Cumulative Skills

Metric Conversions

Understanding and converting between metric units is essential for accurate chemical calculations.

• Metric Prefixes: Common prefixes include kilo- (k, 103), centi- (c, 10-2), milli- (m, 10-3), micro- (μ, 10-6), and nano- (n, 10-9).

• Conversion Example: To convert 250 mg to grams:

Dimensional Analysis

Dimensional analysis is a systematic approach to problem-solving that uses conversion factors to move between units.

• Key Steps: Identify the starting unit, the desired unit, and set up conversion factors so units cancel appropriately.

• Example: Convert 5.0 moles of NaCl to grams (molar mass = 58.44 g/mol):

Chemical Nomenclature

Chemical nomenclature refers to the systematic naming of chemical compounds.

• Ionic Compounds: Name the cation first, then the anion (e.g., NaCl is sodium chloride).

• Molecular Compounds: Use prefixes to indicate the number of each atom (e.g., CO2 is carbon dioxide).

• Acids: Binary acids use the prefix "hydro-" and the suffix "-ic" (e.g., HCl is hydrochloric acid); oxyacids are named based on the polyatomic ion (e.g., H2SO4 is sulfuric acid).

Chapter 4: Chemical Reactions & Stoichiometry

Balancing Non-Redox Chemical Equations

Balancing chemical equations ensures the law of conservation of mass is obeyed.

• Include Phases: Indicate the physical state: (s), (l), (g), (aq).

• Example:

Stoichiometric Calculations Using Dimensional Analysis

Stoichiometry involves quantitative relationships between reactants and products in a chemical reaction.

• Key Steps: Convert given quantities to moles, use mole ratios from the balanced equation, and convert to desired units.

• Example: How many grams of CO2 are produced from 10.0 g of C6H12O6 in the reaction ?

Relating Grams, Moles, and Molecules

• Key Relationships:

• Avogadro's Number: particles/mol

• Example: contains molecules.

Limiting Reagent

The limiting reagent is the reactant that is completely consumed first, limiting the amount of product formed.

• Identification: Calculate the amount of product each reactant can produce; the smallest amount indicates the limiting reagent.

Theoretical, Actual, and Percent Yield

• Theoretical Yield: Maximum amount of product possible, calculated from stoichiometry.

• Actual Yield: Amount of product actually obtained from the reaction.

• Percent Yield Formula:

Molarity and Solution Calculations

• Molarity (M):

• Mass, Volume, and Dilution: for dilution calculations

• Example: To prepare 250 mL of 0.50 M NaCl from 2.0 M stock:

Chapter 5: Aqueous Reactions & Acid–Base Chemistry

Electrolytes and Solubility

• Strong Electrolytes: Completely dissociate in water (e.g., NaCl, HCl).

• Weak Electrolytes: Partially dissociate (e.g., CH3COOH).

• Non-Electrolytes: Do not dissociate (e.g., sugar, ethanol).

Solubility Rules and Precipitation

• Solubility Rules: Guidelines to predict if an ionic compound is soluble in water.

• Precipitate: An insoluble solid formed in a reaction.

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

Molecular, Total Ionic, and Net Ionic Equations

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

• Total Ionic Equation: Shows all strong electrolytes as ions.

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

• Spectator Ions: Ions that do not participate in the reaction.

• Example: For AgNO3 + NaCl → AgCl(s) + NaNO3:

  • Molecular:

  • Total Ionic:

  • Net Ionic:

Acids, Bases, and Conjugate Pairs

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

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

• Weak Acids/Bases: Partially ionize (e.g., CH3COOH, NH3)

• Conjugate Acid–Base Pairs: Differ by one H+ (e.g., NH4+/NH3)

Neutralization Reactions

• General Form: Acid + Base → Salt + Water

• Example:

pH and [H+]

• pH Formula:

• Classifying Solutions: pH < 7: acidic; pH = 7: neutral; pH > 7: basic

• Example: If , then

Balancing Redox Equations

• Redox (Oxidation-Reduction) Reactions: Involve transfer of electrons.

• Balancing: Assign oxidation numbers, balance atoms and charges, often using the half-reaction method.

Chapter 6: Gases

Empirical Gas Laws

Empirical gas laws describe the relationships between pressure, volume, temperature, and amount of gas.

• Boyle's Law: (at constant T and n)

• Charles's Law: (at constant P and n)

• Avogadro's Law: (at constant P and T)

Ideal Gas Law

• Equation:

• Variables: P = pressure (atm), V = volume (L), n = moles, R = 0.0821 L·atm/(mol·K), T = temperature (K)

• Example: Calculate the volume occupied by 2.0 mol of gas at 1.0 atm and 273 K:

Stoichiometry with Gas Laws

• Combine: Use the ideal gas law to relate moles of gas to volume, then apply stoichiometry for reactions involving gases.

• Example: What volume of O2 at STP is needed to react with 4.0 g H2 in ?

Dalton’s Law of Partial Pressures

• Law Statement: The total pressure of a mixture of gases equals the sum of the partial pressures of each gas.

• Equation:

• Example: If a container has O2 at 0.50 atm and N2 at 0.75 atm, atm.

Summary Table: Key Concepts and Formulas

Additional info: This guide covers foundational skills and core concepts from Chapters 4–6, including chemical equations, stoichiometry, solution chemistry, acid–base reactions, and gas laws. Students should also review lab techniques, mathematical operations, and practice integrated problem-solving as indicated by the exam format.