General Chemistry I: Chapters 1–11 Comprehensive Study Notes

Chapter 1 – Matter, Energy, and Measurement

Classification of Matter

• Matter is anything that has mass and occupies space.

• It can be classified as a pure substance or a mixture:

  • Pure substances have a fixed composition and distinct properties. They can be elements or compounds.

  • Mixtures are combinations of two or more substances where each retains its own identity. They can be:

    • Homogeneous mixtures (solutions): uniform throughout (e.g., saltwater).

    • Heterogeneous mixtures: not uniform throughout (e.g., salad).

• Flowchart for Classification:

Type	Uniformity	Separation
Heterogeneous Mixture	No	Physical methods
Homogeneous Mixture (Solution)	Yes	Physical methods
Pure Substance: Element	Yes	Chemical methods not possible
Pure Substance: Compound	Yes	Chemical methods possible

Properties of Matter

• Intensive Properties: Do not depend on the amount of substance (e.g., melting point, boiling point, color, density).

• Extensive Properties: Depend on the amount of substance (e.g., mass, volume).

Precision vs. Accuracy

• Precision: How closely repeated measurements agree with each other.

• Accuracy: How closely a measurement agrees with the true value.

• Example: Hitting the same spot on a target repeatedly is precise; hitting the bullseye is accurate.

Significant Figures (Sig Figs)

• Rules for determining significant figures:

  1. All nonzero digits are significant (e.g., 123.5 has 4 sig figs).

  2. Zeros between nonzero digits are significant (e.g., 2500.3 has 5 sig figs).

  3. Zeros at the beginning are not significant (e.g., 0.02 has 1 sig fig).

  4. Zeros at the end are significant if there is a decimal point (e.g., 0.0200 has 3 sig figs).

  5. For numbers ending in zeros with no decimal, use scientific notation to indicate significance (e.g., 2.00 × 102 has 3 sig figs).

• Operations:

  • Addition/Subtraction: Result has the same number of decimal places as the measurement with the fewest decimal places.

  • Multiplication/Division: Result has the same number of significant figures as the measurement with the fewest significant figures.

Chapter 2 – Atoms, Molecules, and Ions

Atomic Symbols and Structure

• Elements are represented by one- or two-letter symbols (e.g., C for carbon).

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

• Mass number (A): Total number of protons and neutrons.

• For ions: Number of electrons = Atomic number – Charge.

Subatomic Particles

Particle	Charge	Mass (amu)
Proton	+1	1.007
Neutron	0	1.008
Electron	-1	5.486 × 10-4

Atomic Weight

• Weighted average of all isotopes of an element based on their natural abundance.

• Formula:

$\text{Atomic Weight} = \sum (\text{isotope mass} \times \text{fractional natural abundance})$

• All atomic masses are compared to C-12 (exactly 12 amu).

Types of Formulas

• Empirical formula: Lowest whole-number ratio of atoms in a compound.

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

• If the molecular formula is known, the empirical formula can be determined, but not vice versa without more information.

Chapter 3 – Stoichiometry and Chemical Reactions

Stoichiometric Calculations

• Convert between mass, moles, and number of particles using Avogadro's number ($6.022 \times 10^{23}$ particles/mol).

• Calculate formula weight by summing atomic masses of all atoms in the formula.

• Balance chemical equations to obey the law of conservation of mass.

• Use stoichiometry to determine the amount of reactants or products in a reaction.

• Percent composition: $\%\text{Element} = \frac{\text{mass of element in 1 mol compound}}{\text{molar mass of compound}} \times 100$

Chapter 4 – Reactions in Aqueous Solution

Types of Aqueous Reactions

• Precipitation, acid-base, and redox reactions are common in aqueous solutions.

• Solubility guidelines help predict if a precipitate will form.

Soluble Compounds	Important Exceptions
NO3-, CH3COO-, Cl-, Br-, I-, SO42-	Ag+, Pb2+, Hg22+ (for halides); Ba2+, Sr2+, Pb2+ (for sulfates)
Insoluble Compounds	Important Exceptions
CO32-, PO43-, S2-, OH-	Alkali metals, NH4+

Oxidation-Reduction (Redox) Chemistry

• Oxidation: Loss of electrons.

• Reduction: Gain of electrons.

• Oxidation numbers are assigned to track electron transfer.

• Redox activity series helps predict which metals will be oxidized or reduced.

Chapter 5 – Thermochemistry

Enthalpy and Thermochemical Calculations

• Enthalpy (H): Heat content of a system at constant pressure.

• Endothermic: Absorbs heat ($\Delta H > 0$); Exothermic: Releases heat ($\Delta H < 0$).

• Hess's Law: The enthalpy change for a reaction is the same, no matter how many steps the reaction is carried out in.

• Calculate enthalpy change using standard enthalpies of formation:

$\Delta H^\circ_{rxn} = \sum n\Delta H^\circ_f(\text{products}) - \sum n\Delta H^\circ_f(\text{reactants})$

Chapter 6 – Electronic Structure of Atoms

Light and Atomic Structure

• Energy of a photon: $E = h\nu$

• Relationship between wavelength and frequency: $c = \lambda \nu$

• Where $h$ is Planck's constant ($6.626 \times 10^{-34}$ J·s), $c$ is the speed of light ($3.00 \times 10^8$ m/s), $\lambda$ is wavelength, and $\nu$ is frequency.

Quantum Numbers and Electron Configuration

• Quantum numbers describe the properties of atomic orbitals and electrons:

  • Principal ($n$), angular momentum ($l$), magnetic ($m_l$), and spin ($m_s$).

• Electron configurations show how electrons fill atomic orbitals (Aufbau principle, Pauli exclusion, Hund's rule).

Chapter 7 – Periodic Properties of the Elements

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.

Ion Properties

• Cations are smaller, anions are larger than their parent atoms.

• Isoelectronic series: Species with the same number of electrons; size decreases with increasing nuclear charge.

Chapter 8 – Basic Concepts of Chemical Bonding

Ionic vs. Covalent Bonds

• Ionic bonds: Transfer of electrons from metal to nonmetal.

• Covalent bonds: Sharing of electrons between nonmetals.

• Single, double, and triple bonds differ in strength and length (triple > double > single in strength; single > double > triple in length).

Lewis Structures

• Show arrangement of valence electrons in molecules.

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

• Resonance structures: Multiple valid Lewis structures for a molecule; actual structure is a hybrid.

Chapter 9 – Molecular Geometry and Bonding Theories

VSEPR Theory and Molecular Shape

• Electron domains (bonding and lone pairs) determine geometry.

• Common geometries: linear, trigonal planar, tetrahedral, trigonal bipyramidal, octahedral.

• Bond angles depend on geometry (e.g., tetrahedral = 109.5°).

• Hybridization: Mixing of atomic orbitals to form new hybrid orbitals (e.g., sp, sp2, sp3).

• Molecular polarity depends on geometry and bond dipoles.

Chapter 10 – Gases

Gas Laws and Calculations

• Properties: Gases have indefinite shape and volume, are compressible, and mix completely.

• Ideal Gas Law: $PV = nRT$

• Combined Gas Law: $\frac{P_1V_1}{T_1} = \frac{P_2V_2}{T_2}$

• Partial pressure: $P_A = X_A P_{total}$, where $X_A$ is the mole fraction.

• Density: $d = \frac{PM}{RT}$, where $M$ is molar mass.

Chapter 11 – Liquids and Intermolecular Forces

Types of Intermolecular Forces

• London Dispersion: Present in all molecules, especially nonpolar.

• Dipole-Dipole: Between polar molecules.

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

• Ion-Dipole: Between ions and polar molecules.

Effects of Intermolecular Forces

• Stronger intermolecular forces lead to higher boiling points.

• Structural features (polarity, presence of H-bond donors/acceptors) determine the type and strength of forces.