General Chemistry: Practice Checklist and Study Guide for Chemical Reactions, Stoichiometry, and Acid-Base Chemistry

Types of Chemical Reactions

Acid-Base Reactions

Acid-base reactions involve the transfer of protons (H+) between reactants. These reactions are fundamental in chemistry and are classified based on the strength of acids and bases involved.

• Strong Acids and Bases: Completely dissociate in water. Examples: HCl, NaOH.

• Weak Acids and Bases: Partially dissociate in water. Examples: CH3COOH, NH3.

• Net Ionic Equations: Show only the species that participate directly in the reaction.

• Salt Formation: Acids and bases react to form salts and water.

Example: $\mathrm{HCl} + \mathrm{NaOH} \rightarrow \mathrm{NaCl} + \mathrm{H_2O}$

Precipitation Reactions

Precipitation reactions occur when two aqueous solutions combine to form an insoluble product, called a precipitate.

• Solubility Rules: Used to predict whether a precipitate will form.

• Net Ionic Equations: Only include ions and molecules directly involved in the formation of the precipitate.

Example: $\mathrm{AgNO_3} (aq) + \mathrm{NaCl} (aq) \rightarrow \mathrm{AgCl} (s) + \mathrm{NaNO_3} (aq)$

Redox (Oxidation-Reduction) Reactions

Redox reactions involve the transfer of electrons between species. Oxidation is the loss of electrons, while reduction is the gain of electrons.

• Half-Reactions: Show oxidation and reduction separately.

• Activity Series: Predicts whether a single replacement reaction will occur.

• Balancing Redox Equations: Use the half-reaction method to ensure mass and charge are balanced.

Example: $\mathrm{Zn} (s) + \mathrm{Cu^{2+}} (aq) \rightarrow \mathrm{Zn^{2+}} (aq) + \mathrm{Cu} (s)$

Chemical Quantities and Stoichiometry

Mole Concept and Molar Mass

The mole is a fundamental unit in chemistry used to count particles. Molar mass is the mass of one mole of a substance.

• Avogadro's Number: $6.022 \times 10^{23}$ particles per mole.

• Molar Mass: Sum of atomic masses in grams per mole.

Example: $\text{Molar mass of H}_2\text{O} = 2 \times 1.008 + 15.999 = 18.015\ \text{g/mol}$

Stoichiometric Calculations

Stoichiometry involves using balanced chemical equations to calculate the amounts of reactants and products.

• Steps:

  1. Write and balance the chemical equation.

  2. Convert quantities to moles.

  3. Use mole ratios from the equation.

  4. Convert moles to desired units (grams, liters, etc.).

Example: $\mathrm{2H_2} + \mathrm{O_2} \rightarrow \mathrm{2H_2O}$

Solution Chemistry

Concentration Units

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

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

• Percent Composition: $\% = \frac{\text{mass of solute}}{\text{mass of solution}} \times 100$

Preparation of Solutions

Solutions are prepared by dissolving a known amount of solute in solvent to achieve a desired concentration.

• Serial Dilutions: Used to create solutions of lower concentration from a stock solution.

• Mixing Acids and Bases: Requires careful calculation to avoid hazardous reactions.

Thermochemical Aspects of Chemical Reactions

Enthalpy and Calorimetry

Thermochemistry studies the energy changes during chemical reactions, primarily focusing on enthalpy (ΔH).

• Enthalpy Change (ΔH): The heat absorbed or released at constant pressure.

• Calorimetry: Experimental technique to measure heat changes.

Example Equation: $q = m c \Delta T$

Acid-Base Equilibria

pH and pOH Calculations

pH measures the acidity of a solution, while pOH measures its basicity.

• pH Formula: $\mathrm{pH} = -\log[\mathrm{H^+}]$

• pOH Formula: $\mathrm{pOH} = -\log[\mathrm{OH^-}]$

• Relationship: $\mathrm{pH} + \mathrm{pOH} = 14$ (at 25°C)

Titrations and Buffers

Titrations are used to determine the concentration of an unknown acid or base. Buffers resist changes in pH upon addition of small amounts of acid or base.

• Equivalence Point: The point at which the amount of titrant added is stoichiometrically equal to the amount of analyte.

• Buffer Solution: Contains a weak acid and its conjugate base (or vice versa).

Example: $\mathrm{CH_3COOH} + \mathrm{NaOH} \rightarrow \mathrm{CH_3COONa} + \mathrm{H_2O}$

Tables

Solubility Rules Table

This table summarizes common solubility rules for ionic compounds in water.

Ion/Compound	Solubility	Exceptions
Na+, K+, NH4+	Soluble	None
NO3-, C2H3O2-	Soluble	None
Cl-, Br-, I-	Soluble	Ag+, Pb2+, Hg22+
SO42-	Soluble	Ba2+, Pb2+, Ca2+
CO32-, PO43-	Insoluble	Na+, K+, NH4+

Activity Series Table

The activity series ranks metals by their ability to displace other metals in single replacement reactions.

Metal	Reactivity
Li	Most Reactive
K	Highly Reactive
Ca	Reactive
Na	Reactive
Mg	Moderately Reactive
Zn	Less Reactive
Cu	Least Reactive

Additional info:

• Some content inferred and expanded for completeness, including definitions, formulas, and examples.

• Checklist items were grouped and elaborated into textbook-style sections for clarity.