General Chemistry Study Guide: Chemical Reactions, Thermochemistry, and Solutions

Chapter 7: Chemical Reactions & Quantities

Balancing Chemical Equations

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

• Definition: A balanced chemical equation has equal numbers of each type of atom on both sides.

• Steps:

  1. Write the unbalanced equation.

  2. Count the number of atoms of each element on both sides.

  3. Add coefficients to balance the atoms.

  4. Check your work.

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

Stoichiometry and Mole Ratios

Stoichiometry involves using balanced chemical equations to calculate the relationships between reactants and products in chemical reactions.

• Mole Ratio: The ratio of moles of one substance to another in a balanced equation.

• Example: In $2H_2 + O_2 \rightarrow 2H_2O$, the mole ratio of $H_2$ to $O_2$ is 2:1.

• Application: Used to determine how much of a reactant is needed or how much product will form.

Limiting Reactant and Theoretical Yield

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

• Definition: The reactant that determines the maximum amount of product possible.

• Theoretical Yield: The maximum amount of product that can be formed from the limiting reactant.

• Example: If 3 mol $A$ reacts with 2 mol $B$ to produce $C$, and the reaction requires 1 mol $A$ per 1 mol $B$, $B$ is the limiting reactant.

Percent Yield

Percent yield compares the actual yield to the theoretical yield.

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

• Example: If the theoretical yield is 10 g and the actual yield is 8 g, percent yield is $80\%$.

Chapter 9: Thermochemistry

Energy, Work, and Heat

Thermochemistry studies the energy changes that occur during chemical reactions and changes of state.

• System and Surroundings: The system is the part of the universe being studied; the surroundings are everything else.

• Internal Energy ($E$): The sum of all kinetic and potential energies of the system.

• Work ($w$): Energy used to move an object against a force.

• Heat ($q$): Energy transferred due to temperature difference.

• Formula: $\Delta E = q + w$

Enthalpy ($\Delta H$)

Enthalpy is the heat content of a system at constant pressure.

• Definition: $\Delta H$ is the change in enthalpy, equal to heat at constant pressure.

• Formula: $\Delta H = H_{products} - H_{reactants}$

• Exothermic Reaction: Releases heat ($\Delta H < 0$).

• Endothermic Reaction: Absorbs heat ($\Delta H > 0$).

Calorimetry

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

• Specific Heat Capacity ($c$): Amount of heat required to raise the temperature of 1 g of a substance by 1°C.

• Formula: $q = m \cdot c \cdot \Delta T$

• Example: If 100 g of water is heated from 20°C to 30°C, $q = 100 \times 4.18 \times 10 = 4180$ J.

Hess's Law

Hess's Law states that the total enthalpy change for a reaction is the same, no matter how many steps the reaction is carried out in.

• Formula: $\Delta H_{rxn} = \sum \Delta H_{products} - \sum \Delta H_{reactants}$

• Application: Used to calculate enthalpy changes for reactions not easily measured directly.

Chapter 8: Solutions & Aqueous Reactions

Solubility Rules

Solubility rules help predict whether an ionic compound will dissolve in water.

Ion	Solubility	Exceptions
Li+, Na+, K+, NH4+	soluble	none
Cl-, Br-, I-	soluble	Ag+, Pb2+, Hg22+
SO42-	soluble	Ca2+, Sr2+, Ba2+, Pb2+
CO32-, PO43-	insoluble	Li+, Na+, K+, NH4+
OH-	insoluble	Li+, Na+, K+, NH4+, Ca2+, Sr2+, Ba2+

Types of Chemical Reactions in Solution

Common types of reactions in aqueous solution include precipitation, acid-base, and redox reactions.

• Precipitation Reaction: Forms an insoluble product (precipitate).

• Acid-Base Reaction: Involves transfer of protons ($H^+$).

• Redox Reaction: Involves transfer of electrons.

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

Electrolytes

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

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

• Weak Electrolyte: Partially dissociates (e.g., acetic acid).

• Nonelectrolyte: Does not produce ions (e.g., sugar).

Oxidation States and Redox Reactions

Oxidation states indicate the degree of oxidation of an atom in a compound. Redox reactions involve changes in oxidation states.

• Rules for Assigning Oxidation States:

  1. Elemental form: 0

  2. Monatomic ion: charge of ion

  3. Group 1A: +1; Group 2A: +2

  4. Fluorine: -1; Oxygen: -2 (except in peroxides); Hydrogen: +1 (except in metal hydrides)

  5. Sum of oxidation states in a neutral compound is 0; in a polyatomic ion, equals the ion charge

• Example: In $H_2O$, H is +1, O is -2.

Reference Data and Useful Equations

Periodic Table

The periodic table organizes elements by increasing atomic number and groups elements with similar chemical properties.

• Groups: Vertical columns; elements in the same group have similar properties.

• Periods: Horizontal rows; properties change progressively across a period.

Useful Equations

• $q = \text{specific heat} \times \text{mass} \times \Delta T$

• $\Delta E = q + w$

• $w = -P\Delta V$

• $\Delta H = \Delta E + P\Delta V$

• $\Delta H_{rxn} = \sum \Delta H_{f,\text{products}} - \sum \Delta H_{f,\text{reactants}}$

Additional info:

• Questions cover a range of topics from chemical equations, stoichiometry, thermochemistry, solution chemistry, and redox reactions, all central to a General Chemistry course.

• Tables and diagrams (e.g., periodic table, solubility rules) are included for reference and application in problem-solving.