GOB Chemistry Exam 1 Study Guide: Chapters 1, 2, 3, 4, & 6

Chapter 1 – Chemistry & Problem Solving

Scientific Method

The scientific method is a systematic approach to understanding the natural world through observation, hypothesis formation, experimentation, and the development of laws and theories.

• Observation: Gathering information through the senses or measurement. Example: "The solution turned blue."

• Hypothesis: A testable explanation or prediction. Example: "The solution turns blue because substance X reacts with it."

• Experiment: A controlled procedure to test a hypothesis.

• Law: A statement describing a consistent pattern in nature (often mathematical).

• Theory: A well-supported explanation of why something happens, based on evidence.

Example: If a student observes that a plant grows taller near a window (observation) and predicts that sunlight causes this (hypothesis), she can design an experiment to test this idea.

Strategies for Learning Chemistry

• Review material regularly and avoid cramming.

• Practice solving problems and check your units and significant figures.

• Understand concepts before memorizing details.

Scientific Notation

Scientific notation expresses numbers as $a \times 10^n$, where $1 \leq a < 10$.

• Large numbers: positive exponent (e.g., $45,000 = 4.5 \times 10^4$)

• Small decimals: negative exponent (e.g., $0.00072 = 7.2 \times 10^{-4}$)

Percentages

Percentages are calculated as:

$\% = \frac{\text{part}}{\text{whole}} \times 100$

Example: If 15 out of 20 questions are correct: $\frac{15}{20} \times 100 = 75\%$

Decimal, Fraction, and Percentage Conversions

• To convert a decimal to a fraction: 0.375 = $\frac{3}{8}$

• To convert a decimal to a percentage: 0.375 = 37.5%

• To convert a percentage to a decimal: 65% = 0.65

Chapter 2 – Measurements & Calculations

Exact vs. Measured Numbers

• Exact numbers: Have no uncertainty (e.g., 12 eggs in a dozen).

• Measured numbers: Obtained by measurement and have some uncertainty (e.g., 12.5 g).

Significant Figures

• Nonzero digits are always significant.

• Zeros between nonzero digits are significant.

• Leading zeros are not significant.

• Trailing zeros after a decimal are significant.

• Trailing zeros without a decimal can be ambiguous; use scientific notation for clarity.

Multiplication/Division: The answer has the same number of significant figures as the measurement with the fewest significant figures. Addition/Subtraction: The answer has the same number of decimal places as the measurement with the fewest decimal places.

Dimensional Analysis

Dimensional analysis uses conversion factors to change units:

• Write the given value.

• Multiply by conversion factors so units cancel appropriately.

• Calculate and check units.

Example: Convert 2.5 m to cm: $2.5\,\text{m} \times \frac{100\,\text{cm}}{1\,\text{m}} = 250\,\text{cm}$

Metric & SI Units

• Length: meter (m)

• Mass: kilogram (kg)

• Volume: liter (L)

• Time: second (s)

• Temperature: kelvin (K)

Common conversions: $1\,\text{m} = 100\,\text{cm}$, $1\,\text{kg} = 1000\,\text{g}$, $1\,\text{L} = 1000\,\text{mL}$

Density

Density is defined as mass per unit volume:

$D = \frac{m}{V}$

Rearranged:

• $m = D \times V$

• $V = \frac{m}{D}$

Example: If mass = 20.0 g and volume = 5.0 mL, $D = \frac{20.0\,g}{5.0\,mL} = 4.0\,g/mL$

Chapter 3 – Matter & Energy

Classification of Matter

• Pure substances: Constant composition (elements or compounds).

• Element: One type of atom (e.g., Fe, O2).

• Compound: Two or more elements chemically bonded (e.g., H2O, CO2).

• Mixture: Physical combination of substances (homogeneous or heterogeneous).

States of Matter

• Solid: Fixed shape and volume; particles vibrate in place.

• Liquid: Fixed volume, takes shape of container; particles move past each other.

• Gas: No fixed volume or shape; particles move freely.

Physical vs. Chemical Changes

• Physical change: Identity does not change (e.g., melting ice).

• Chemical change: New substance formed (e.g., burning wood).

Temperature Scales

• Celsius (°C), Fahrenheit (°F), Kelvin (K)

Conversions:

• $°F = \frac{9}{5}(°C) + 32$

• $°C = \frac{5}{9}(°F - 32)$

• $K = °C + 273.15$

Changes of State

• Melting: solid → liquid

• Freezing: liquid → solid

• Evaporation/Boiling: liquid → gas

• Condensation: gas → liquid

• Sublimation: solid → gas

• Deposition: gas → solid

Calories & Joules

• 1 cal = 4.184 J

• 1 Cal (nutritional Calorie) = 1 kcal = 1000 cal = 4184 J

Chapter 4 – Atoms & the Periodic Table

Groups & Periods

• Group: Vertical column (e.g., Group 1 = alkali metals)

• Period: Horizontal row

• Elements in the same group have similar chemical properties.

Metals, Nonmetals, and Metalloids

• Metals: Shiny, conduct electricity, malleable, ductile (left/center of table).

• Nonmetals: Poor conductors, brittle, many are gases (upper right).

• Metalloids: Properties between metals and nonmetals (along stair-step line).

Subatomic Particles

Particle	Charge	Location	Approx. Mass
Proton	+1	Nucleus	1 amu
Neutron	0	Nucleus	1 amu
Electron	−1	Outside nucleus	~0 amu

Atomic Number & Mass Number

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

• Mass number (A): Protons + neutrons.

• Neutrons = A − Z

Example: Atomic number = 17, Mass number = 37 → Protons = 17, Neutrons = 20

Isotopes

• Atoms of the same element with different numbers of neutrons (different mass numbers).

• Isotope notation: $^{A}_{Z}X$

Example: Carbon-14: 6 protons, 8 neutrons

Chapter 6 – Chemical Compounds & Bonding

Types of Compounds

• Element: Only one type of atom (e.g., Fe, O2).

• Ionic compound: Metal + nonmetal or contains polyatomic ions (e.g., NaCl).

• Molecular (covalent) compound: Nonmetal + nonmetal (e.g., CO2).

Ions

• Cation: Positive ion (atom loses electrons).

• Anion: Negative ion (atom gains electrons).

Example: $\mathrm{Na} \rightarrow \mathrm{Na}^+ + e^-$ (cation); $\mathrm{Cl} + e^- \rightarrow \mathrm{Cl}^-$ (anion)

Fixed vs. Variable Charges

• Group 1 metals: +1 charge

• Group 2 metals: +2 charge

• Some metals (e.g., Fe, Cu, Sn) have variable charges; use Roman numerals in names.

Naming Ionic Compounds

• Fixed-charge metal: Name = metal + anion Example: sodium chloride ($\mathrm{NaCl}$)

• Variable-charge metal: Name = metal (Roman numeral) + anion Example: iron(II) chloride ($\mathrm{FeCl_2}$), iron(III) chloride ($\mathrm{FeCl_3}$)

Polyatomic Ions

• A polyatomic ion is a group of atoms with an overall charge.

• Keep the polyatomic ion's name in the compound name. Example: sodium nitrate ($\mathrm{NaNO_3}$)

Writing Molecular Compound Formulas

• Prefixes indicate the number of atoms (e.g., CO2 is carbon dioxide).

• Common prefixes: mono-, di-, tri-, tetra-, penta-, hexa-, hepta-, octa-, nona-, deca-

Diatomic Molecules

• Seven elements naturally exist as diatomic molecules: H2, N2, O2, F2, Cl2, Br2, I2

• Memory aid: "Have No Fear Of Ice Cold Beer"

Electronegativity

• Electronegativity is an atom's ability to attract shared electrons in a bond.

• Increases left to right and bottom to top on the periodic table; fluorine is the most electronegative.

Polar vs. Nonpolar Covalent Bonds

• Nonpolar: Electrons shared equally; small or zero electronegativity difference.

• Polar: Electrons shared unequally; larger electronegativity difference.

• Difference: $\Delta EN = |EN_1 - EN_2|$

Intermolecular Attractions

• London Dispersion Forces: Present in all molecules; stronger in larger/more polarizable particles.

• Dipole-Dipole: Between polar molecules; positive end of one attracts negative end of another.

• Hydrogen Bonding: Strong dipole-dipole attraction when H is bonded to N, O, or F.

Relative strength: Hydrogen bonding > dipole-dipole > dispersion forces (for comparable molecules).

Key Formulas to Know

• $D = \frac{m}{V}$

• $m = D \times V$

• $V = \frac{m}{D}$

• $A = p + n$

• $n = A - Z$

• $Z = p$

• $°F = \frac{9}{5}(°C) + 32$

• $°C = \frac{5}{9}(°F - 32)$

• $K = °C + 273.15$

• $1\,\text{cal} = 4.184\,\text{J}$

• $1\,\text{Cal} = 1000\,\text{cal}$

Highest-Priority Exam Skills

• Solve: significant figures, rounding, dimensional analysis, density, temperature, and energy conversions, atomic structure, writing/naming compounds, bond polarity.

• Identify/explain: scientific method, types of numbers, matter classification, changes, periodic table groups, subatomic particles, isotopes, ions, bonding, polyatomic ions, electronegativity, intermolecular forces.