General Chemistry Study Guide: Periodic Trends, Lewis Structures, Molecular Geometry, and Bonding

Periodic Trends

Core and Valence Electrons; Effective Nuclear Charge

Understanding the distribution of electrons in atoms is essential for predicting chemical behavior. The effective nuclear charge (Zeff) is the net positive charge experienced by valence electrons, accounting for shielding by core electrons.

• Core electrons: Electrons in filled inner shells that shield valence electrons from the nucleus.

• Valence electrons: Electrons in the outermost shell, involved in chemical bonding.

• Effective Nuclear Charge Formula: $Z_{eff} = Z - S$, where Z is the atomic number and S is the number of core electrons.

• Example: For Mg (atomic number 12), core electrons = 10, valence electrons = 2, so $Z_{eff} = 12 - 10 = 2$.

Atomic Radius Trends

Atomic radius varies across the periodic table due to changes in nuclear charge and electron shielding.

• Trend: Atomic radius decreases across a period and increases down a group.

• Ion Size: Anions are larger than their parent atoms; cations are smaller.

• Example: Rank F-, F, F+ from smallest to largest radius: F+ < F < F-.

Ionization Energy

Ionization energy is the energy required to remove an electron from an atom in the gas phase. Successive ionization energies increase as electrons are removed.

• High jumps in ionization energy indicate removal of a core electron.

• Example: If IE3 is much higher than IE2, the element has two valence electrons.

Electron Configuration and Magnetism

Electron configurations determine if an ion is paramagnetic (unpaired electrons) or diamagnetic (all electrons paired).

• Example: Mg2+ has a noble gas configuration (diamagnetic); Al has unpaired electrons (paramagnetic).

Electron Affinity

Electron affinity is the energy change when an atom gains an electron. It generally increases across a period.

• Trend: Halogens have the highest electron affinity.

• Example: Cl > Br > I; N < F < O.

Naming Compounds

Writing Names and Formulas

Chemical nomenclature follows systematic rules for naming compounds and writing formulas.

• Ammonium sulfide: $(NH_4)_2S$

• SF6: Sulfur hexafluoride

• Silver(I) nitrate: $AgNO_3$

• CuSO4: Copper(II) sulfate

Lewis Structures

Drawing Lewis Structures and Assigning Formal Charges

Lewis structures represent the arrangement of electrons in molecules and ions. Formal charges help identify the most stable structure.

• Steps:

  1. Count total valence electrons.

  2. Arrange atoms and connect with single bonds.

  3. Complete octets and assign formal charges: $\text{Formal charge} = \text{valence electrons} - (\text{lone electrons} + \frac{1}{2}\text{bonding electrons})$

  4. Draw resonance structures if applicable.

• Examples: CO, NH4+, SO3, O2, ClO-, NO-, NO2, SF4

Resonance Structures

Some molecules have multiple valid Lewis structures, called resonance structures. The true structure is a hybrid.

• Example: Azo Diazide (N3N2): Draw all possible resonance forms, ensuring no rings/cyclic structures.

Molecular Geometry and Polarity

Electron Geometry, Molecular Geometry, Bond Angles, Dipoles, and Polarity

The shape of a molecule is determined by the arrangement of electron pairs around the central atom (VSEPR theory).

• Electron geometry: Arrangement of all electron groups (bonding and lone pairs).

• Molecular geometry: Arrangement of atoms (ignoring lone pairs).

• Bond angles: Typical angles for common geometries (e.g., tetrahedral: 109.5°).

• Dipole moment: Separation of charge in a molecule; determines polarity.

• Polarity: Molecules with net dipole moments are polar.

• Examples: BCl3 (trigonal planar, nonpolar), ClO2- (bent, polar), H2S (bent, polar), KrF2 (linear, nonpolar), HCN (linear, polar), XeO4 (tetrahedral, nonpolar).

Sigma and Pi Bonds / Valence Bond Theory

Sigma and Pi Bonds

Covalent bonds are classified as sigma (σ) or pi (π) bonds. Sigma bonds result from head-on orbital overlap; pi bonds result from side-to-side overlap.

• Single bond: 1 sigma bond

• Double bond: 1 sigma + 1 pi bond

• Triple bond: 1 sigma + 2 pi bonds

• Example: In the given structure, count the number of sigma and pi bonds and identify hybridization of numbered atoms.

Valence Bond Theory and Hybridization

Valence bond theory explains bonding using hybrid orbitals formed from atomic orbitals.

• Hybridization: Mixing of s, p, and sometimes d orbitals to form equivalent hybrid orbitals.

• Common types: sp (linear), sp2 (trigonal planar), sp3 (tetrahedral).

• Example: BeCl2 (sp hybridization), SCl2 (sp3 hybridization).

Orbital Diagrams

Orbital diagrams visually represent electron configurations and hybridization.

• Ground state: Electrons occupy lowest energy orbitals.

• Hybridized state: Orbitals mix to form hybrids (e.g., sp, sp3).

• Example: Fill in orbital diagrams for beryllium and sulfur in both ground and hybridized states.

Atom	Ground State Orbitals	Hybridized Orbitals
Beryllium	1s2 2s2	sp hybrids
Sulfur	1s2 2s2 2p6 3s2 3p4	sp3 hybrids

Additional info: Academic context and examples have been added to clarify and expand upon the original questions, making the guide self-contained and suitable for exam preparation.