General Chemistry I: Practice Problems and Key Concepts for Exam 2

Practice Problems for General Chemistry I: Exam 2 Review

Naming Compounds and Writing Formulas

Understanding how to name chemical compounds and write their formulas is fundamental in general chemistry. This includes ionic compounds, hydrates, and molecular compounds.

• Ionic Compounds: Composed of cations (usually metals) and anions (nonmetals or polyatomic ions). The name typically includes the cation first, followed by the anion.

• Hydrates: Compounds that include water molecules in their crystalline structure. The number of water molecules is indicated by a prefix (e.g., heptahydrate for 7 H2O).

• Molecular Compounds: Composed of nonmetals; prefixes indicate the number of each atom (e.g., dichlorine heptoxide for Cl2O7).

Examples:

• CuCl2·6H2O: Copper(II) chloride hexahydrate

• CuSO4·5H2O: Copper(II) sulfate pentahydrate

• MgCl2: Magnesium chloride

• HNO3: Nitric acid

• P4O10: Tetraphosphorus decoxide

• Cl2O7: Dichlorine heptoxide

Stoichiometry and Chemical Quantities

Stoichiometry involves calculations based on chemical equations, including determining moles, mass, and percent composition.

• Mole Concept: The mole is a counting unit for atoms, molecules, or ions. Avogadro's number ($6.022 \times 10^{23}$) is used to convert between moles and number of particles.

• Percent Composition: The percent by mass of each element in a compound.

• Empirical and Molecular Formulas: The empirical formula gives the simplest whole-number ratio of atoms; the molecular formula gives the actual number of atoms in a molecule.

Key Equations:

• Moles = $\frac{\text{mass (g)}}{\text{molar mass (g/mol)}}$

• Number of particles = moles $\times$ Avogadro's number

• Percent composition = $\frac{\text{mass of element}}{\text{total mass}} \times 100\%$

Example: Calculate the number of moles of glucose (C6H12O6) in a 1.80 g sample.

Lewis Structures, Formal Charge, and Molecular Geometry

Lewis structures represent the arrangement of electrons in molecules. Formal charge helps determine the most stable structure. Molecular geometry describes the 3D arrangement of atoms.

• Lewis Structure: Shows all valence electrons as dots or lines (bonds).

• Formal Charge: $\text{Formal charge} = \text{valence electrons} - \text{nonbonding electrons} - \frac{1}{2}(\text{bonding electrons})$

• Electron Groups: Regions of electron density (bonds or lone pairs) around a central atom.

• Molecular Geometry: Determined by the number of electron groups and lone pairs (VSEPR theory).

Examples:

• CO2: Linear geometry, 2 electron groups, 180° bond angle.

• BF5: Square pyramidal geometry, 6 electron groups.

• SO2: Bent geometry, 3 electron groups.

Bonding and Hybridization

Bonding involves the sharing or transfer of electrons. Hybridization explains the mixing of atomic orbitals to form new hybrid orbitals for bonding.

• Sigma (σ) Bonds: Single covalent bonds formed by head-on overlap of orbitals.

• Pi (π) Bonds: Formed by side-on overlap of p orbitals; present in double and triple bonds.

• Hybridization: The process of mixing atomic orbitals to form new hybrid orbitals (e.g., sp, sp2, sp3).

Example: In the structure of Tylenol, count the number of sigma bonds present.

Molecular Orbitals and Bond Order

Molecular orbital theory describes the distribution of electrons in molecules. Bond order indicates the strength and stability of a bond.

• Molecular Orbitals: Formed by the combination of atomic orbitals from bonded atoms.

• Bond Order: $\text{Bond order} = \frac{1}{2}(\text{number of bonding electrons} - \text{number of antibonding electrons})$

Example: Calculate the bond order for N2 using its molecular orbital diagram.

Table: Lewis Structures and Molecular Properties

The following table summarizes the Lewis structures, formal charges, electron groups, molecular geometry, hybridization, and dipole moments for selected molecules:

Molecule	Lewis structure with formal charges	# of electron groups	Molecular geometry	Hybridization	Dipole moment (yes/no)
SO2	O=S=O (with one lone pair on S, formal charges: S=0, O=0)	3	Bent	sp2	Yes
NO3-	Resonance structures with one N=O and two N–O bonds, formal charges: N=0, O=0 or -1	3	Trigonal planar	sp2	No
SF4	Central S with four F atoms and one lone pair, formal charges: S=0, F=0	5	Seesaw	sp3d	Yes

Additional info: Table entries inferred based on standard Lewis structures and VSEPR theory.

Key Concepts and Applications

• Assigning Formal Charges: Used to determine the most stable Lewis structure.

• Identifying Covalent Bonds: Covalent bonds are typically found between nonmetals.

• Predicting Molecular Shape: Use VSEPR theory to predict geometry based on electron groups.

• Hybridization and Bonding: The type of hybridization affects molecular geometry and bond angles.

Sample Problems

1. Calculate the number of oxygen atoms in 32.5 g of CaCO3.

2. Determine the mass in kg of $6.89 \times 10^{22}$ molecules of CO2 (molar mass = 44.01 g/mol).

3. Determine the empirical formula for a compound with 52.14% C, 34.73% O, and 13.13% H by mass.

4. Draw Lewis structures and predict geometry, hybridization, and dipole moment for SO2, NO3-, and SF4.

5. Distribute valence electrons in N2 using molecular orbitals, calculate bond order, and write electron configuration.

Additional info: These problems cover key skills in chemical nomenclature, stoichiometry, molecular structure, and bonding, which are essential for success in General Chemistry I.