General Chemistry Study Guide: Chapters 4, 5, and 6 (Aqueous Reactions, Thermochemistry, Electronic Structure)

Chapter 4: Reactions in Aqueous Solution

4.1 General Properties of Aqueous Solutions

Aqueous solutions are mixtures where water acts as the solvent. Understanding their properties is essential for predicting chemical behavior in water.

• Aqueous Solution: A solution in which water is the solvent.

• Electrolyte: A substance that dissolves in water to produce ions, allowing the solution to conduct electricity.

• Nonelectrolyte: A substance that dissolves in water but does not produce ions; the solution does not conduct electricity.

• Strong vs. Weak Electrolyte: Strong electrolytes dissociate completely into ions; weak electrolytes only partially dissociate.

• Examples: NaCl (strong electrolyte), CH3COOH (weak electrolyte), sugar (nonelectrolyte).

4.2 Precipitation Reactions

Precipitation reactions occur when two solutions are mixed and an insoluble product (precipitate) forms.

• Solubility Chart: Used to determine if a compound is soluble (aq) or insoluble (s) in water.

• Precipitation Reaction: A reaction in which an insoluble solid forms and separates from the solution.

• Precipitate: The solid product formed; its phase is (s).

• Ionic Equation vs. Net Ionic Equation: Ionic equations show all ions; net ionic equations show only those involved in the reaction.

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

• Example: Mixing AgNO3 (aq) and NaCl (aq) forms AgCl (s) as a precipitate.

Solubility Guidelines Table

Ion	Soluble Compounds	Exceptions
NO3-	All	None
Cl-, Br-, I-	Most	Ag+, Hg22+, Pb2+
SO42-	Most	Ba2+, Pb2+, Ca2+, Sr2+
CO32-, PO43-	NH4+, alkali metals	Others insoluble
OH-	NH4+, alkali metals	Others insoluble

4.3 Acids, Bases, and Neutralization Reactions

Acids and bases are defined by their ability to donate or accept protons. Their reactions in water are fundamental to chemistry.

• Acid: Substance that donates H+ ions in solution.

• Base: Substance that accepts H+ ions or donates OH- ions.

• Strong vs. Weak Acids/Bases: Strong acids/bases dissociate completely; weak ones only partially.

• 7 Strong Acids: HCl, HBr, HI, HNO3, HClO4, HClO3, H2SO4

• 8 Strong Bases: Group 1 and 2 hydroxides (e.g., NaOH, KOH, Ca(OH)2)

• Neutralization Reaction: Acid + Base → Salt + Water

• Net Ionic Equation: Shows only the ions involved in forming water.

• Spectator Ions: Remain unchanged in solution.

• Example: HCl (aq) + NaOH (aq) → NaCl (aq) + H2O (l)

4.4 Oxidation-Reduction (Redox) Reactions

Redox reactions involve the transfer of electrons between species, changing their oxidation states.

• Oxidation: Loss of electrons (increase in oxidation number).

• Reduction: Gain of electrons (decrease in oxidation number).

• OIL RIG: "Oxidation Is Loss, Reduction Is Gain" (of electrons).

• Oxidation Number: Indicates the degree of oxidation of an atom in a compound.

• Example: Zn (s) + Cu2+ (aq) → Zn2+ (aq) + Cu (s)

4.5 Concentrations of Solutions

Concentration describes the amount of solute dissolved in a given quantity of solvent.

• Molarity (M): $M = \frac{\text{mol}}{\text{L}}$

• Dilution Equation: $M_1 V_1 = M_2 V_2$

• Unit Conversion: 1 L = 1000 mL

• Ion Concentration: For ionic compounds, the concentration of each ion depends on the formula.

• Example: In Na2SO4, [Na+] is twice [SO42-].

4.6 Solution Stoichiometry and Chemical Analysis

Stoichiometry in solutions involves using concentrations and volumes to calculate amounts of reactants and products.

• Solution Stoichiometry: Use $M$, $V$, and balanced equations to solve for unknowns.

• Titration: Analytical technique to determine concentration by reacting with a standard solution.

• Neutralization, Redox, Precipitation: All can be analyzed using stoichiometry.

• Example: Calculate volume of acid needed to neutralize a base.

Chapter 5: Thermochemistry

5.1 First Law of Thermodynamics

The first law states that energy cannot be created or destroyed, only transferred.

• First Law: $\Delta E = q + w$ (change in internal energy equals heat plus work)

• System: The part of the universe being studied.

• Surroundings: Everything else outside the system.

• State Function: Property dependent only on the current state, not the path (e.g., energy, enthalpy).

• Signs: $q$ and $w$ are positive if energy flows into the system, negative if out.

5.3 Enthalpy

Enthalpy is a measure of heat content at constant pressure.

• Enthalpy (H): $\Delta H$ is the heat change at constant pressure.

• Endothermic Process: Absorbs heat ($\Delta H > 0$).

• Exothermic Process: Releases heat ($\Delta H < 0$).

• Direction of Heat Flow: Endothermic: into system; Exothermic: out of system.

5.4 Enthalpies of Reaction

Calculating the enthalpy change for a reaction is essential for understanding energy changes.

• Enthalpy of Reaction: $\Delta H_{rxn} = \Sigma \Delta H_f^{\circ} (\text{products}) - \Sigma \Delta H_f^{\circ} (\text{reactants})$

• Application: Use molar amounts and chemical equations to calculate $\Delta H_{rxn}$.

5.5 Calorimetry

Calorimetry measures heat flow in chemical reactions.

• Basic Formula: $q = m \times C \times \Delta T$

• Units: Mass (g), Specific heat (J/g·K), Temperature change (K or °C).

• Application: Used to determine heat absorbed or released.

5.7 Enthalpies of Formation

Standard enthalpy of formation is the enthalpy change when one mole of a compound forms from its elements in their standard states.

• Standard Enthalpy of Formation ($\Delta H_f^{\circ}$): Used to calculate reaction enthalpy.

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

• Practice: Pay attention to coefficients and signs.

Chapter 6: Electronic Structure of Atoms

6.1 Electromagnetic Radiation

Light and other forms of electromagnetic radiation are characterized by wavelength, frequency, and energy.

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

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

• Speed of Light ($c$): $c = \lambda \nu$

• Photon Energy: $E = h \nu$

• Relationship: Higher frequency = higher energy; shorter wavelength = higher energy.

• Electromagnetic Spectrum: Includes radio, microwave, infrared, visible, ultraviolet, X-ray, gamma ray.

Electromagnetic Spectrum Table

Type	Wavelength (m)	Frequency (Hz)
Radio	10-1 to 103	106 to 109
Microwave	10-3 to 10-1	109 to 1011
Infrared	10-6 to 10-3	1011 to 1014
Visible	400–700 nm	4.3 × 1014 to 7.5 × 1014
Ultraviolet	10-8 to 10-7	1015 to 1016
X-ray	10-10 to 10-8	1016 to 1018
Gamma ray	10-12 to 10-10	1018 to 1020

6.2 Atomic Spectra and the Bohr Model

The Bohr model explains atomic emission spectra using quantized energy levels.

• Energy of Electron Transition: $\frac{1}{\lambda} = R_H \left( \frac{1}{n_f^2} - \frac{1}{n_i^2} \right)$

• Quantization: Electrons occupy discrete energy levels.

• Photon Emission: Occurs when electrons drop to lower energy levels.

6.5/6.6 Quantum Mechanics and Atomic Orbitals

Quantum mechanics describes electrons in atoms using quantum numbers and orbitals.

• Quantum Numbers:

  • Principal (n): Energy level (n = 1, 2, 3...)

  • Angular Momentum (l): Subshell (l = 0 to n-1)

  • Magnetic (ml): Orientation (-l to +l)

  • Spin (ms): Electron spin (+1/2 or -1/2)

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

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

• Electron Configuration: Distribution of electrons among orbitals (e.g., 1s2 2s2 2p6).

• Condensed Electron Configuration: Uses noble gas core to simplify notation (e.g., [Ne] 3s2).

• Periodic Table: Used to determine electron configurations and element properties.

Periodic Table Reference

The periodic table organizes elements by atomic number and electron configuration. It is essential for predicting chemical behavior and writing electron configurations.

Additional info: Some details and examples have been expanded for clarity and completeness. Practice problems and homework assignments referenced in the notes are recommended for mastery of these topics.