CHEM 102 Test 1 Study Guide: Chapters 1-3 (Matter, Atoms, Compounds)

Chapter 1: Chemistry Basics - Matter and Measurement

Classification of Matter

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

• Pure Substances: Consist of only one type of particle. Examples: Elements (e.g., O2, Fe), Compounds (e.g., H2O, NaCl).

• Mixtures: Contain two or more substances physically combined. Can be separated by physical means.

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

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

The Periodic Table and Its Organization

The periodic table arranges elements by increasing atomic number and similar properties.

• Groups: Vertical columns; elements in a group have similar chemical properties.

• Periods: Horizontal rows; properties change progressively across a period.

• Metals: Good conductors, malleable, ductile (e.g., Fe, Cu).

• Nonmetals: Poor conductors, brittle (e.g., O2, S).

• Metalloids: Properties intermediate between metals and nonmetals (e.g., Si, B).

• Chemical Formulas: Indicate the number and type of atoms in a compound (e.g., H2O has 2 H and 1 O atom).

Changes in Matter and Chemical Equations

• Physical Changes: Do not alter the chemical composition (e.g., melting, boiling).

• Chemical Reactions: Substances are transformed into new substances.

• Balancing Chemical Equations: The number of atoms of each element must be the same on both sides of the equation.

Measurement and Math in Chemistry

• Metric Units: Common units include meter (m), gram (g), liter (L), second (s).

• 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 (e.g., $6.02 \times 10^{23}$).

• Percent Calculations: $\text{Percent} = \frac{\text{Part}}{\text{Whole}} \times 100$

Mass, Volume, Density, and Temperature

• Mass: Amount of matter; SI unit is kilogram (kg).

• Volume: Space occupied; SI unit is liter (L).

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

• Specific Gravity: Ratio of the density of a substance to the density of water.

• Temperature Conversions:

  • Celsius to Kelvin: $K = ^\circ C + 273.15$

  • Celsius to Fahrenheit: $^\circ F = (^\circ C \times \frac{9}{5}) + 32$

  • Fahrenheit to Celsius: $^\circ C = (^\circ F - 32) \times \frac{5}{9}$

• Kinetic vs Potential Energy: Kinetic is energy of motion; potential is stored energy.

• Energy Units: Joule (J), calorie (cal); $1 \text{ cal} = 4.184 \text{ J}$

• Specific Heat: Amount of heat required to raise the temperature of 1 g of a substance by 1°C.

  • Formula: $q = m \times c \times \Delta T$

• States of Matter: Solids (fixed shape/volume), liquids (fixed volume, variable shape), gases (variable shape/volume).

Health Measurements and Unit Conversions

• Accuracy: Closeness to the true value.

• Precision: Reproducibility of measurements.

• SI and U.S. Units: Know how to convert between systems (e.g., 1 inch = 2.54 cm).

• Conversion Factors: Used to convert between units; cancel units as needed.

Chapter 2: Atoms and Radioactivity

Structure of the Atom

Atoms are composed of subatomic particles: protons, neutrons, and electrons.

• Protons: Positive charge, located in the nucleus.

• Neutrons: No charge, located in the nucleus.

• Electrons: Negative charge, located in orbitals around the nucleus.

Atomic Number, Mass Number, and Isotopes

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

• Mass Number (A): Number of protons plus neutrons.

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

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

• Symbolic Notation: $^{A}_{Z}\text{X}$, where X is the element symbol.

Radioactivity and Types of Ionizing Radiation

• Radioactivity: Spontaneous emission of particles or energy from unstable nuclei.

• Types of Radiation:

  • Alpha (α): 2 protons + 2 neutrons; low penetration.

  • Beta (β): Electron emission; moderate penetration.

  • Gamma (γ): High-energy photons; high penetration.

  • Positron (β+): Positive electron; used in PET scans.

Nuclear Equations and Decay

• Balancing Nuclear Equations: Mass numbers and atomic numbers must be conserved.

• Example: $^{14}_{6}\text{C} \rightarrow ^{14}_{7}\text{N} + ^{0}_{-1}\beta$ (beta decay)

Radiation Units and Half-Life

• Units of Radioactivity: Becquerel (Bq), Curie (Ci).

• Units of Biological Damage: Sievert (Sv), rem.

• Half-Life (t1/2): Time for half of a radioactive sample to decay.

• Half-Life Formula: $N = N_0 \times (\frac{1}{2})^{\frac{t}{t_{1/2}}}$

Medical Applications of Radioactivity

• Diagnosis: Radioisotopes used in imaging (e.g., PET scans).

• Treatment: Radiation therapy for cancer.

Chapter 3: Compounds - How Elements Combine

Electron Arrangement and the Octet Rule

• Electron Shells: Electrons occupy energy levels (shells) around the nucleus.

• Valence Electrons: Electrons in the outermost shell; determine chemical reactivity.

• Octet Rule: Atoms tend to gain, lose, or share electrons to achieve 8 valence electrons.

Ion Formation and Polyatomic Ions

• Ions: Charged particles formed by gaining or losing electrons.

• Cations: Positively charged (loss of electrons).

• Anions: Negatively charged (gain of electrons).

• Polyatomic Ions: Groups of atoms with a net charge (e.g., SO42−).

Ionic Compounds

• Formation: Combination of cations and anions to form neutral compounds.

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

• Transition Metals: May have variable charges; use Roman numerals (e.g., FeCl2: iron(II) chloride).

Covalent Compounds and Lewis Structures

• Covalent Bonds: Atoms share electrons to achieve octet.

• Lewis Structures: Diagrams showing bonding and lone pairs.

• Naming Binary Covalent Compounds: Use prefixes (e.g., CO2: carbon dioxide).

The Mole and Avogadro's Number

• Mole: Amount of substance containing $6.022 \times 10^{23}$ particles.

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

• Conversions:

  • Moles to grams: $\text{mass (g)} = \text{moles} \times \text{molar mass (g/mol)}$

  • Moles to particles: $\text{particles} = \text{moles} \times 6.022 \times 10^{23}$

Molecular Shape and Polarity

• VSEPR Theory: Predicts 3D shape of molecules based on electron pair repulsion.

• Wedge-and-Dash Notation: Shows 3D arrangement of atoms.

• Bond Polarity: Determined by difference in electronegativity between atoms.

• Molecular Polarity: Depends on bond polarities and molecular shape.

Table: Comparison of Ionic and Covalent Compounds

Property	Ionic Compounds	Covalent Compounds
Bond Type	Transfer of electrons	Sharing of electrons
Constituents	Metals + Nonmetals	Nonmetals
Physical State	Usually solid	Solid, liquid, or gas
Melting Point	High	Low to moderate
Electrical Conductivity	Conducts when molten or dissolved	Usually does not conduct

Example: Calculating Moles and Mass

• How many grams are in 2.00 moles of NaCl? Molar mass of NaCl = 58.44 g/mol.

• Calculation: $2.00 \text{ mol} \times 58.44 \text{ g/mol} = 116.88 \text{ g}$

Additional info: Some explanations and examples have been expanded for clarity and completeness based on standard general, organic, and biological chemistry curricula.