General Chemistry Exam 2 Study Guide: Atomic Structure, Bonding, and Chemical Formulas

Chapter 3: Atomic Structure and Periodicity

Valence Electron and Electron Configurations

Understanding electron configurations is essential for predicting chemical behavior and bonding. The Pauli exclusion principle, Aufbau principle, and Hund's Rule guide the arrangement of electrons in atomic orbitals.

• Pauli Exclusion Principle: No two electrons in an atom can have the same set of four quantum numbers.

• Aufbau Principle: Electrons fill the lowest energy orbitals first.

• Hund's Rule: Electrons occupy degenerate orbitals singly before pairing.

• Electron Configuration: The distribution of electrons among the orbitals of an atom. Example: The electron configuration of oxygen is 1s2 2s2 2p4.

• Orbital Diagram: A graphical representation showing the arrangement of electrons in orbitals.

Shielding, Zeff, and Penetration

Electrons in inner shells shield outer electrons from the full nuclear charge, affecting atomic properties.

• Shielding: The reduction in effective nuclear charge on the electron cloud due to inner electrons.

• Effective Nuclear Charge (Zeff): The net positive charge experienced by valence electrons. $Z_{eff} = Z - S$ where $Z$ is the atomic number and $S$ is the shielding constant.

• Penetration: The ability of an electron to get close to the nucleus, affecting its energy.

• Periodic Trends: Zeff increases across a period and decreases down a group.

Trends in Periodic Table: Ionization Energy, Electron Affinity, Atomic Radius

Periodic trends help predict element reactivity and properties.

• Ionization Energy: The energy required to remove an electron from a gaseous atom. Increases across a period, decreases down a group.

• Electron Affinity: The energy change when an electron is added to a gaseous atom. Generally becomes more negative across a period.

• Atomic Radius: The distance from the nucleus to the outermost electron. Decreases across a period, increases down a group.

• Successive Ionization Energies: Each subsequent electron removed requires more energy.

Magnetic Properties of Atoms

Atoms can exhibit magnetic behavior based on their electron configurations.

• Paramagnetic: Atoms with unpaired electrons; attracted to magnetic fields.

• Diamagnetic: Atoms with all electrons paired; weakly repelled by magnetic fields.

• Example: Oxygen is paramagnetic due to two unpaired electrons in its 2p orbitals.

Chapter 4: Chemical Nomenclature and Formulas

Nomenclature of Ionic and Covalent Compounds

Systematic naming of compounds ensures clear communication in chemistry.

• Ionic Compounds: Composed of cations and anions. Name the cation first, then the anion. Example: NaCl is sodium chloride.

• Polyatomic Ions: Charged species composed of multiple atoms. Example: $\text{SO}_4^{2-}$ is sulfate.

• Covalent Compounds: Use prefixes to indicate the number of atoms. Example: CO2 is carbon dioxide.

• Acids and Oxoacids: Named based on the anion. Example: HNO3 is nitric acid.

Mole Concept and Calculations

The mole is a fundamental unit for quantifying substances in chemistry.

• Mole: $6.022 \times 10^{23}$ entities (Avogadro's number).

• Molar Mass: Mass of one mole of a substance (g/mol).

• Conversions: $\text{mass} \rightarrow \text{moles} \rightarrow \text{number of particles}$

• Example: To find the number of molecules in 18 g of H2O: $\text{moles} = \frac{18\,g}{18.02\,g/mol} = 1\,mol$ $\text{molecules} = 1\,mol \times 6.022 \times 10^{23} = 6.022 \times 10^{23}$ molecules

Percent Composition and Empirical/Molecular Formulas

Determining the composition and formulas of compounds is crucial for chemical analysis.

• Percent Composition: The percentage by mass of each element in a compound. $\%\,\text{element} = \frac{\text{mass of element in 1 mol}}{\text{molar mass}} \times 100$

• Empirical Formula: The simplest whole-number ratio of atoms in a compound.

• Molecular Formula: The actual number of atoms of each element in a molecule.

• Combustion Analysis: Used to determine empirical formulas from combustion data (e.g., CO2 and H2O produced).

• Example: If a compound contains 40% C, 6.7% H, and 53.3% O, its empirical formula is CH2O.

Chapter 5: Chemical Bonding and Molecular Structure

Electronegativity and Periodic Table Trends

Electronegativity is the tendency of an atom to attract electrons in a chemical bond.

• Electronegativity: Increases across a period, decreases down a group. Fluorine is the most electronegative element.

• Periodic Table: Used to predict bond type and polarity.

Bond Types: Covalent, Polar Covalent, Ionic, and Bond Polarity

Bonds form between atoms to achieve stable electron configurations.

• Covalent Bond: Sharing of electron pairs between atoms.

• Polar Covalent Bond: Unequal sharing of electrons due to differences in electronegativity.

• Ionic Bond: Transfer of electrons from one atom to another, forming ions.

• Bond Polarity: Determined by the difference in electronegativity ($\Delta EN$). $\Delta EN = |EN_{A} - EN_{B}|$

• Example: HCl is polar covalent; NaCl is ionic.

Lewis Structures, Formal Charge, and Resonance

Lewis structures represent the arrangement of electrons in molecules.

• Lewis Structure: Shows bonding and lone pairs in a molecule.

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

• Resonance: Occurs when more than one valid Lewis structure can be drawn for a molecule.

• Example: Ozone (O3) has resonance structures.

Octet Rule and Exceptions

The octet rule states that atoms tend to have eight electrons in their valence shell, but exceptions exist.

• Octet Rule: Most atoms form bonds to achieve eight valence electrons.

• Exceptions: Hydrogen (2 electrons), Boron (6 electrons), expanded octets for elements in period 3 and beyond.

• Example: SF6 has 12 electrons around sulfur.

Bond Energies and Bond Lengths

Bond energy is the energy required to break a bond; bond length is the distance between nuclei of bonded atoms.

• Bond Energy: Higher for multiple bonds (triple > double > single).

• Bond Length: Shorter for multiple bonds.

• Example: C≡C bond is shorter and stronger than C=C or C–C.

Electronic and Molecular Geometry: VSEPR Model

The VSEPR (Valence Shell Electron Pair Repulsion) model predicts the shapes of molecules based on electron pair repulsion.

• VSEPR Model: Electron pairs arrange themselves to minimize repulsion.

• Common Geometries:

  • Linear: 180° bond angle (e.g., CO2)

  • Trigonal planar: 120° (e.g., BF3)

  • Tetrahedral: 109.5° (e.g., CH4)

  • Trigonal bipyramidal: 90°, 120° (e.g., PCl5)

  • Octahedral: 90° (e.g., SF6)

• Bond Angles: Determined by the number of bonding and lone pairs.

Summary Table: Bond Types and Properties

Bond Type	Electron Sharing	Polarity	Example
Covalent	Shared equally	Nonpolar	H2, O2
Polar Covalent	Shared unequally	Polar	HCl, H2O
Ionic	Transferred	Very polar	NaCl, KBr

Additional info: Academic context and examples have been added to expand on the brief points in the original study guide.