General Chemistry I: Comprehensive Study Notes and Reference Tables

Atoms, Elements, and the Periodic Table

Atomic Structure and the Periodic Table

The periodic table organizes elements based on atomic number, electron configuration, and recurring chemical properties. Understanding the arrangement of elements is fundamental to predicting their behavior in chemical reactions.

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

• Mass Number (A): Sum of protons and neutrons in the nucleus.

• Groups and Periods: Vertical columns are groups (similar chemical properties); horizontal rows are periods (increasing atomic number).

• Metals, Nonmetals, Metalloids: Elements are classified based on their physical and chemical properties.

Example: Sodium (Na) is in Group 1 (alkali metals), Period 3, atomic number 11.

Chemical Bonding and Molecular Structure

Lewis Structures and Molecular Geometry

Lewis structures represent the arrangement of valence electrons in molecules. Molecular geometry describes the three-dimensional arrangement of atoms.

• Valence Electrons: Electrons in the outermost shell involved in bonding.

• Lewis Structure: Diagram showing bonds and lone pairs in a molecule.

• Electron Domain Geometry: Arrangement of electron domains (bonding and lone pairs) around the central atom.

• Molecular Geometry: Arrangement of atoms (ignoring lone pairs).

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

• Polarity: A molecule is polar if it has a net dipole moment due to unequal sharing of electrons.

Example: For ICl3, the Lewis structure shows 5 electron domains (3 bonds, 2 lone pairs), leading to a T-shaped molecular geometry and sp3d hybridization.

Chemical Reactions and Stoichiometry

Balancing Chemical Equations

Chemical equations must be balanced to obey the law of conservation of mass. Stoichiometry involves calculations based on balanced equations to determine the amounts of reactants and products.

• Balancing: Adjust coefficients to have equal numbers of each atom on both sides.

• Stoichiometry: Use mole ratios from the balanced equation to calculate quantities.

Example: Balancing the reaction: HClO3 (aq) + Cl2 (g) → HCl (aq) + ClO2- (aq) + Cl- (aq).

Chemical Quantities and Aqueous Reactions

Solution Concentrations and Ionic Equations

Solutions are homogeneous mixtures. Concentration is often expressed in molarity (mol/L). Ionic equations show the species present in solution.

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

• Molecular Equation: Shows all reactants and products as compounds.

• Complete Ionic Equation: Shows all strong electrolytes as ions.

• Net Ionic Equation: Shows only the species that change during the reaction.

Example: Mixing K3PO4 (aq) and AgNO3 (aq) forms Ag3PO4 (s).

Gases and Gas Laws

Properties of Gases and Gas Law Calculations

Gases have unique properties described by several laws. The ideal gas law relates pressure, volume, temperature, and amount of gas.

• Ideal Gas Law: $PV = nRT$

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

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

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

• Dalton's Law of Partial Pressures: $P_{\text{total}} = P_1 + P_2 + \ldots$

• Graham's Law of Effusion: $\frac{\text{rate}_1}{\text{rate}_2} = \sqrt{\frac{M_2}{M_1}}$

Example: Calculating the volume of CO2 produced from a reaction at a given temperature and pressure.

Thermochemistry

Enthalpy, Calorimetry, and Bond Energies

Thermochemistry studies energy changes in chemical reactions, focusing on enthalpy (ΔH). Calorimetry measures heat transfer, and bond energies estimate reaction enthalpy changes.

• Enthalpy Change (ΔH): $\Delta H = \sum \Delta H_f^\circ(\text{products}) - \sum \Delta H_f^\circ(\text{reactants})$

• Bond Energy Method: $\Delta H = \sum D(\text{bonds broken}) - \sum D(\text{bonds formed})$

• Calorimetry: $q = mc\Delta T$ (where m = mass, c = specific heat, ΔT = temperature change)

Example: Calculating ΔH for CH4 + 3O2 → 2CO2 + 2H2O using bond energies and standard enthalpies of formation.

Quantum Mechanics and Atomic Structure

Electromagnetic Radiation and Atomic Orbitals

Quantum mechanics explains the behavior of electrons in atoms. Electromagnetic radiation is characterized by wavelength, frequency, and energy.

• Energy of a Photon: $E = h\nu = \frac{hc}{\lambda}$

• Frequency and Wavelength: $c = \lambda \nu$

• Quantum Numbers: Describe the energy, shape, and orientation of atomic orbitals.

Example: Calculating the energy of radiation with a wavelength of 365.1 nm.

Molecular Orbital Theory

Bond Order and Magnetism

Molecular orbital (MO) theory describes the distribution of electrons in molecules. Bond order indicates the strength and stability of a bond, while paramagnetism or diamagnetism depends on unpaired electrons.

• Bond Order: $\text{Bond order} = \frac{\text{number of bonding electrons} - \text{number of antibonding electrons}}{2}$

• Paramagnetic: Molecules with unpaired electrons.

• Diamagnetic: Molecules with all electrons paired.

Example: Filling MO diagrams for NO, NO-, NO+, and NO2+ to determine bond order and magnetism.

Reference Tables and Constants

Physical Constants and Conversion Factors

Essential constants and conversion factors are used throughout general chemistry calculations.

Constant	Symbol	Value
Avogadro's Number	NA	6.022 × 1023 mol-1
Planck's Constant	h	6.626 × 10-34 J·s
Speed of Light	c	2.998 × 108 m/s
Ideal Gas Constant	R	0.08206 L·atm·mol-1·K-1

Bond Energies and Bond Lengths

Bond energies and lengths are used to estimate reaction enthalpies and molecular structures.

Bond	Bond Energy (kJ/mol)	Bond Length (pm)
H–H	436	74
C–H	414	109
O=O	498	121
N≡N	941	110
Cl–Cl	243	199
Additional info:	See full tables for more values.	—

Standard Enthalpies of Formation (ΔHf°)

Standard enthalpy values are used to calculate reaction enthalpies.

Compound	ΔHf° (kJ/mol)
CO2 (g)	-393.5
H2O (l)	-285.8
CH4 (g)	-74.8
Additional info:	See full tables for more values.

Solubility Rules for Ionic Compounds

Solubility rules help predict whether a precipitate will form in aqueous reactions.

• All salts containing Group 1 cations and NH4+ are soluble.

• All nitrates (NO3-), acetates (CH3COO-), and most perchlorates (ClO4-) are soluble.

• Most chlorides, bromides, and iodides are soluble except with Ag+, Pb2+, and Hg22+.

• Most sulfates are soluble except with Ba2+, Sr2+, Pb2+, and Ca2+.

• Most carbonates, phosphates, and sulfides are insoluble except with Group 1 and NH4+.

Selected Formulas

Formula	Description
$E = h\nu = \frac{hc}{\lambda}$	Energy of a photon
$PV = nRT$	Ideal gas law
$q = mc\Delta T$	Heat transfer (calorimetry)
$\Delta H = \sum D(\text{bonds broken}) - \sum D(\text{bonds formed})$	Bond energy method for enthalpy
$\Delta H = \sum \Delta H_f^\circ(\text{products}) - \sum \Delta H_f^\circ(\text{reactants})$	Standard enthalpy of reaction
$\text{Bond order} = \frac{\text{bonding electrons} - \text{antibonding electrons}}{2}$	Molecular orbital theory

Additional info:

• Reference tables include a periodic table, solubility table, and physical constants for quick access during exams.

• Practice problems cover Lewis structures, thermochemistry, gas laws, solution stoichiometry, molecular orbitals, and quantum mechanics, reflecting the breadth of a General Chemistry I curriculum.