Exam 2 Study Guide: Chemical Bonding, Compounds, Chemical Reactions, and Solutions

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Chemical bonds are the forces that hold atoms together in compounds. The energy required to break a bond is called bond energy.

Bond Energy: The amount of energy needed to break one mole of a specific bond in a molecule in the gas phase.
Calculation: To estimate the enthalpy change () for a reaction, sum the bond energies of bonds broken and subtract the bond energies of bonds formed:
Example: Breaking H2 and forming HCl in the reaction:

Compounds are classified based on the nature of their bonding.

Covalent Compounds: Formed by sharing electrons between nonmetal atoms.
Ionic Compounds: Formed by transfer of electrons from metal to nonmetal, resulting in positive and negative ions held together by electrostatic forces.
Atomic Elements: Elements whose particles are single atoms (e.g., Ne, Fe).
Molecular Elements: Elements whose particles are molecules (e.g., O2, N2).
Ionic Compounds: Composed of ions (e.g., NaCl).
Molecular Compounds: Composed of molecules (e.g., H2O).

Nomenclature is the system for naming chemical compounds.

Ionic Compounds: Name the cation first, then the anion. For transition metals, indicate the charge with Roman numerals (e.g., FeCl2: iron(II) chloride).
Covalent Compounds: Use prefixes to indicate the number of atoms (e.g., CO2: carbon dioxide).
Acids: If the anion ends in -ide, use "hydro-" and "-ic" (e.g., HCl: hydrochloric acid). If the anion ends in -ate or -ite, use "-ic" or "-ous" (e.g., HNO3: nitric acid).
Hydrates: Name the compound, then add the prefix for the number of water molecules and "hydrate" (e.g., CuSO4·5H2O: copper(II) sulfate pentahydrate).

The chemical formula shows the types and numbers of atoms in a compound. The formula mass is the sum of atomic masses in the formula.

Molar Mass: The mass of one mole of a compound, calculated by summing the atomic masses of all atoms in the formula.
Example: Molar mass of H2O: g/mol

Formulas represent the composition of compounds in different ways.

Empirical Formula: Shows the simplest whole-number ratio of atoms (e.g., CH2O).
Molecular Formula: Shows the actual number of atoms in a molecule (e.g., C6H12O6).
Structural Formula: Shows the arrangement of atoms (e.g., H–O–H for water).
Writing Ionic Formulas: Use ion charges to balance for electrical neutrality. Example: Na+ and Cl- combine to form NaCl.

Quantities of compounds can be converted between mass, moles, and number of molecules.

Mass percent composition shows the percentage by mass of each element in a compound.

Chemical reactions are represented by chemical equations, which must be balanced to obey the law of conservation of mass.

Chemical Equation: Shows reactants and products, e.g.,
Balancing: Adjust coefficients to ensure equal numbers of each atom on both sides.

Stoichiometry is the calculation of reactant and product quantities in chemical reactions.

Mole Relationships: Use balanced equations to relate moles of reactants and products.
Empirical and Molecular Formulas from Data: Use experimental mass or percent composition to determine formulas.
Limiting Reactant: The reactant that is completely consumed first, limiting the amount of product formed.
Theoretical Yield: The maximum amount of product that can be formed from given reactants.
Percent Yield:
Calculating Amounts: Use stoichiometric ratios to calculate masses, moles, or molecules of reagents, products, and excess reactants.
Example: If 10 g of A reacts with 15 g of B, determine the limiting reactant and calculate the theoretical yield of product C.

Mass percent composition is used to determine the empirical formula from experimental data.

Calculation: Convert mass percent to grams, then to moles, and find the simplest ratio.

Solutions are homogeneous mixtures of solute and solvent.

Solution: Homogeneous mixture of two or more substances.
Solvent: The substance present in the greatest amount (usually water in aqueous solutions).
Solute: The substance dissolved in the solvent.
Aqueous Solution: Solution where water is the solvent.
Molarity (M):

Molarity is a measure of concentration, used to quantify the amount of solute in a solution.

When a solution is diluted, the amount of solute remains constant, but the volume increases.

Dilution Equation:
Example: To dilute 100 mL of 1.0 M solution to 0.5 M, add enough water to make the final volume 200 mL.

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