Atomic Structure, Electron Configuration, and Chemical Bonding: General Chemistry Study Notes

Atomic Structure and Quantum Numbers

Quantum Numbers and Orbitals

Quantum numbers describe the properties of atomic orbitals and the electrons in them. Each electron in an atom is defined by four quantum numbers:

• Principal quantum number (n): Indicates the energy level and size of the orbital.

• Angular momentum quantum number (l): Defines the shape of the orbital (s, p, d, f).

• Magnetic quantum number (ml): Specifies the orientation of the orbital.

• Spin quantum number (ms): Indicates the spin direction of the electron (+1/2 or -1/2).

Example table of quantum numbers for selected orbitals:

Orbital	n	l	ml	ms
2p	2	1	-1, 0, +1	-1/2, +1/2
3s	3	0	0	-1/2, +1/2
4d	4	2	-2, -1, 0, 1, 2	-1/2, +1/2

Electron Capacity of Orbitals

• Maximum electrons in 3d orbitals: 10 electrons (5 orbitals × 2 electrons each)

• Quantum number combinations: For n = 1, l = 0: 2 electrons (1s orbital)

• Principal quantum number of 2: 8 electrons (2s and 2p orbitals)

Electron Configuration

Writing Electron Configurations

Electron configuration describes the arrangement of electrons in an atom's orbitals. The Aufbau principle, Pauli exclusion principle, and Hund's rule guide the filling order.

• Silicon: $1s^2 2s^2 2p^6 3s^2 3p^2$

• Condensed configuration for Cu2+: Cu (neutral): $[Ar] 4s^2 3d^9$; Cu2+: $[Ar] 3d^9$ (electrons are lost from the 4s before the 3d)

• Cation isoelectronic to Xe: Examples: Cs+, Ba2+, La3+, Ce4+, Pr4+

Periodic Trends in Electron Configuration

• Elements with configurations ending in ns2np6 are found in Group 18 (noble gases).

• Elements with configurations ending in 3s23p6 are also noble gases (e.g., Argon).

Periodic Trends: Atomic Radius and Electron Affinity

Atomic Radius

Atomic radius is influenced by the number of electron shells and effective nuclear charge ($Z_{eff}$).

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

• Example: $K^+$, Ca, K (smallest to largest): $K^+$ < Ca < K

• Explanation: More shells result in a larger radius; higher $Z_{eff}$ pulls electrons closer, reducing radius.

Electron Affinity

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

• Trend: Most negative electron affinity is for elements that strongly attract electrons (e.g., F).

• Order: F < O < Ne (most negative to most positive)

• Explanation: Fluorine has the highest electron affinity due to its high $Z_{eff}$ and nearly full outer shell.

Ionization Energy

Successive Ionization Energies

Ionization energy is the energy required to remove an electron from an atom.

• Trend: Each successive ionization requires more energy, especially after removing all valence electrons.

• Example: The jump in ionization energy after the removal of valence electrons indicates a new shell is being accessed.

• Charge after fifth ionization: $5^+$ (each ionization removes one electron)

Chemical Bonding and Molecular Geometry

Bond Order and Resonance

Bond order indicates the number of chemical bonds between a pair of atoms. Resonance structures affect bond order.

• Bond order formula: $\text{Bond order} = \frac{\text{sum of individual bond orders in all resonance structures}}{\text{number of resonance structures}}$

• SO3: All S–O bonds are double bonds; bond order = 2.

• SO42−: Resonance gives bond order = 1.5.

Electronic and Molecular Geometry

Geometry is determined by the arrangement of electron pairs around the central atom.

Species	Electronic Geometry	Molecular Geometry	Ideal Bond Angle
SO3	Trigonal planar	Trigonal planar	120°
SO32−	Tetrahedral	Trigonal pyramidal	109.5°
SO42−	Tetrahedral	Tetrahedral	109.5°

Polarity and Lewis Structures

Polarity of Molecules

Polarity depends on the presence of polar bonds and molecular geometry.

Molecule	Polar Bonds?	Polar Molecule?
SO3	Yes	No
S8	No	No
H2S	Yes	Yes

Lewis Structures and Formal Charge

Lewis structures show the arrangement of electrons in a molecule. Formal charge helps determine the most stable structure.

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

• SF2O: 26 valence electrons; S can expand its octet.

• PO3−: 24 valence electrons; P can expand its octet.

• H2N2: 12 valence electrons.

Resonance Structures

Properties of Resonance

• Resonance structures are different possible arrangements of electrons in a molecule.

• Actual structure is a hybrid of all resonance forms.

• Some molecules can shift between resonance structures.

Electronegativity and Lattice Energy

Electronegativity

Electronegativity is the tendency of an atom to attract electrons in a bond. It increases across a period and decreases down a group.

Element	Electronegativity
H	2.1
Li	1.0
Be	1.5
B	2.0
C	2.5
N	3.0
O	3.5
F	4.0

Lattice Energy

Lattice energy is the energy released when ions form a solid lattice.

• Lattice energy formula: $\text{Lattice energy} = \frac{kQ_1Q_2}{d_{radius}}$

Additional info: These notes expand on the original exam questions by providing definitions, formulas, and context for each topic, making them suitable for comprehensive exam review.