General Chemistry: Core Concepts, Calculations, and Applications

General Chemistry Study Guide

Overview and Structure

This study guide covers essential topics in General Chemistry, including atomic structure, chemical bonding, thermodynamics, kinetics, acids and bases, redox reactions, and inorganic chemistry. It is structured to provide definitions, explanations, formulas, and examples relevant for college-level exam preparation.

Atomic Structure and Chemical Bonding

Electron Configuration and Ions

Understanding electron configurations is fundamental to predicting chemical behavior and bonding.

• Electron Configuration: The arrangement of electrons in an atom's orbitals, following the Aufbau principle, Pauli exclusion principle, and Hund's rule.

• Example: Selenium (Se) has atomic number 34. Its ground-state electron configuration is: $1s^2 2s^2 2p^6 3s^2 3p^6 4s^2 3d^{10} 4p^4$ The selenide ion (Se2−) gains two electrons: $1s^2 2s^2 2p^6 3s^2 3p^6 4s^2 3d^{10} 4p^6$

• Ion Formation: Atoms gain or lose electrons to achieve noble gas configurations, forming cations or anions.

Periodic Table and Trends

The periodic table organizes elements by increasing atomic number and recurring chemical properties.

• Groups and Periods: Vertical columns are groups (similar valence electron configurations); horizontal rows are periods.

• Trends:

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

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

  • Electronegativity: Increases across a period, decreases down a group.

Chemical Bonding and Molecular Structure

Bond Types and Geometry

Chemical bonds include ionic, covalent, and metallic bonds. Molecular geometry is determined by the VSEPR (Valence Shell Electron Pair Repulsion) theory.

• Ionic Bonds: Formed by transfer of electrons between metals and nonmetals.

• Covalent Bonds: Formed by sharing electrons between nonmetals.

• Lewis Structures: Diagrams showing bonding and lone pairs.

• Resonance Structures: Multiple valid Lewis structures for a molecule or ion, differing in electron arrangement.

• Example: Monoborate anion resonance structures should show correct geometry, bond angles, and lone pairs.

Thermodynamics

Key Concepts and Formulas

Thermodynamics studies energy changes in chemical processes.

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

• Entropy ($S$): Measure of disorder or randomness.

• Gibbs Free Energy ($G$): Determines spontaneity of a process. $G = H - TS$

• Standard Values: Standard enthalpy and entropy values are used for calculations at 298 K.

• Phase Changes: Melting, vaporization, and sublimation involve enthalpy and entropy changes.

• Clausius-Clapeyron Equation: Relates vapor pressure and temperature: $\ln P = -\frac{\Delta H_{vap}}{R} \cdot \frac{1}{T} + C$

Example Table: Thermodynamic Data for Water

Property	Phase
Melting Enthalpy	Ice
Vaporization Enthalpy	Liquid Water
Standard Gibbs Energy	Liquid Water
Standard Gibbs Energy	Gaseous Water
Standard Enthalpy	Liquid Water
Standard Enthalpy	Gaseous Water

Kinetics

Reaction Rates and Arrhenius Equation

Chemical kinetics studies the speed of reactions and factors affecting it.

• Rate Law: Expresses reaction rate as a function of reactant concentrations. $\text{Rate} = k[A]^m[B]^n$

• Order of Reaction: Sum of exponents in the rate law.

• Arrhenius Equation: Describes temperature dependence of rate constant: $k = A e^{-E_a/(RT)}$ Where $E_a$ is activation energy, $R$ is gas constant, $T$ is temperature.

• Graphical Analysis: Plotting $\ln k$ vs. $1/T$ yields a straight line with slope $-E_a/R$.

Example Table: Temperature Dependence of Rate Constant

Temperature (K)	Rate Constant (k)
387	7.2 × 10−4
407	2.2 × 10−3
447	1.7 × 10−2
457	0.11

Acids, Bases, and Precipitation

Acid-Base Equilibria and pH

Acids donate protons (H+), bases accept protons. The pH measures hydrogen ion concentration.

• pH Calculation: $\text{pH} = -\log[H^+]$

• Speciation Diagrams: Show predominant species at different pH values.

• Example: For oxalic acid and its conjugate bases, the diagram shows HOX, Ox2−, H2Ox, and H+ species as a function of pH.

Redox Reactions

Oxidation and Reduction

Redox reactions involve electron transfer between species.

• Oxidation: Loss of electrons.

• Reduction: Gain of electrons.

• Oxidation States: Assigned to atoms to track electron transfer.

• Example: Oxygen typically has oxidation state −2, except in peroxides and OF2.

Inorganic Chemistry and Complexes

Coordination Compounds

Complexes consist of a central metal ion bonded to ligands.

• Ligands: Molecules or ions that donate electron pairs to the metal.

• Coordination Number: Number of ligand attachments to the metal.

• Example: KBr forms a crystal lattice with K+ and Br− ions.

Constants and Reference Values

Fundamental Constants

• Elementary Charge: $e = 1.602 \times 10^{-19}$ C

• Gas Constant: $R = 8.314$ J mol−1 K−1

• Avogadro's Number: $N_A = 6.022 \times 10^{23}$ mol−1

• Molar Volume: $V_m = 24.79$ L mol−1 (at 298.15 K and 100 kPa)

• Faraday Constant: $F = 96485$ C mol−1

Sample Table: True/False Statements in Chemistry

Statement	Correctness
Selenium does not exceed its electron octet in compounds.	True
Oxygen only occurs in oxidation states −1 and −2.	False
Fe2+ forms Fe3+ by losing a proton.	False
Aluminum ion is smaller than silver ion.	True
Ionization energy of Au is greater than Ag.	True
Germanium has a higher atomic radius than arsenic.	True
Magnesium's third ionization energy is greater than its second.	True
Potassium and bromine are adjacent in KBr crystal lattice.	True
Threefold bonding is observed in triple bonds.	True

Additional info:

• Some context and explanations have been expanded for clarity and completeness.

• Tables and diagrams referenced are described in text for accessibility.