Atomic Structure, Electron Configuration, and Atomic Orbitals: Study Notes

Atomic Models and Structure

Historical Development of Atomic Models

The understanding of atomic structure has evolved through several key models, each contributing to our knowledge of the atom's composition and behavior.

• Dalton Model (1803): Atoms are small, indivisible spheres.

• Thomson Model (1897): "Plum Pudding" model; electrons are suspended in a positive gel.

• Rutherford Model (1911): Atoms are mostly empty space with a very small, dense, positively charged nucleus; electrons move around the nucleus.

• Bohr Model (1913): Electrons travel in concentric rings (orbits) around the nucleus, each with a specific energy level.

Example: The Bohr model explains why elements emit light at specific wavelengths (atomic spectra).

Bohr Model and Energy Levels

The Bohr model introduced the concept of quantized energy levels for electrons in atoms.

• Electrons occupy specific energy levels (orbits) around the nucleus.

• Each orbit corresponds to a fixed amount of energy ($E$).

• Electrons can jump to a higher energy level when energy is absorbed, and fall to a lower energy level when energy is emitted.

• The energy difference between levels is called a quantum of energy.

• Electrons emit electromagnetic radiation (light) when they lose energy.

Formula:

$E = h\nu$

where $E$ is energy, $h$ is Planck's constant, and $\nu$ is frequency.

Additional info: The Bohr model is most accurate for hydrogen-like atoms and does not fully explain the behavior of multi-electron atoms.

Quantum Mechanical Model and Atomic Orbitals

Quantum Mechanical Model

The quantum mechanical model describes electrons as existing in probable regions called orbitals, rather than fixed paths.

• Electrons are restricted to certain energy levels ($n$).

• Each energy level contains one or more sublevels (s, p, d, f).

• Orbitals are regions where electrons are likely to be found; they have different shapes and orientations.

• There is no exact path for electron movement; only probability clouds.

Example: The 1s orbital is spherical, while the 2p orbitals are dumbbell-shaped.

Principal Quantum Number ($n$)

The principal quantum number ($n$) indicates the primary energy level of an electron.

• $n$ can be any integer from 1 to 7 for known elements.

• As $n$ increases, the electron's average distance from the nucleus increases.

• Energy levels get closer together as $n$ increases.

Formula:

$n = 1, 2, 3, \ldots$

Atomic Orbitals and Sublevels

Each energy level contains sublevels, which are associated with different orbital shapes.

• s orbital: Spherical shape

• p orbital: Dumbbell (peanut) shape

• d orbital: Double peanut shape

• f orbital: Flower shape

Each sublevel contains a specific number of orbitals:

Sublevel	Number of Orbitals
s	1
p	3
d	5
f	7

Additional info: Each orbital can hold up to 2 electrons.

Energy Level and Sublevel Table

The following table summarizes the number of sublevels, orbitals, and electrons for the first four energy levels:

n	# of Sublevels	Sublevels	# of Orbitals	# of Electrons
1	1	s	1	2
2	2	s, p	4	8
3	3	s, p, d	9	18
4	4	s, p, d, f	16	32

Additional info: Orbitals within a specific subshell (e.g., the 3 orbitals in the 3p sublevel) are equal in energy.

Electron Configuration

Electron Configuration Notation

Electron configuration describes the arrangement of electrons in an atom's orbitals.

• Electrons fill orbitals in order of increasing energy (Aufbau principle).

• Each orbital can hold a maximum of 2 electrons (Pauli exclusion principle).

• Electrons occupy orbitals singly before pairing (Hund's rule).

Example:

• Hydrogen (H): 1s1

• Helium (He): 1s2

• Carbon (C): 1s2 2s2 2p2

• Neon (Ne): 1s2 2s2 2p6

Orbital Diagrams

Orbital diagrams visually represent electron configurations using boxes and arrows.

• Each box represents an orbital; arrows indicate electrons and their spins.

• Electrons fill lower energy orbitals first, then higher ones.

Additional info: The order of filling is: 1s, 2s, 2p, 3s, 3p, 4s, 3d, 4p, 5s, 4d, 5p, 6s, 4f, 5d, 6p, 7s, 5f, 6d, 7p.

Noble Gas Shortcut

The noble gas shortcut simplifies electron configuration notation by using the previous noble gas as a reference point.

• Write the symbol of the previous noble gas in brackets, then continue with the configuration.

Example: Silver (Ag): [Kr] 5s1 4d10

Periodic Table and Electron Configuration

The periodic table is organized according to electron configurations, with blocks corresponding to s, p, d, and f sublevels.

• s-block: Groups 1-2

• p-block: Groups 13-18

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

• f-block: Lanthanides and actinides

Additional info: The position of an element in the periodic table can be used to determine its electron configuration.

Summary Table: Electron Configuration Examples

Element	Electron Configuration	Noble Gas Shortcut
Si (#14)	1s2 2s2 2p6 3s2 3p2	[Ne] 3s2 3p2
Ag (#47)	1s2 2s2 2p6 3s2 3p6 4s2 3d10 4p6 5s1 4d10	[Kr] 5s1 4d10
W (#74)	1s2 2s2 2p6 3s2 3p6 4s2 3d10 4p6 5s2 4d10 5p6 6s2 4f14 5d4 6p6 7s2 5f14 6d4	[Rn] 7s2 5f14 6d4

Key Terms and Concepts

• Quantum: The smallest discrete amount of energy that can be absorbed or emitted.

• Orbital: A region in space where there is a high probability of finding an electron.

• Sublevel: A set of orbitals within an energy level (s, p, d, f).

• Electron Configuration: The arrangement of electrons in an atom's orbitals.

• Noble Gas Shortcut: A method to abbreviate electron configurations using noble gases.