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; the rest are core electrons.

Electron Configuration and the Aufbau Principle

Electron configurations describe the arrangement of electrons in an atom. The Aufbau principle states that electrons fill orbitals starting with the lowest energy first. The Pauli exclusion principle and Hund's rule further govern electron placement.

• Aufbau principle: Electrons occupy the lowest energy orbitals available.

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

• Hund's rule: Electrons fill 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 the electrons in them:

• 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

Atomic Radius, Ionization Energy, and Electron Affinity

Periodic trends help predict element behavior:

• 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.

• Example: Fluorine has a smaller atomic radius and higher ionization energy than sodium.

Transition Metals and f-block Elements

Transition metals (d-block) and f-block elements (lanthanides and actinides) have unique electron configurations and properties.

• Transition metals: Elements in groups 3-12, characterized by partially filled d orbitals.

• f-block elements: Lanthanides and actinides, characterized by filling of 4f and 5f orbitals.

Magnetism and Electron Pairing

Paramagnetism and Diamagnetism

Magnetic properties depend on electron pairing:

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

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

• Example: O2 is paramagnetic due to two unpaired electrons.

Electronic Structure and Spectroscopy

Photoelectric Effect and Electromagnetic Spectrum

The photoelectric effect demonstrates the particle nature of light. The electromagnetic spectrum includes all types of electromagnetic radiation, characterized by wavelength and frequency.

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

• Electromagnetic spectrum: Range from radio waves to gamma rays.

• Equation: $E = h\nu$ (energy of a photon)

• Equation: $c = \lambda \nu$ (speed of light relation)

Wave-Particle Duality and Interference

Light and electrons exhibit both wave-like and particle-like properties. Constructive and destructive interference occur when waves overlap.

• Constructive interference: Waves add together, increasing amplitude.

• Destructive interference: Waves cancel each other, reducing amplitude.

Thermochemistry

Heat, Work, and Energy

Thermochemistry studies energy changes in chemical reactions.

• Heat (q): Energy transferred due to temperature difference.

• Work (w): Energy transferred when an object is moved by a force.

• Internal energy (U): Total energy of a system.

• Equation: $\Delta U = q + w$

Enthalpy and Calorimetry

Enthalpy (H) is the heat content of a system at constant pressure. Calorimetry measures heat changes.

• Equation: $q = m c \Delta T$ (heat absorbed or released)

• Specific heat capacity: Amount of heat required to raise the temperature of 1 g of substance by 1°C.

• Example: Calculating heat absorbed by water when heated.

Heats of Formation and Reaction

Heats of formation and heats of reaction are used to calculate enthalpy changes.

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

• Equation: $\Delta H_{rxn} = \sum \Delta H_f^\circ (\text{products}) - \sum \Delta H_f^\circ (\text{reactants})$

Chemical Reactions and Stoichiometry

Types of Chemical Reactions

Chemical reactions can be classified as acid-base, gas evolution, oxidation-reduction, and more.

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

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

• Oxidation-reduction (redox) reactions: Involve transfer of electrons.

Writing and Balancing Equations

Equations must be balanced to obey the law of conservation of mass. Ionic equations show the species involved in reactions in aqueous solution.

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

• Complete ionic equation: Shows all ions present.

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

Stoichiometry and Molar Calculations

Stoichiometry involves calculations based on balanced chemical equations.

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

• Equation: $\text{Moles} = \frac{\text{Mass (g)}}{\text{Molar mass (g/mol)}}$

• Example: Calculating moles of NaCl from 58.44 g sample.

Concentration and Molarity

Molarity (M) is a measure of concentration in moles per liter.

• Equation: $M = \frac{n}{V}$, where n = moles of solute, V = volume of solution in liters.

• Example: Preparing 1.0 M NaCl solution by dissolving 58.44 g NaCl in 1.0 L water.

Identifying Spectator Ions

Spectator ions do not participate in the chemical reaction and remain unchanged in solution.

• Example: In the reaction between NaCl and AgNO3, Na+ and NO3- are spectator ions.

Acids and Bases

Strong vs. Weak Acids

Acids are classified based on their degree of ionization in water.

• Strong acids: Completely ionize in water (e.g., HCl, HNO3).

• Weak acids: Partially ionize in water (e.g., CH3COOH).

Oxidation-Reduction (Redox) Reactions

Assigning Oxidation States

Oxidation states help track electron transfer in redox reactions.

• Oxidizing agent: Causes oxidation by accepting electrons.

• Reducing agent: Causes reduction by donating electrons.

• Example: In Zn + Cu2+ → Zn2+ + Cu, Zn is the reducing agent, Cu2+ is the oxidizing agent.

Summary Table: Key Concepts and Equations

Concept	Key Equation	Example/Application
Energy of photon	$E = h\nu$	Calculate energy for light of frequency 5.0 x 1014 Hz
Speed of light	$c = \lambda \nu$	Find wavelength for frequency 6.0 x 1014 Hz
Heat change	$q = m c \Delta T$	Calculate heat absorbed by 100 g water heated by 10°C
Molarity	$M = \frac{n}{V}$	Prepare 0.5 M solution from 0.25 mol solute in 0.5 L
Enthalpy of reaction	$\Delta H_{rxn} = \sum \Delta H_f^\circ (\text{products}) - \sum \Delta H_f^\circ (\text{reactants})$	Calculate $\Delta H_{rxn}$ for combustion of methane

Additional info: Some topics (e.g., double-slit experiment, electromagnetic spectrum, quantum numbers) are expanded for completeness and context. This guide covers atomic structure, periodicity, chemical reactions, thermochemistry, and related calculations as outlined in the study guide.