General Chemistry I: Final Exam Study Guide – Matter, Measurement, Atomic Structure, Periodicity, Bonding, and Stoichiometry

Matter & Measurement Techniques

Significant Figures

Significant figures (sig figs) are the digits in a measurement that are known with certainty plus one digit that is estimated. Correct use of significant figures ensures that calculated results reflect the precision of the measurements used.

• Nonzero digits are always significant.

• Any zeros between significant digits are significant.

• Leading zeros are not significant.

• Trailing zeros in a number with a decimal point are significant.

Example: 0.00450 has three significant figures (4, 5, and the trailing 0).

Rounding and Calculations with Significant Figures

When performing calculations, the number of significant figures in the result depends on the operation:

• Addition/Subtraction: Round the result to the same number of decimal places as the measurement with the fewest decimal places.

• Multiplication/Division: Round the result to the same number of significant figures as the measurement with the fewest significant figures.

Example: 2.54 cm × 1.2 cm = 3.0 cm2 (rounded to two significant figures).

Temperature Conversion

Temperature can be converted between Fahrenheit (°F) and Celsius (°C) using the following formulas:

• From Fahrenheit to Celsius: $T(°C) = \frac{5}{9} [T(°F) - 32]$

• From Celsius to Fahrenheit: $T(°F) = \frac{9}{5} T(°C) + 32$

Unit Conversion (Dimensional Analysis)

Dimensional analysis is used to convert between units using conversion factors.

• Example: To convert 25.4 mm to inches, use the conversion factor 1 in = 25.4 mm.

• $25.4\ \text{mm} \times \frac{1\ \text{in}}{25.4\ \text{mm}} = 1\ \text{in}$

Separation Techniques

Separation techniques are used to separate mixtures based on physical properties.

• Distillation: Used to separate components of a homogeneous liquid mixture (e.g., water and alcohol) based on differences in boiling points.

• Filtration: Used to separate solids from liquids in heterogeneous mixtures.

Phases and Phase Changes

Matter exists in different phases: solid, liquid, and gas. Phase changes occur when matter transitions between these states.

• Melting: Solid to liquid

• Freezing: Liquid to solid

• Vaporization: Liquid to gas

• Condensation: Gas to liquid

• Sublimation: Solid to gas

• Deposition: Gas to solid

Classification of Matter

Matter can be classified as:

• Pure Element: Substance made of only one type of atom (e.g., O2).

• Pure Compound: Substance made of two or more elements chemically combined (e.g., H2O).

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

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

Solubility Rules

Solubility rules help predict whether an ionic compound will dissolve in water. Refer to your periodic table for common rules (e.g., most nitrates are soluble, most chlorides are soluble except AgCl, PbCl2, etc.).

Physical vs. Chemical Change

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

• Chemical Change: Alters the chemical composition (e.g., combustion, rusting).

Laboratory Glassware

Common laboratory glassware includes beakers, Erlenmeyer flasks, graduated cylinders, burettes, and pipettes.

Safety Symbols (WHMIS Pictograms)

WHMIS pictograms identify hazards associated with chemicals. Examples include:

• Flame: Flammable materials

• Corrosion: Corrosive to metals and skin

• Skull and Crossbones: Acute toxicity

• Exclamation Mark: Irritant or less severe toxicity

Additional info: Students should be able to match pictograms to their meanings and understand basic laboratory safety.

Atomic Structure and Quantum Theory

Dalton's Atomic Theory

Dalton's atomic theory laid the foundation for modern chemistry by proposing that matter is composed of indivisible atoms, each element has unique atoms, and atoms combine in fixed ratios to form compounds.

Isotopes and Atomic Mass

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

• Atomic Mass: Weighted average of all naturally occurring isotopes.

Example: Carbon-12 and Carbon-14 are isotopes of carbon.

Electromagnetic Radiation

Electromagnetic radiation includes visible light, X-rays, and radio waves. Key properties include:

• Wavelength (λ): Distance between successive peaks.

• Frequency (ν): Number of cycles per second.

• Speed of light (c): $c = \lambda \nu$

Quantum Numbers

Quantum numbers describe the properties of atomic orbitals and electrons:

• Principal quantum number (n): Energy level (1, 2, 3, ...)

• Angular momentum quantum number (l): Shape of orbital (0 = s, 1 = p, 2 = d, 3 = f)

• Magnetic quantum number (ml): Orientation of orbital (−l to +l)

• Spin quantum number (ms): Electron spin (+½ or −½)

Quantum Number Rules

• For a given n, l can be 0 to n−1.

• For a given l, ml can be −l to +l.

• Each orbital can hold two electrons with opposite spins.

Orbitals and Subshells

Know the maximum number of electrons in each subshell:

• s: 2 electrons

• p: 6 electrons

• d: 10 electrons

• f: 14 electrons

Electron Configuration

Write the long and short (noble gas) electron configurations for elements. Use the Aufbau principle and Hund's rule to fill orbitals:

• Aufbau Principle: Fill lowest energy orbitals first.

• Hund's Rule: Fill degenerate orbitals singly before pairing electrons.

Bonding and Periodicity

Nomenclature

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

• Covalent Compounds: Use prefixes to indicate number of atoms (e.g., CO2 = carbon dioxide).

Lewis Structures

Draw Lewis structures to show valence electrons and bonding in molecules and polyatomic ions.

Periodic Trends

• Atomic Radius: Increases down a group, decreases across a period.

• Ionization Energy: Decreases down a group, increases across a period.

• Electronegativity: Decreases down a group, increases across a period.

Intermolecular Forces (IMFs)

• London Dispersion Forces: Present in all molecules, stronger in larger/heavier molecules.

• Dipole-Dipole Forces: Present in polar molecules.

• Hydrogen Bonding: Strongest IMF, occurs when H is bonded to F, O, or N.

Physical Properties: Stronger IMFs lead to higher melting and boiling points.

Chemical Reactions and Equations

Balancing Chemical Equations

Ensure the same number of each atom on both sides of the equation. Use coefficients to balance.

Types of Chemical Reactions

• Combination (Synthesis)

• Decomposition

• Single Displacement

• Double Displacement

• Combustion

Stoichiometry and Chemical Formulas

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

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

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

Example: Empirical formula of C6H12O6 is CH2O.

Limiting Reactant and Yield

• Identify the limiting reactant (the reactant that is completely consumed first).

• Calculate theoretical yield (maximum possible amount of product).

• Calculate actual yield (amount actually obtained).

• Calculate percent yield: $\text{Percent Yield} = \frac{\text{Actual Yield}}{\text{Theoretical Yield}} \times 100\%$

Titration and Molarity

• Molarity (M): $M = \frac{\text{moles of solute}}{\text{liters of solution}}$

• Use titration data to calculate concentration of analyte.

• Use the formula: $M_1V_1 = M_2V_2$ for dilution and titration calculations.

Appendix: Common Tables

Phase Change	Transition
Melting	Solid to Liquid
Freezing	Liquid to Solid
Vaporization	Liquid to Gas
Condensation	Gas to Liquid
Sublimation	Solid to Gas
Deposition	Gas to Solid

Quantum Number	Symbol	Meaning
Principal	n	Energy level
Angular Momentum	l	Shape of orbital
Magnetic	ml	Orientation of orbital
Spin	ms	Electron spin

Intermolecular Force	Occurs In	Relative Strength
London Dispersion	All molecules	Weakest
Dipole-Dipole	Polar molecules	Intermediate
Hydrogen Bonding	H bonded to F, O, or N	Strongest