General Chemistry Final Exam Review: Key Concepts and Practice

Final Exam Review: General Chemistry

Significant Figures and Measurement

Understanding significant figures is essential for reporting measurements accurately in chemistry.

• Significant Figures: The digits in a measurement that are known with certainty plus one estimated digit.

• Rules:

  • All nonzero digits are significant.

  • Zeros between nonzero digits are significant.

  • Leading zeros are not significant.

  • Trailing zeros are significant only if there is a decimal point.

• Example: The number 1.00243 has 6 significant figures.

Atoms, Moles, and Avogadro's Number

Chemists use the mole to count atoms, molecules, or ions. Avogadro's number ($6.022 \times 10^{23}$) is the number of particles in one mole.

• Calculating Atoms:

  • Number of atoms = (moles) × (Avogadro's number)

  • Example: How many atoms of Cu are in 5.81 × 10-2 g of copper?

Wavelength, Frequency, and the Electromagnetic Spectrum

The relationship between wavelength ($\lambda$), frequency ($\nu$), and the speed of light ($c$) is fundamental in quantum chemistry.

• Equation: $c = \lambda \nu$

• Example: To find the wavelength of light with a frequency of $6.52 \times 10^{14}$ s-1:

• $\lambda = \frac{c}{\nu}$

Quantum Numbers and Electron Configuration

Quantum numbers describe the properties of atomic orbitals and the electrons in them.

• Principal quantum number (n): Energy level (n = 1, 2, 3, ...)

• Angular momentum quantum number (l): Shape of orbital (l = 0 to n-1)

• Magnetic quantum number (ml): Orientation (ml = -l to +l)

• Spin quantum number (ms): Spin direction (±1/2)

• Example: A 4p electron: n = 4, l = 1, ml = -1, 0, or 1, ms = ±1/2

Periodic Trends

Understanding periodic trends helps predict element properties.

• Atomic Radius: Increases down a group, decreases across a period.

• Ionization Energy: Energy required to remove an electron; increases across a period, decreases down a group.

• Example: Which element has the largest atomic radius?

Ionic Compounds and Formulas

Ionic compounds are formed from metals and nonmetals. Their formulas must be electrically neutral.

• Example: Which of the following ionic compound formulas is incorrect?

Lewis Structures and Formal Charge

Lewis structures represent the arrangement of electrons in molecules. Formal charge helps determine the most stable structure.

• Formal Charge Equation: $\text{Formal charge} = \text{Valence electrons} - \text{Nonbonding electrons} - \frac{1}{2}(\text{Bonding electrons})$

• Example: Draw the Lewis structure for NO3- and calculate the formal charge on N.

Molecular Geometry and Bond Angles

The shape of a molecule is determined by the number of bonding and lone pairs on the central atom (VSEPR theory).

• Common Geometries:

  • Linear: 180°

  • Trigonal planar: 120°

  • Tetrahedral: 109.5°

  • Trigonal bipyramidal: 90°, 120°

  • Octahedral: 90°

• Example: What are the bond angles in a molecule with three bonding groups and zero lone pairs?

Polarity of Molecules

Molecular polarity depends on the difference in electronegativity and the symmetry of the molecule.

• Polar Molecule: Has a net dipole moment due to unequal sharing of electrons.

• Nonpolar Molecule: Symmetrical distribution of charge.

• Example: Which of the following molecules is polar?

Hybridization

Hybridization describes the mixing of atomic orbitals to form new hybrid orbitals for bonding.

• sp: Linear geometry

• sp2: Trigonal planar geometry

• sp3: Tetrahedral geometry

• Example: What is the hybridization of the central atom in PH3?

Stoichiometry and Chemical Reactions

Stoichiometry involves calculations based on balanced chemical equations.

• Example: Sodium azide (NaN3) decomposes to produce nitrogen gas for airbags:

• $2\text{NaN}_3 (s) \rightarrow 2\text{Na} (s) + 3\text{N}_2 (g)$

• Calculate the mass of sodium azide needed to produce a certain mass of N2.

Solutions and Molarity

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

• Equation: $M = \frac{\text{moles of solute}}{\text{liters of solution}}$

• Example: What is the molarity of a solution made by dissolving 7.5 g of CaCl2 in enough water to make 150 mL of solution?

Electron Configuration and Ions

Electron configuration describes the arrangement of electrons in an atom or ion.

• Example: How many electrons does a Fe3+ ion have?

Covalent Bond Polarity

A covalent bond is polar if the atoms have different electronegativities.

• Example: Which of the following pairs of atoms will form a polar covalent bond?

Molecular Orbital Theory

Molecular orbital diagrams show the combination of atomic orbitals to form molecular orbitals, which can be bonding or antibonding.

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

• Paramagnetism: Molecules with unpaired electrons are paramagnetic.

• Example: Use the molecular orbital diagram to determine bond order and paramagnetism.

Successive Ionization Energies

Successive ionization energies increase as more electrons are removed, with a large jump when removing a core electron.

• Example: Use ionization energy data to determine the group of an element.

Chemiluminescence and Limiting Reactants

Chemiluminescence is light produced by a chemical reaction. The limiting reactant is the reactant that is completely consumed first.

• Example: Calculate the limiting reactant in a reaction involving luminol.

Isotopes and Atomic Mass

Isotopes are atoms of the same element with different numbers of neutrons. The average atomic mass is calculated using the relative abundances and masses of isotopes.

• Equation: $\text{Average atomic mass} = \sum (\text{isotope mass} \times \text{fractional abundance})$

• Example: Calculate the atomic mass of rubidium given two isotopes and their abundances.

Naming and Writing Chemical Formulas

Chemical compounds are named according to IUPAC rules. The chemical formula represents the types and numbers of atoms in a compound.

Name	Chemical Formula
Magnesium hydroxide	Mg(OH)2
Sodium nitrate	NaNO3
Vanadium (III) bromide	VBr3
Dimethyl tetrasulfide	(CH3)2S4
Copper(II) nitrate	Cu(NO3)2

Empirical and Molecular Formulas

The empirical formula is the simplest whole-number ratio of atoms in a compound. The molecular formula is a multiple of the empirical formula.

• Example: Determine the empirical formula from percent composition data.

Gas Laws and Stoichiometry

Gas laws relate the pressure, volume, temperature, and amount of gas. Stoichiometry can be applied to reactions involving gases.

• Example: Calculate the mass of CO2 produced from a given mass of butane and oxygen.

Physical and Chemical Properties

Physical properties can be observed without changing the substance's identity. Chemical properties describe a substance's ability to undergo chemical changes.

• Example: Melting point is a physical property; reactivity with chlorine gas is a chemical property.

Pure Substances and Mixtures

Pure substances have a fixed composition; mixtures contain two or more substances physically combined.

• Example: Sodium hydrogen carbonate is a pure substance; baking powder is a mixture.

Acid-Base and Redox Reactions

Acid-base reactions involve proton transfer; redox reactions involve electron transfer.

• Example: Identify acid-base and redox reactions from given equations.

Lewis Structures, Geometry, and Polarity Practice

Drawing Lewis structures and determining geometry, bond angles, and polarity are essential skills.

• Practice: Draw Lewis structures for molecules/ions and determine:

  • Electron geometry

  • Molecular geometry

  • Bond angles

  • Polarity

  • Hybridization

Additional info:

• This review covers topics from Ch.1 (Introduction to Chemistry) through Ch.11 (Liquids, Solids, and Intermolecular Forces), Ch.13 (Solutions), Ch.14 (Kinetics), Ch.15 (Equilibrium), Ch.16 (Acid-Base), Ch.17 (Aqueous Equilibria), Ch.18 (Thermodynamics), Ch.19 (Electrochemistry), and Ch.20 (Nuclear Chemistry), as well as lab techniques and mathematical operations.

• Some questions require application of multiple concepts, such as stoichiometry, molecular structure, and periodic trends.