General Chemistry I: Study Guide for Exam 2 – Atomic Structure, Periodicity, Chemical Reactions, and Thermochemistry

Atomic Structure and Electron Configuration

Valence Electrons and Core Electrons

Understanding the distinction between valence electrons and core electrons is essential for predicting chemical behavior. Valence electrons are those found in the outermost shell of an atom and are involved in chemical bonding, while core electrons are located in inner shells and do not participate directly in bonding.

• Valence electrons: Electrons in the highest principal energy level (n).

• Core electrons: Electrons in all lower energy levels.

• Example: For sodium (Na, atomic number 11), the electron configuration is 1s2 2s2 2p6 3s1. The single 3s electron is the valence electron.

Electron Configuration Principles

Electron configurations are determined by several key principles:

• Aufbau Principle: Electrons fill orbitals starting with the lowest energy first.

• Pauli Exclusion Principle: No two electrons in an atom can have the same set of four quantum numbers.

• Hund's Rule: Electrons occupy degenerate orbitals singly before pairing.

• Example: Carbon (C, atomic number 6): 1s2 2s2 2p2. The two 2p electrons occupy separate orbitals.

Quantum Numbers

Quantum numbers describe the properties of atomic orbitals and electrons:

• Principal quantum number (n): Indicates the energy level.

• Angular momentum quantum number (l): Indicates the shape of the orbital (s, p, d, f).

• Magnetic quantum number (ml): Indicates the orientation of the orbital.

• Spin quantum number (ms): Indicates the spin direction (+1/2 or -1/2).

• Example: For a 2p electron: n = 2, l = 1, ml = -1, 0, or +1, ms = +1/2 or -1/2.

Periodic Properties of the Elements

Trends in the Periodic Table

Periodic trends help predict element properties:

• 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.

• Metallic character: Increases down a group, decreases across a period.

Classification of Elements

• Transition elements: d-block elements, typically metals with variable oxidation states.

• f-block elements: Lanthanides and actinides.

Electronic Structure and Magnetism

Paramagnetism and Diamagnetism

Magnetic properties depend on electron configuration:

• Paramagnetic: Atoms or ions with unpaired electrons; attracted to magnetic fields.

• Diamagnetic: Atoms or ions with all electrons paired; weakly repelled by magnetic fields.

Wave-Particle Duality and Quantum Mechanics

Electromagnetic Radiation

Light exhibits both wave-like and particle-like properties.

• Wavelength ($\lambda$): Distance between successive peaks.

• Frequency ($\nu$): Number of cycles per second.

• Energy of a photon: $E = h\nu$

• Speed of light: $c = \lambda \nu$

Photoelectric Effect and Double-Slit Experiment

• Photoelectric effect: Emission of electrons from a metal when light shines on it.

• Double-slit experiment: Demonstrates wave-particle duality of light.

Chemical Reactions and Stoichiometry

Types of Chemical Equations

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

• Complete ionic equation: Shows all strong electrolytes as ions.

• Net ionic equation: Shows only species that change during the reaction.

Acid-Base and Gas Evolution Reactions

• Acid-base reactions: Involve transfer of protons (H+).

• Gas evolution reactions: Produce a gas as a product.

• Example: $\text{Na}_2\text{CO}_3 + 2\text{HCl} \rightarrow 2\text{NaCl} + \text{H}_2\text{O} + \text{CO}_2$

Oxidation-Reduction (Redox) Reactions

• Oxidation: Loss of electrons.

• Reduction: Gain of electrons.

• Oxidizing agent: Causes oxidation, is itself reduced.

• Reducing agent: Causes reduction, is itself oxidized.

Thermochemistry

Energy and Heat in Chemical Reactions

• Kinetic energy: Energy due to motion.

• Potential energy: Energy due to position or composition.

• Chemical energy: Energy stored in chemical bonds.

• Heat capacity ($C$): Amount of heat required to raise temperature by 1°C.

• Specific heat ($c$): Heat required to raise 1 g of substance by 1°C.

• Enthalpy ($\Delta H$): Heat change at constant pressure.

• Heats of reaction ($\Delta H_{rxn}$): Energy change during a chemical reaction.

• Heats of formation ($\Delta H_f$): Energy change when 1 mole of compound forms from its elements.

Calculations in Chemistry

Mole and Molar Mass Calculations

• Mole: Amount of substance containing Avogadro's number ($6.022 \times 10^{23}$) of particles.

• Molar mass: Mass of 1 mole of a substance (g/mol).

• Example: Calculate moles: $\text{moles} = \frac{\text{mass (g)}}{\text{molar mass (g/mol)}}$

Concentration and Molarity

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

• Example: 0.5 moles NaCl in 1 L solution: $M = 0.5 \text{ mol/L}$

Identifying Spectator Ions

• Spectator ions: Ions that do not participate in the actual chemical change.

Summary Table: Key Concepts and Calculations

Concept	Definition	Key Formula
Valence Electrons	Electrons in the outermost shell	N/A
Electron Configuration	Arrangement of electrons in orbitals	N/A
Energy of Photon	Energy carried by a photon	$E = h\nu$
Mole Calculation	Converting mass to moles	$\text{moles} = \frac{\text{mass}}{\text{molar mass}}$
Molarity	Concentration of solution	$M = \frac{\text{moles}}{\text{liters}}$
Enthalpy Change	Heat change at constant pressure	$\Delta H = \sum \Delta H_f(\text{products}) - \sum \Delta H_f(\text{reactants})$

Additional info: Some topics (e.g., double-slit experiment, electromagnetic spectrum) are expanded for completeness and context. The study guide covers material from chapters on atomic structure, periodicity, chemical reactions, thermochemistry, and solution chemistry, all relevant to General Chemistry I.