General Chemistry I: Comprehensive Study Notes (Chapters 1–11)

Chapter 1: Chemical Tools – Experimentation & Measurement

SI Units and Prefixes

The International System of Units (SI) is the standard for scientific measurement. Prefixes are used to denote multiples or fractions of base units.

• Base Units: meter (m), kilogram (kg), second (s), mole (mol), kelvin (K), ampere (A), candela (cd)

• Common Prefixes:

  • kilo- (k): $10^3$

  • centi- (c): $10^{-2}$

  • milli- (m): $10^{-3}$

  • micro- (\mu): $10^{-6}$

  • nano- (n): $10^{-9}$

Dimensional Analysis

Dimensional analysis is a method for converting between units using conversion factors.

• Conversion Factor: A ratio that expresses how many of one unit are equal to another unit.

• Example: To convert 5.0 cm to meters: $5.0\ \text{cm} \times \frac{1\ \text{m}}{100\ \text{cm}} = 0.050\ \text{m}$

Significant Figures

Significant figures reflect the precision of a measured quantity.

• All nonzero digits are significant.

• Zeros between nonzero digits are significant.

• Leading zeros are not significant; trailing zeros are significant only if there is a decimal point.

Chapter 2: Atoms, Molecules & Ions

Properties of Matter

• Matter: Anything that has mass and occupies space.

• Physical Properties: Characteristics that can be observed without changing the substance (e.g., melting point, density).

• Chemical Properties: Characteristics that describe a substance's ability to change into different substances.

Atomic Structure

• Atoms: The smallest unit of an element that retains its identity in a chemical reaction.

• Subatomic Particles: Protons (+), neutrons (0), electrons (−)

• Atomic Number (Z): Number of protons in the nucleus.

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

Isotopes

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

• Example: $^{12}\text{C}$ and $^{14}\text{C}$ are isotopes of carbon.

Compounds and Formulas

• Molecular Formula: Shows the exact number of atoms of each element in a molecule.

• Empirical Formula: Shows the simplest whole-number ratio of atoms.

Chapter 3: Mass Relationships in Chemical Reactions

Balancing Chemical Equations

Chemical equations must be balanced to obey the law of conservation of mass.

• Adjust coefficients to ensure the same number of each atom on both sides.

Stoichiometry

• Relates quantities of reactants and products in a chemical reaction.

• Mole Concept: $1\ \text{mol} = 6.022 \times 10^{23}$ particles (Avogadro's number)

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

• Example: To find grams of product formed from a given amount of reactant, use:

  • Convert grams to moles

  • Use mole ratio from balanced equation

  • Convert moles to grams

Limiting Reactant and Percent Yield

• Limiting Reactant: The reactant that is completely consumed first, limiting the amount of product formed.

• Theoretical Yield: Maximum amount of product possible.

• Percent Yield: $\frac{\text{Actual Yield}}{\text{Theoretical Yield}} \times 100\%$

Chapter 4: Reactions in Aqueous Solution

Types of Chemical Reactions

• Precipitation Reactions: Formation of an insoluble product (precipitate).

• Acid-Base Reactions: Transfer of protons (H+).

• Redox Reactions: Transfer of electrons.

Solubility Rules

• Used to predict whether a precipitate will form in a reaction.

Net Ionic Equations

• Show only the species that actually change during the reaction.

Chapter 5: Periodicity & Electronic Structure of Atoms

Periodic Trends

• Atomic Radius: Decreases across a period, increases down a group.

• Ionization Energy: Increases across a period, decreases down a group.

• Electron Affinity: Generally becomes more negative across a period.

• Electronegativity: Tendency of an atom to attract electrons in a bond.

Quantum Numbers

• Describe the energy and shape of atomic orbitals.

• Principal (n), angular momentum (l), magnetic (ml), spin (ms).

Electron Configurations

• Filling order: Aufbau principle, Pauli exclusion principle, Hund's rule.

• Example: $1s^2 2s^2 2p^6 3s^2 3p^6 4s^1$ for potassium (K).

Chapter 6: Ionic Compounds – Periodic Trends and Bonding Theory

Ionic Bonding

• Formed by transfer of electrons from metals to nonmetals.

• Lattice Energy: Energy required to separate one mole of an ionic solid into gaseous ions.

Properties of Ionic Compounds

• High melting and boiling points, conduct electricity when molten or dissolved in water.

Chapter 7: Covalent Bonding and Electron-Dot Structures

Covalent Bonding

• Sharing of electron pairs between atoms.

• Bond Length and Strength: Shorter bonds are stronger; multiple bonds (double, triple) are shorter and stronger than single bonds.

Lewis Structures

• Show arrangement of valence electrons in molecules.

• Follow the octet rule for main group elements.

Resonance

• Some molecules can be represented by two or more valid Lewis structures.

Chapter 8: Covalent Compounds – Bonding Theories and Molecular Structure

VSEPR Theory

• Valence Shell Electron Pair Repulsion (VSEPR) theory predicts molecular shapes based on electron pair repulsion.

• Example: $\text{CH}_4$ is tetrahedral, $\text{NH}_3$ is trigonal pyramidal.

Bond Polarity and Molecular Polarity

• Bond polarity arises from differences in electronegativity.

• Molecular polarity depends on both bond polarity and molecular shape.

Hybridization

• Atomic orbitals mix to form new hybrid orbitals (e.g., $sp^3$, $sp^2$, $sp$).

Chapter 9: Thermochemistry – Chemical Energy

Energy Changes in Chemical Reactions

• System vs. Surroundings: The system is the part of the universe being studied; everything else is the surroundings.

• Endothermic: Absorbs heat; Exothermic: Releases heat.

First Law of Thermodynamics

• Energy cannot be created or destroyed, only transferred or converted.

• $\Delta E = q + w$ (change in internal energy = heat + work)

Enthalpy

• $\Delta H$ is the heat change at constant pressure.

• Standard enthalpy of formation ($\Delta H_f^\circ$): Enthalpy change for forming 1 mol of a compound from its elements in their standard states.

Calorimetry

• Measurement of heat flow.

• $q = mc\Delta T$ (heat = mass × specific heat × temperature change)

Chapter 10: Gases – Their Properties & Behavior

Gas Laws

• Boyle's Law: $P_1V_1 = P_2V_2$ (at constant T and n)

• Charles's Law: $\frac{V_1}{T_1} = \frac{V_2}{T_2}$ (at constant P and n)

• Avogadro's Law: $\frac{V_1}{n_1} = \frac{V_2}{n_2}$ (at constant P and T)

• Ideal Gas Law: $PV = nRT$

Partial Pressures and Dalton's Law

• Total pressure of a gas mixture is the sum of the partial pressures of each component.

• $P_{\text{total}} = P_1 + P_2 + \ldots$

Kinetic Molecular Theory

• Explains gas behavior in terms of particle motion.

• Assumes particles are in constant, random motion and collisions are elastic.

Chapter 11: Liquids & Phase Changes

Properties of Liquids

• Intermolecular forces: dispersion, dipole-dipole, hydrogen bonding.

• Viscosity, surface tension, vapor pressure.

Phase Changes

• Heat of fusion (melting), heat of vaporization (boiling).

• Calculating heat for temperature changes: $q = mc\Delta T$

• Calculating heat for phase changes: $q = n\Delta H$

Phase Diagrams

• Show the state of a substance at various temperatures and pressures.

• Triple point: All three phases coexist.

• Critical point: End of the liquid-gas boundary.

Summary Table: SI Prefixes

Prefix	Symbol	Multiplier
kilo-	k	$10^3$
centi-	c	$10^{-2}$
milli-	m	$10^{-3}$
micro-	μ	$10^{-6}$
nano-	n	$10^{-9}$

Additional info: These notes are structured to cover the main learning objectives and foundational concepts for a first-semester General Chemistry course, as outlined in the provided syllabus. For each chapter, students should practice applying these concepts to problems and laboratory scenarios.