CHEM 30301 Inorganic Chemistry 1: Atomic Structure, Quantum Mechanics, and Electron Configuration

Introduction to Inorganic Chemistry

Definition and Scope

Inorganic chemistry is the branch of chemistry concerned with the properties and reactions of inorganic compounds. It encompasses all chemical compounds except those based primarily on chains or rings of carbon atoms (organic compounds). The distinction between organic and inorganic chemistry is not absolute, with significant overlap in organometallic chemistry.

• Organic Chemistry: Focuses on hydrocarbons and their derivatives.

• Inorganic Chemistry: Describes the chemistry of "everything else," including all elements and all modes of bonding.

• Applications: Medicine (MRI, X-ray imaging), Biochemistry/Biology (metalloenzymes, O2 binding), Physics/Materials Science (semiconductors, superconductors), Geology/Geochemistry (mineral synthesis), and Organometallic Chemistry (new compounds, catalysis).

Comparison: Organic vs. Inorganic Chemistry

Organic Chemistry	Inorganic Chemistry
Few elements, mostly covalent/polar bonds	All elements, all bonding modes
Molecular solids (except polymers)	Ionic, extended-network, molecular solids
Air-stable, soluble in nonpolar solvents	Varied stability and solubility
Distillable, crystallizable	Wide solubility range
Bonding involves s & p electrons	Bonding involves s, p, d, f electrons

Bonding in Inorganic Chemistry

Bonding Concepts

Bonding in inorganic chemistry is more diverse than in organic chemistry, involving ionic, covalent, metallic, and extended-network structures. Common conceptions of bonding are often insufficient, especially for complex molecules like boranes (e.g., B2H6).

• Example: B2H6 features unusual bonding where hydrogen forms more than one bond, and the B-H bond length differs from typical values.

• Bonding Types: Covalent, ionic, metallic, and network solids.

Atoms and Quantum Mechanics

Quantum Mechanics in Chemistry

Understanding atoms requires quantum mechanics, which describes the wave properties of electrons in atoms. The Schrödinger Equation is central to this understanding:

• Schrödinger Equation: $\hat{H}\Psi = E\Psi$

• Energy Levels: $E_n = -\frac{hcRZ^2}{n^2}$, where $n = 1, 2, 3, ...$ and $hcR \approx 13.6$ eV

• Wavefunction ($\Psi$): Describes where electrons exist and their quantum numbers.

Orbitals and Quantum Numbers

Quantum Numbers

Quantum mechanics provides three main quantum numbers that define atomic orbitals:

• Principal Quantum Number ($n$): Defines energy level and size of orbital ($n = 1, 2, 3, ...$).

• Angular Momentum Quantum Number ($l$): Defines shape/type of orbital (subshell: s, p, d, f; $l = 0$ to $n-1$).

• Magnetic Quantum Number ($m_l$): Defines orientation of orbital in space ($m_l = -l$ to $+l$).

l	Designation	Shape (3D)
0	s	Sphere
1	p	Dumbbell
2	d	Four-lobed
3	f	Eight-lobed

Electron Spin and Pauli Exclusion Principle

• Electron Spin Quantum Number ($m_s$): $+\frac{1}{2}$ or $-\frac{1}{2}$

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

Electron Configuration

Rules and Principles

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

• Aufbau Principle: Orbitals are filled in order of increasing energy.

• Relative Energies: $1s < 2s < 2p < 3s < 3p < 4s < 3d < 4p < 5s < 4d < 5p < 6s < 4f < 5d < 6p < 7s < 5f < 6d < 7p$

Electron Configurations of Atoms and Ions

• Neutral Atoms: Fill orbitals according to the above order.

• Ions: For cations, remove electrons from the highest $n$ orbital first.

• Valence Electrons: For main group, highest $n$; for transition metals, $(n-1)d$ and $ns$.

Periodic Table Blocks

Block	Orbitals
s-block	ns
p-block	np
d-block	(n-1)d
f-block	(n-2)f

Exceptions and Special Cases

• Half-filled and fully-filled d or f subshells: Special stability, leading to exceptions in electron configurations (e.g., Mo, W).

• Group 10: Ni follows expected rules, Pd and Pt are exceptions.

• f-block elements: Arrangement affects number of exceptions; alternative arrangements can reduce exceptions.

Electronegativity (EN)

Definition and Trends

• Electronegativity: Tendency of an atom to attract electrons in a chemical bond. Pauling scale is most common.

• EN Trends: Increases across a period, decreases down a group. Most EN: nonmetals; least EN: metals.

• EN and Bonding: Difference in EN ($\Delta EN$) determines bond polarity and dipole moments.

• Metallic Character: Inversely related to EN.

Special Considerations for Hydrogen

• Hydrogen's EN is close to the middle of the scale (2.1).

• Hydrogen is not part of Groups 1 or 17 due to its EN and bonding behavior.

• Hydrogen can form both H+ and H-, but is not a metalloid.

EN Trends in the Periodic Table

• Main group (MG) elements show a "plane" or "dip" in EN values.

• Transition metals (Mt) have high EN and can form covalent or polar covalent bonds with nonmetals.

• EN is used to evaluate the ionic component of bonds; many metals make covalent bonds based on EN.

Summary Table: Quantum Numbers and Orbitals

n	l	ml	# of orbitals	Type of orbitals
1	0	0	1	1s
2	0	0	1	2s
2	1	-1, 0, +1	3	2p
3	0	0	1	3s
3	1	-1, 0, +1	3	3p
3	2	-2, -1, 0, +1, +2	5	3d

Additional info: These notes expand on the provided slides with definitions, examples, and context for quantum mechanics, atomic structure, and periodic trends, suitable for college-level inorganic chemistry students.