General Chemistry I: Core Topics and Study Guide

Stoichiometry

Reaction Stoichiometry

Stoichiometry involves the quantitative relationships between reactants and products in a chemical reaction, based on the balanced chemical equation.

• Mole Ratios: Use coefficients from the balanced equation to relate moles of different substances.

• Conversions: Convert between grams, moles, and molecules using molar mass and Avogadro's number.

• Example: For the reaction $2H_2 + O_2 \rightarrow 2H_2O$, 2 moles of hydrogen react with 1 mole of oxygen to produce 2 moles of water.

Limiting Reactants

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

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

• Example: If 5.0 g of A reacts with 10.0 g of B, determine which produces less product to find the limiting reactant.

Percent Yield

Percent yield measures the efficiency of a reaction by comparing the actual yield to the theoretical yield.

• Formula: $\text{Percent Yield} = \left( \frac{\text{Actual Yield}}{\text{Theoretical Yield}} \right) \times 100\%$

• Example: If the theoretical yield is 10.0 g and the actual yield is 8.5 g, percent yield is $\left( \frac{8.5}{10.0} \right) \times 100\% = 85\%$.

Empirical and Molecular Formulas

The empirical formula gives the simplest whole-number ratio of atoms in a compound, while the molecular formula shows the actual number of atoms.

• Empirical Formula: Convert percent composition to moles, divide by the smallest value, and round to nearest whole number.

• Molecular Formula: $\text{Molecular Formula} = n \times \text{Empirical Formula}$, where $n = \frac{\text{Molar Mass}}{\text{Empirical Formula Mass}}$

• Example: A compound with 40% C, 6.7% H, and 53.3% O has an empirical formula of CH2O. If its molar mass is 180 g/mol, the molecular formula is C6H12O6.

Reactions in Solution (Aqueous Reactions)

Molarity and Solution Calculations

Molarity (M) is the concentration of a solution, defined as moles of solute per liter of solution.

• Formula: $M = \frac{\text{moles of solute}}{\text{volume of solution (L)}}$

• Calculations: Use $M \times V = \text{moles}$ to find unknowns in dilution or reaction problems.

• Example: To prepare 0.5 L of 1.0 M NaCl, dissolve 0.5 mol NaCl in enough water to make 0.5 L.

Titration and Neutralization

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

• Neutralization: Acid and base react to form water and a salt.

• Formula: $M_1V_1 = M_2V_2$ (for monoprotic acid-base reactions)

• Example: 25.0 mL of 0.100 M HCl neutralizes 50.0 mL of NaOH. Find NaOH concentration: $0.100 \times 0.025 = M_{NaOH} \times 0.050$; $M_{NaOH} = 0.050$ M.

Precipitation Reactions

Precipitation occurs when two aqueous solutions react to form an insoluble solid (precipitate).

• Predicting Precipitates: Use solubility rules to determine if a solid forms.

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

Acid-Base Reactions and Net Ionic Equations

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

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

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

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

Electrolytes

Electrolytes are substances that conduct electricity when dissolved in water.

• Strong Electrolytes: Completely dissociate (e.g., NaCl, KNO3).

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

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

Types of Chemical Reactions

Oxidation-Reduction (Redox) Reactions

Redox reactions involve the transfer of electrons between species, resulting in changes in oxidation numbers.

• Assigning Oxidation Numbers: Use rules to determine the oxidation state of each atom in a compound.

• Identifying Redox Reactions: Look for changes in oxidation numbers between reactants and products.

• Oxidized: Species that loses electrons (increase in oxidation number).

• Reduced: Species that gains electrons (decrease in oxidation number).

• Balancing Redox Equations: Use the half-reaction method to balance electrons lost and gained.

• Example: $\text{Zn} (s) + \text{Cu}^{2+} (aq) \rightarrow \text{Zn}^{2+} (aq) + \text{Cu} (s)$ Zn is oxidized, Cu2+ is reduced.

Thermochemistry

Calorimetry

Calorimetry measures the heat absorbed or released during a chemical or physical process.

• Formula: $q = m \times c \times \Delta T$ Where $q$ = heat (J), $m$ = mass (g), $c$ = specific heat (J/g·°C), $\Delta T$ = temperature change (°C).

• Example: 100 g of water ($c = 4.18$ J/g·°C) heated from 25°C to 35°C: $q = 100 \times 4.18 \times (35-25) = 4180$ J.

Hess's Law

Hess's Law states that the total enthalpy change for a reaction is the sum of the enthalpy changes for individual steps.

• Application: Manipulate and add equations to find $\Delta H$ for a target reaction.

• Example: If $A \rightarrow B$ ($\Delta H_1$) and $B \rightarrow C$ ($\Delta H_2$), then $A \rightarrow C$ has $\Delta H = \Delta H_1 + \Delta H_2$.

Combustion Analysis and Balancing

Combustion Reactions

Combustion involves the reaction of a substance with oxygen to produce oxides and energy, often requiring careful balancing, especially with organic compounds.

• General Form: $\text{CxHy} + O_2 \rightarrow CO_2 + H_2O$

• Balancing Steps: Balance C, then H, then O atoms last.

• Example: $C_3H_8 + 5O_2 \rightarrow 3CO_2 + 4H_2O$

Recommended Study Strategy

• Memorize Key Formulas: $M = n/V$, $q = mc\Delta T$, and Hess's Law equations.

• Practice Limiting Reactant Problems: Work through step-by-step calculations to identify the limiting reactant and calculate product yield.

• Master Solution Stoichiometry: Convert between volume, molarity, and mass using molar mass and concentration formulas.

• Review Oxidation Number Rules: Practice assigning oxidation numbers to various compounds and ions.

• Focus on Open-Ended Questions: Practice showing all steps in calculations for stoichiometry, titration, and redox problems.