Introduction to Inorganic Chemistry: Atomic Structure, Bonding, and Electron Configuration

What is Inorganic Chemistry?

Definition and Scope

• Inorganic chemistry is the branch of chemistry concerned with the properties and reactions of inorganic compounds, which includes all chemical compounds except those based upon chains or rings of carbon atoms (organic compounds).

• It is described as the chemistry of "everything else" outside of organic chemistry, making it an extremely broad field.

• There is significant overlap with organometallic chemistry, which studies compounds containing metal-carbon bonds.

Applications and Interdisciplinary Connections

• Inorganic chemistry is foundational to fields such as medicine (e.g., MRI contrast agents), biochemistry (e.g., metalloproteins), materials science (e.g., semiconductors), geochemistry, and catalysis.

• Examples of important inorganic compounds include hemoglobin, cisplatin (a chemotherapy drug), and vitamin B12.

Organic vs. Inorganic Chemistry

Key Differences

Organic Chemistry	Inorganic Chemistry
Few elements, mostly covalent/polar covalent bonds	All elements, all modes of bonding
Molecular solids (except polymers)	Ionic, extended-network, and molecular solids
Usually air-stable	Varied stability with air/water
Soluble in nonpolar solvents	Wide range of solubilities
Bonding involves s & p electrons	Bonding can involve s, p, d, and f electrons

Bonding in Inorganic Chemistry

Complexity of Bonding

• Bonding in inorganic compounds can be more complex than in organic compounds, often involving multi-center bonds (e.g., in boranes like B2H6).

• Common conceptions of bonding (e.g., two-center, two-electron bonds) are not always sufficient to describe inorganic molecules.

The Scope of Inorganic Chemistry

Major Areas

• Medicine: Imaging agents, drugs

• Biochemistry/Biology: Metalloproteins, catalysis

• Organic Chemistry: Organometallics

• Physics/Materials Science: Semiconductors, superconductors

• Geology/Geochemistry: Mineral synthesis, solar evolution

• Organometallic Chemistry: New compounds, catalysis

Atoms and Quantum Mechanics

Introduction

• Understanding atoms requires quantum mechanics, which describes the wave properties of electrons in atoms.

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

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

Orbitals and Quantum Numbers

Quantum Numbers

• 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 ($l = 0$ to $n-1$; s, p, d, f).

• Magnetic quantum number (ml): Defines orientation of orbital in space ($m_l = -l$ to $+l$).

• Spin quantum number (ms): $+\frac{1}{2}$ or $-\frac{1}{2}$; distinguishes electrons in the same orbital.

Allowed Combinations

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

Shapes of Orbitals

• s-orbitals: Spherical

• p-orbitals: Dumbbell-shaped, three orientations (px, py, pz)

• d-orbitals: Four-lobed shapes, five orientations

• f-orbitals: Eight lobes, seven orientations

• g-orbitals: 8-12 lobes, higher energy

Pauli Exclusion Principle

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

• Each orbital can hold a maximum of two electrons with opposite spins.

Electron Configuration

Filling Order and Rules

• Orbitals are filled in order of increasing energy (Aufbau principle): $1s < 2s < 2p < 3s < 3p < 4s < 3d < 4p < 5s < 4d < 5p < 6s < 4f < 5d < 6p < 7s < 5f < 6d < 7p$

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

• Electron configurations of ions differ from neutral atoms; electrons are removed from the highest $n$ orbitals first for cations.

Examples

• Oxygen: O = [He] 2s2 2p4

• Silicon: Si = [Ne] 3s2 3p2

• Molybdenum: Mo = [Kr] 5s1 4d5 (half-filled d-subshell stability)

• Tungsten: W = [Xe] 6s2 4f14 5d4

Periodic Table Blocks

• s-block: Groups 1-2

• p-block: Groups 13-18

• d-block: Transition metals (Groups 3-12)

• f-block: Lanthanides and actinides

Exceptions and Special Cases

• Half-filled and fully-filled d or f subshells confer extra stability (e.g., Cr, Mo, Cu, Ag, Au).

• Alternative arrangements of the f-block can reduce the number of exceptions in electron configurations.

Electronegativity (EN)

Definition and Trends

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

• Pauling scale is most common; EN increases across a period and decreases down a group.

• EN is inversely related to metallic character.

• Dipole moments are determined by differences in EN ($\Delta EN$).

Special Considerations

• Hydrogen's EN is intermediate (2.1), making its placement in the periodic table ambiguous.

• Main group (MG) elements show a "plane" or "dip" in EN trends, especially in Groups 13 and 14.

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

Applications

• EN is used to evaluate the ionic or covalent character of bonds.

• Many metals form polar covalent bonds, not strictly ionic.

Summary Table: Quantum Numbers and Orbitals

Quantum Number	Symbol	Meaning	Possible Values
Principal	n	Energy level, shell	1, 2, 3, ...
Angular Momentum	l	Shape, subshell	0 to n-1
Magnetic	ml	Orientation	-l to +l
Spin	ms	Spin direction	+1/2, -1/2

Additional info: This guide covers foundational concepts in atomic structure, quantum mechanics, and periodic trends, which are essential for further study in inorganic chemistry.