versa.

• Identify formulas from element/ion combinations.

Compound Calculations

• mass ↔ mol: $\text{mass (g)} \div \text{molar mass (g/mol)} = \text{mol}$

• mol ↔ number of molecules/formula units/ions/atoms: $\text{mol} \times 6.022 \times 10^{23}$

• Formula mass: Sum of atomic masses in a compound

• Mass percent (mass%): $\text{mass\%} = \frac{\text{mass of element in 1 mol of compound}}{\text{molar mass of compound}} \times 100\%$

• Determining chemical formula from experimental data (mass%, decomposition, combustion analysis)

• Conversion factors based on mass% and chemical formula

Organic vs. Inorganic Compounds

• Organic Compounds: Contain carbon, often with hydrogen, oxygen, nitrogen, etc.

• Inorganic Compounds: All other compounds.

Families of Organic Compounds and Functional Groups

• Recognize and name functional groups (e.g., alcohols, aldehydes, ketones, carboxylic acids, amines, etc.)

• Name simple, straight-chain hydrocarbons (alkanes, alkenes, alkynes)

Appendix: Key Table – Subatomic Particles

Particle	Symbol	Charge	Approximate Mass (amu)
Proton	p+	+1	1
Neutron	n0	0	1
Electron	e-	-1	~0

Additional info: This study guide synthesizes the main learning objectives and key concepts from the first three chapters of a general chemistry course, providing definitions, examples, and essential equations for exam preparation.