General Chemistry Study Notes: Solutions, Electrolytes, Acids/Bases, and Redox Reactions

Chapter 4: Solutions and Their Properties

Molarity and Dissolved Ions

Understanding the concentration of ions in solution is essential in chemistry, especially for health-related issues and laboratory analysis. Molarity is the most common unit used to express solution concentration.

• Molarity (M): Defined as the number of moles of solute per liter of solution.

• Formula: $M = \frac{\text{moles of solute}}{\text{liters of solution}}$

• Application: Used to calculate the concentration of specific ions in a solution, such as sodium ions in sodium sulfate or nitrate ions in mixed solutions.

• Example: What is the molarity of sodium ion in 25.0 mL of 0.100 M sodium sulfate solution?

Dilutions

Preparing solutions of desired concentration from a concentrated stock is a common laboratory practice. The dilution equation allows calculation of the final concentration after dilution.

• Dilution Equation: $M_1 V_1 = M_2 V_2$

• Key Principle: The number of moles of solute remains constant during dilution; only the volume and concentration change.

• Example: A 5.00 mL aliquot of 12.1 M HCl is diluted to 250.00 mL. What is the final molarity?

Classification of Solutions: Electrolytes and Non-Electrolytes

Electrolytes

Electrolytes are substances that dissolve in water to produce ions, allowing the solution to conduct electricity. They are essential for physiological processes and laboratory analysis.

• Types:

  • Strong Electrolytes: Ionize completely in water (e.g., soluble ionic compounds, strong acids, strong bases).

  • Weak Electrolytes: Ionize partially in water (e.g., weak acids, weak bases, slightly soluble salts).

• Examples: NaCl, KNO3 (strong); CH3COOH, NH3 (weak)

• Conductivity: Strong electrolytes conduct electricity well; weak electrolytes conduct poorly.

Non-Electrolytes

Non-electrolytes dissolve in water but do not produce ions, so their solutions do not conduct electricity.

• Examples: Alcohols (e.g., CH3OH), sugars (e.g., C12H22O11), pure water

• Properties: Molecular compounds that do not ionize in solution.

Acids and Bases

Definitions

Acids and bases are defined by their behavior in water and their ability to donate or accept protons.

• Arrhenius Definition:

  • Acid: Increases H+ in water

  • Base: Increases OH- in water

• Brønsted-Lowry Definition:

  • Acid: Proton (H+) donor

  • Base: Proton (H+) acceptor

Strong Acids and Bases

Strong acids and bases ionize completely in water and are strong electrolytes.

• 6 Strong Acids (memorize): HCl, HBr, HI, HNO3, HClO4, H2SO4

• Strong Bases: Alkali metal hydroxides (LiOH, NaOH, KOH, RbOH, CsOH) and Ba(OH)2

Acid/Base Properties

• Acids: Taste sour, turn litmus red, react with metals, conduct electricity

• Bases: Taste bitter, feel slippery, turn litmus blue, conduct electricity

Acid/Base Neutralization Reactions

When acids and bases react, they form a salt and water in a neutralization reaction.

• General Equation: $\text{Acid} + \text{Base} \rightarrow \text{Salt} + \text{Water}$

• Example: $\text{HCl}(aq) + \text{NaOH}(aq) \rightarrow \text{NaCl}(aq) + \text{H}_2\text{O}(l)$

Titrations

Titration is a technique used to determine the concentration of an unknown solution by reacting it with a solution of known concentration.

• Key Terms:

  • Equivalence Point: The point at which stoichiometrically equivalent amounts of acid and base have reacted.

  • Endpoint: The observable change (often color change) indicating completion.

• Indicator: A dye (e.g., phenolphthalein) used to signal the endpoint.

• Stoichiometry: Use balanced chemical equations to relate moles of reactants and products.

• Example: 30.00 mL of 0.1000 M NaOH neutralizes 50.00 mL of H2SO4. Find the concentration of H2SO4.

Solubility and Precipitation Reactions

Solubility Rules

Solubility rules help predict whether an ionic compound will dissolve in water or form a precipitate.

Soluble Compounds	Exceptions
Group 1A ions (Li+, Na+, K+, etc.)	None
Ammonium ion (NH4+)	None
Nitrates (NO3-), Acetates (C2H3O2-), Bicarbonates (HCO3-)	None
Chlorides, Bromides, Iodides	Ag+, Pb2+, Hg22+
Sulfates (SO42-)	Ag+, Ca2+, Sr2+, Ba2+, Pb2+

Insoluble Compounds	Exceptions
Carbonates (CO32-), Phosphates (PO43-), Sulfides (S2-), Hydroxides (OH-)	Group 1A ions, NH4+, Ba(OH)2 (slightly soluble)

Additional info: Table entries inferred and summarized from standard solubility rules.

Precipitation Reactions

When two solutions of soluble ionic compounds are mixed, an insoluble compound may form as a precipitate.

• General Pattern: Double displacement (partner exchange) reactions: $AB + CD \rightarrow AD + CB$

• Example: $\text{NaCl}(aq) + \text{AgNO}_3(aq) \rightarrow \text{AgCl}(s) + \text{NaNO}_3(aq)$

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

Oxidation-Reduction (Redox) Reactions

Redox Concepts

Redox reactions involve the transfer of electrons between substances. They are fundamental in processes such as battery operation and corrosion.

• Oxidation: Loss of electrons

• Reduction: Gain of electrons

• Mnemonic: OIL RIG (Oxidation Is Loss, Reduction Is Gain)

• Example: $2\text{Ca}(s) + \text{O}_2(g) \rightarrow 2\text{CaO}(s)$

Assigning Oxidation Numbers

Oxidation numbers are used to keep track of electron transfer in compounds and reactions.

• Rules:

  1. Elements in their elemental form have oxidation number 0.

  2. Group 1A metals: +1; Group 2A metals: +2

  3. Fluorine: always -1

  4. Oxygen: usually -2 (except in peroxides: -1)

  5. Hydrogen: +1 (except in metal hydrides: -1)

  6. Sum of oxidation numbers in a neutral compound is 0; in a polyatomic ion, equals the ion's charge.

• Example: Assign oxidation numbers to all elements in KMnO4.

Redox Agents

• Oxidizing Agent: Causes another substance to be oxidized; is itself reduced.

• Reducing Agent: Causes another substance to be reduced; is itself oxidized.

Single Displacement and Activity Series

Single displacement reactions involve a metal displacing another metal or hydrogen from a compound, guided by the activity series.

• General Reaction: $M(s) + 2HCl(aq) \rightarrow MCl_2(aq) + H_2(g)$

• Activity Series: Metals higher in the series are more easily oxidized and can displace metals lower in the series or hydrogen from acids.

• Application: Predicting whether a reaction will occur based on the relative activity of metals.

Additional info: Some table and example content inferred from standard chemistry curriculum.