General Chemistry Study Guide: Atomic Structure, Periodic Trends, Chemical Bonding, and Molecular Geometry

Atomic Structure and Periodic Properties

Electron Configuration

Electron configuration describes the arrangement of electrons in an atom's orbitals. It follows the Aufbau principle, Pauli exclusion principle, and Hund's rule.

• Aufbau Principle: Electrons fill orbitals starting with the lowest energy level.

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

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

• Example: The ground state electron configuration for Rb (Rubidium) is $[Kr]5s^1$.

• Example: The ground state electron configuration for Zn^{2+} is $[Ar]3d^{10}$.

Ionization Energy

Ionization energy is the energy required to remove an electron from a gaseous atom or ion.

• Trend: Increases across a period (left to right), decreases down a group.

• Highest First Ionization Energy: Noble gases and elements on the right of the periodic table have the highest values.

• Example: Among Si, S, P, Cl, Mg, Cl has the highest first ionization energy.

Atomic and Ionic Radius

Atomic radius is the distance from the nucleus to the outermost electron shell. Ionic radius refers to the size of an ion.

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

• Ionic Radius: Anions are larger than their parent atoms; cations are smaller.

• Example: Among Na+, Ga3+, K+, Mg2+, Ca2+, K+ has the largest radius.

Electron Affinity

Electron affinity is the energy change when an atom gains an electron.

• Trend: Becomes more negative across a period; halogens have the most negative electron affinity.

• Example: Among B, C, N, O, F, F has the most negative electron affinity.

Chemical Bonding and Molecular Structure

Ionic Compounds: Cations and Anions

Ionic compounds consist of positively charged cations and negatively charged anions. The ratio of ions balances the overall charge.

Compound	Cations	Number of Cations	Anions	Number of Anions
KF	K+	1	F-	1
Mg3N2	Mg2+	3	N3-	2
Na2O	Na+	2	O2-	1
SnCl4	Sn4+	1	Cl-	4
PbI2	Pb2+	1	I-	2
Fe2O3	Fe3+	2	O2-	3
FeCl3	Fe3+	1	Cl-	3
CuNO3	Cu+	1	NO3-	1
Mg[Cl(H2O)6]2	Mg2+	1	Cl(H2O)6-	2
NH4I	NH4+	1	I-	1
NaBrO4	Na+	1	BrO4-	1
Fe(OH)3	Fe3+	1	OH-	3
KClO3	K+	1	ClO3-	1
Na3PO4	Na+	3	PO43-	1
CaF2	Ca2+	1	F-	2
CaCO3	Ca2+	1	CO32-	1

Formal Charge and Resonance

Formal charge helps determine the most stable Lewis structure for a molecule. Resonance structures are alternative Lewis structures for the same molecule.

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

• Best Lewis Structure: The structure with the lowest formal charges and negative charges on the most electronegative atoms is preferred.

• Resonance: SO2 has 2 resonance structures obeying the octet rule.

Bonding and Molecular Geometry

Molecular geometry describes the three-dimensional arrangement of atoms in a molecule, determined by the number of bonding and lone pairs around the central atom.

• VSEPR Theory: Electron pairs around a central atom arrange themselves to minimize repulsion.

• Common Geometries:

  • Linear: 180° bond angle

  • Trigonal planar: 120° bond angle

  • Tetrahedral: 109.5° bond angle

  • Trigonal bipyramidal: 90°, 120° bond angles

  • Octahedral: 90° bond angles

• Example: SF6 has an octahedral geometry with 90° bond angles.

• Example: H2O has a bent geometry due to two lone pairs on oxygen.

Hybridization

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

• sp: Linear geometry

• sp2: Trigonal planar geometry

• sp3: Tetrahedral geometry

• sp3d: Trigonal bipyramidal geometry

• sp3d2: Octahedral geometry

• Example: The central atom in PH3 is sp3 hybridized.

Sigma and Pi Bonds

Sigma (σ) bonds are single covalent bonds formed by head-on overlap of orbitals. Pi (π) bonds are formed by side-to-side overlap and exist in double and triple bonds.

• Single bond: 1 σ bond

• Double bond: 1 σ and 1 π bond

• Triple bond: 1 σ and 2 π bonds

• Example: Benzene (C6H6) has 12 σ bonds and 3 π bonds.

Polarity of Molecules

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

• Polar Molecule: Has a net dipole moment due to asymmetrical charge distribution.

• Nonpolar Molecule: Symmetrical molecules with no net dipole moment.

• Example: CH4 is nonpolar; HCN is polar.

Naming and Writing Formulas for Compounds

Nomenclature

Chemical nomenclature is the system for naming chemical compounds.

Formula	Name
NiSO4	Nickel(II) sulfate
	Phosphorous pentafluoride
Li2S	Lithium sulfide
	Aluminum carbonate
N2O4	Dinitrogen tetroxide
	Copper(I) fluoride

Lewis Structures and Expanded Octet

Lewis Structures

Lewis structures represent the arrangement of electrons in a molecule, showing bonds and lone pairs.

• Octet Rule: Atoms tend to have eight electrons in their valence shell.

• Expanded Octet: Elements in period 3 and beyond can have more than eight electrons (e.g., SF6, XeF4).

• Example: SF4 has a central atom with an expanded octet.

Practice: Drawing Lewis Structures

Practice drawing Lewis structures for the following molecules:

• CF4

• NF3

• SF4

• BH3

• NO

• ICL2-

• OPBr3

• XeF4

Summary Table: Key Periodic Trends

Property	Across a Period	Down a Group
Atomic Radius	Decreases	Increases
Ionization Energy	Increases	Decreases
Electron Affinity	Becomes more negative	Less negative
Electronegativity	Increases	Decreases

Additional info:

• Some compound names and formulas in tables were inferred based on standard chemical nomenclature.

• Lewis structure drawing boxes are for student practice and not filled in here.