Stoichiometry and Aqueous Reactions: Study Notes for General Chemistry

Stoichiometry: Calculations with Chemical Formulas and Equations

Balancing Chemical Equations

Balancing chemical equations ensures that the number of atoms for each element is the same on both sides of the equation, reflecting the law of conservation of mass.

• Diatomic Elements: Elements that naturally exist as molecules of two atoms: H2, O2, F2, Cl2, Br2, I2.

• Writing Chemical Formulas: Use the charge criss-cross method to write formulas for ionic compounds (e.g., magnesium nitride).

• Balancing Steps: Adjust coefficients (not subscripts) to balance equations. Never change the chemical formulas of the reactants or products.

• Reference: See Section 3.1 in your textbook for more examples.

Simple Patterns of Chemical Reactivity

Chemical reactions can be classified into several types based on the changes that occur.

• Combination (Synthesis) Reactions: Two or more substances combine to form a single product. General form: $A + B \rightarrow C$

• Decomposition Reactions: A single compound breaks down into two or more simpler substances. General form: $AB \rightarrow A + B$

• Combustion Reactions: A substance reacts with oxygen, producing a flame, usually forming CO2 and H2O. Example: Hydrocarbon + O2 $\rightarrow$ CO2 + H2O

Formula Weights and Percent Composition

The formula weight of a compound is the sum of the atomic weights of all atoms in its formula.

• Formula Weight: For ionic compounds, also called formula units; for covalent compounds, called molecular weight.

• Calculation Example: The formula weight of Ca(NO2)3 is 164.1 amu.

• Percent Composition: The percentage by mass of each element in a compound. Formula: $\%\ \text{Element X} = \frac{\text{atomic weight of X} \times \text{number of atoms of X}}{\text{formula weight of compound}} \times 100\%$

Avogadro’s Number and the Mole

The mole is a fundamental unit for counting particles in chemistry, relating mass to number of particles.

• Avogadro’s Number: $6.02 \times 10^{23}$ particles per mole.

• Mole Definition: The amount of substance whose mass in grams equals its atomic or formula weight.

• Molar Mass: The mass of one mole of a substance (g/mol), numerically equal to the formula weight in amu.

• Conversions: Grams $\leftrightarrow$ Moles $\leftrightarrow$ Particles (must convert through moles).

Empirical Formula and Its Determination

The empirical formula shows the simplest whole-number ratio of elements in a compound.

• From Percent Composition:

  1. Assume 100 g of compound (percentages become grams).

  2. Calculate moles of each element: $\text{moles} = \frac{\text{grams}}{\text{atomic mass}}$

  3. Divide by the smallest number of moles to get subscripts.

• From Combustion Data:

  1. Find grams of C from CO2 produced.

  2. Find grams of H from H2O produced.

  3. Find grams of O by subtracting C and H from total mass.

Quantitative Information from Balanced Equations

Balanced equations provide the mole ratios needed for stoichiometric calculations.

• Coefficients: Indicate the number of moles (or molecules) of each substance involved.

• Mass Conservation: Total mass is conserved, but the number of moles may differ between reactants and products.

• Stoichiometry Steps:

  1. Convert mass of given substance to moles using molar mass.

  2. Use coefficients to find moles of unknown substance.

  3. Convert moles of unknown to mass using its molar mass.

Limiting Reactants and Percent Yield

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

• Identifying Limiting Reactant: Calculate how much of one reactant is needed to react with the other; the reactant that is used up first is limiting.

• Theoretical Yield: Maximum amount of product possible from the limiting reactant.

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

• Percent Yield Formula:

  • $\%\ \text{Yield} = \frac{\text{Actual Yield}}{\text{Theoretical Yield}} \times 100\%$

Aqueous Reactions and Solution Stoichiometry

General Properties of Aqueous Solutions

Aqueous solutions are homogeneous mixtures where water is the solvent.

• Solution: Homogeneous mixture of two or more substances.

• Solute: Substance dissolved (smaller amount).

• Solvent: Substance doing the dissolving (larger amount, usually water).

Electrolytic Properties

Electrolytes are substances that produce ions in solution, allowing the solution to conduct electricity.

• Strong Electrolytes: Completely ionize in water (strong acids, strong bases, soluble salts).

• Strong Acids to Know: HCl, HBr, HI, HNO3, H2SO4, HClO4, HClO3.

• Strong Bases: Hydroxides of Group IA and IIA metals (except Be).

• Weak Electrolytes: Partially ionize (weak acids and bases, e.g., NH3, CH3NH2).

• Nonelectrolytes: Do not produce ions (most molecular substances that are not acids or bases).

Precipitation Reactions

Precipitation reactions occur when mixing two solutions produces an insoluble solid (precipitate).

• Precipitate: The insoluble solid formed.

• Solubility Guidelines: Used to predict whether a precipitate will form (see Table 4.1 in your textbook).

• Exchange (Metathesis) Reactions: Ions in two compounds exchange partners: $AX + BY \rightarrow AY + BX$

• Types of Equations:

  • Molecular Equation: All species written as molecules.

  • Complete Ionic Equation: All strong electrolytes written as ions.

  • Net Ionic Equation: Only species that change during the reaction; spectator ions omitted.

Acid-Base Reactions

Acid-base reactions involve the transfer of protons (H+) between reactants.

• Neutralization Reaction: Acid reacts with base to form salt and water. General form: $\text{Acid} + \text{Base} \rightarrow \text{Salt} + \text{Water}$ $HA + MOH \rightarrow MA + H_2O$

• Example: $HClO_4 + KOH \rightarrow KClO_4 + H_2O$

• Gas Formation: Some acid-base reactions produce a gas as a product (e.g., CO2).

Oxidation-Reduction (Redox) Reactions

Redox reactions involve the transfer of electrons between substances.

• Oxidation: Loss of electrons.

• Reduction: Gain of electrons.

Concentrations of Solutions

Concentration expresses the amount of solute in a given amount of solution.

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

• Dissociation: When electrolytes dissolve, they produce multiple moles of ions per mole of compound (e.g., 1 mole of Ca(NO3)2 produces 3 moles of ions).

• Dilution Equation: $M_{conc} \times V_{conc} = M_{dil} \times V_{dil}$

Solution Stoichiometry and Chemical Analysis

Stoichiometric calculations in solution are based on the volume and concentration of solutions.

• Moles from Solution: $\text{Moles} = \text{Molarity} \times \text{Volume (L)}$

• Titrations: Analytical technique to determine the concentration of a solute using a reaction with a standard solution.

• Procedure: Use balanced equations and stoichiometric relationships to relate volumes and concentrations of reactants and products.

Additional info:

• For more practice, refer to the sample exercises and figures mentioned in your textbook (e.g., Figure 3.13, Figure 4.19, Sample Exercises 3.15, 4.7, 4.11).

• Understanding the difference between molecular, complete ionic, and net ionic equations is crucial for mastering solution chemistry.