General Chemistry Study Guide: Chemical Reactions, Stoichiometry, Solutions, and Gases

Chapter 4. Chemical Reactions and Chemical Quantities

4.2 Writing and Balancing Chemical Equations

Chemical equations are symbolic representations of chemical reactions, showing the reactants and products involved. Balancing these equations ensures the conservation of mass and atoms.

• Reactants are substances that undergo change during a reaction; products are the substances formed.

• A chemical equation uses chemical formulas to represent the reactants and products.

• Balancing equations requires equal numbers of each atom on both sides.

• Example: $2H_2 + O_2 \rightarrow 2H_2O$

4.3 Reaction Stoichiometry: How Much Carbon Dioxide?

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

• Defines the mole relationships among chemical amounts in a balanced reaction.

• Calculates the mass of a reactant needed or the mass of product formed.

• Example: Using $C_6H_{12}O_6 + 6O_2 \rightarrow 6CO_2 + 6H_2O$ to find moles of $CO_2$ produced from glucose.

4.4 Stoichiometric Relationships: Limiting Reactant, Theoretical Yield, Percent Yield, and Reactant in Excess

Stoichiometry also determines which reactant limits the reaction, the maximum possible yield, and the efficiency of the reaction.

• Limiting reactant: The reactant that is completely consumed first, limiting the amount of product.

• Theoretical yield: The maximum amount of product possible from the limiting reactant.

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

• Reactant in excess: The reactant left over after the reaction is complete.

4.5 Three Examples of Chemical Reactions: Combustion, Alkali Metals, and Halogens

Chemical reactions can be classified by the types of reactants and products involved. Combustion, reactions of alkali metals, and halogen reactions are common examples.

• Combustion reactions involve a substance reacting with oxygen to produce energy, often forming $CO_2$ and $H_2O$.

• Alkali metals react vigorously with water to produce hydrogen gas and a metal hydroxide.

• Halogens react with metals to form ionic halide salts.

• Example: $2Na + Cl_2 \rightarrow 2NaCl$

Chapter 5. Solution Chemistry

5.2 Solution Concentration

Concentration describes the amount of solute dissolved in a given quantity of solvent. Molarity is a common unit.

• Solution: Homogeneous mixture of solute and solvent.

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

• Calculating concentration after dilution: $M_1V_1 = M_2V_2$

5.3 Solution Stoichiometry

Solution stoichiometry uses concentration and volume to determine the amounts of reactants and products in reactions involving solutions.

• Calculate moles of reactants/products using molarity and volume.

• Example: How many moles of $NaCl$ are in 0.5 L of 1.0 M $NaCl$? $0.5 \times 1.0 = 0.5$ moles.

5.4 Types of Aqueous Solutions and Solubility

Aqueous solutions can contain electrolytes or nonelectrolytes, and substances can be classified as soluble or insoluble.

• Electrolyte: Substance that conducts electricity when dissolved in water.

• Strong electrolyte: Completely dissociates in water (e.g., $NaCl$).

• Weak electrolyte: Partially dissociates (e.g., $CH_3COOH$).

• Nonelectrolyte: Does not conduct electricity (e.g., sugar).

• Solubility rules help predict which compounds dissolve in water.

5.5 Precipitation Reactions

Precipitation reactions occur when two solutions are mixed and an insoluble product (precipitate) forms.

• Precipitate: Solid formed from a reaction in solution.

• Use solubility rules to predict precipitate formation.

• Example: $AgNO_3 + NaCl \rightarrow AgCl (s) + NaNO_3$

5.6 Representing Aqueous Reactions: Molecular, Ionic, and Complete Ionic Equations

Chemical reactions in solution can be represented in different ways to show the species involved.

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

• Complete ionic equation: Shows all strong electrolytes as ions.

• Net ionic equation: Shows only the species that actually change during the reaction.

• Example: Net ionic for precipitation of $AgCl$: $Ag^+ (aq) + Cl^- (aq) \rightarrow AgCl (s)$

5.7 Acid-Base and Gas-Evolution Reactions

Acid-base reactions involve the transfer of protons, while gas-evolution reactions produce a gaseous product.

• Recognize diprotic and polyprotic acids.

• Identify strong and weak acids and bases.

• Example: $HCl + NaOH \rightarrow NaCl + H_2O$

Chapter 6. Gases

6.2 Pressure: The Result of Molecular Collisions

Pressure is caused by the collisions of gas molecules with the walls of their container.

• Measured in units such as mmHg, atm, and Pa.

• Example: Atmospheric pressure is about 1 atm at sea level.

6.3 The Simple Gas Laws: Boyle's Law, Charles's Law, and Avogadro's Law

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

• Boyle's Law: $P_1V_1 = P_2V_2$ (at constant T and n)

• Charles's Law: $\frac{V_1}{T_1} = \frac{V_2}{T_2}$ (at constant P and n)

• Avogadro's Law: $\frac{V_1}{n_1} = \frac{V_2}{n_2}$ (at constant P and T)

6.4 The Ideal Gas Law

The ideal gas law combines the simple gas laws into one equation relating pressure, volume, temperature, and amount of gas.

• $PV = nRT$

• R is the gas constant: $0.0821 \frac{L \cdot atm}{mol \cdot K}$

• Used to calculate unknown properties of gases.