Quantum Mechanics: Electron Transitions and Numbers

Definition: Movement of electrons between energy levels in an atom.
Types:
- Absorption: Electron gains energy and moves to a higher energy level.
- Emission: Electron loses energy and falls to a lower energy level, often releasing a photon.
Energy Levels: Quantized; electrons can only occupy specific energy levels.
Significance: Electron transitions are responsible for spectral lines in atomic emission/absorption spectra.

Purpose: Describe the unique quantum state of an electron in an atom.
Types:
1. Principal Quantum Number (n): Indicates the energy level and size of the orbital; n = 1, 2, 3, ...
2. Angular Momentum Quantum Number (l): Defines the shape of the orbital; l = 0 to (n-1).
3. Magnetic Quantum Number (m_l): Specifies orientation of the orbital; m_l = -l to +l.
4. Spin Quantum Number (m_s): Describes the spin of the electron; m_s = +1/2 or -1/2.
Principle: No two electrons in an atom can have the same set of quantum numbers (Pauli Exclusion Principle).

Types of Orbitals:
- s-Orbital: Spherical shape, l = 0.
- p-Orbitals: Dumbbell shape, l = 1; three orientations (px, py, pz).
- d-Orbitals: More complex shapes, l = 2; five orientations.
- f-Orbitals: Even more complex, l = 3; seven orientations.
Significance: Orbital shapes affect chemical bonding and properties of elements.

Definition: Arrangement of electrons when one or more electrons are promoted to higher energy levels than the ground state.
Characteristics:
- Higher energy than ground state configurations.
- Often unstable; electrons tend to return to ground state, releasing energy (often as light).
Example: Ground state of hydrogen: 1s¹; excited state can be 2s¹ or 2p¹.

Trends:
- Atomic Radius: Decreases across a period (more protons pull electrons closer); increases down a group (more energy levels).
- Ionization Energy: Energy required to remove an electron; increases across a period (greater nuclear charge); decreases down a group (increased shielding).
- Electronegativity: Ability of an atom to attract electrons; increases across a period; decreases down a group.
- Electron Affinity: Energy change when an electron is added; generally increases across a period.
Group Characteristics: Elements in the same group have similar chemical properties due to similar valence electron configurations.

Movement of electrons between defined energy levels in atoms.
Absorption: An electron gains energy, resulting in an increase to a higher energy level.
Emission: An electron loses energy and transitions to a lower energy level, usually releasing a photon in the process.
Energy levels are quantized, meaning electrons are restricted to specific levels without existing in between.
Electron transitions account for the presence of spectral lines in atomic emission and absorption spectra.

Quantum numbers define the unique state of an electron within an atom.
Principal Quantum Number (n): Indicates both the energy level and size of the electron orbital; values begin at 1 and increase whole numbers (1, 2, 3,...).
Angular Momentum Quantum Number (l): Determines the shape of an orbital; values range from 0 to n-1.
Magnetic Quantum Number (m_l): Describes the orientation of the orbital within a given energy level; values span from -l to +l.
Spin Quantum Number (m_s): Indicates the electron's spin; can take values of +1/2 or -1/2.
According to the Pauli Exclusion Principle, no two electrons can share the same set of quantum numbers.

s-Orbital: Characterized by a spherical shape with l = 0.
p-Orbitals: Exhibit a dumbbell shape with l = 1 and exist in three orientations (px, py, pz).
d-Orbitals: Possess more intricate shapes with l = 2 and contain five different orientations.
f-Orbitals: Have the most complex shapes with l = 3, comprising seven orientations.
The shapes of orbitals significantly influence the chemical bonding and properties of elements.

Refers to the arrangement of electrons when one or more electrons occupy energy levels higher than the ground state.
These configurations are characterized by increased energy compared to ground state configurations.
Excited states are typically unstable, prompting a return to the ground state while releasing energy, often in the form of light.
For instance, the ground state electron configuration of hydrogen is 1s¹, while excited states include configurations like 2s¹ or 2p¹.

Atomic Radius: Decreases across a period due to increased nuclear charge pulling electrons closer; increases down a group because of added energy levels.
Ionization Energy: The energy necessary to remove an electron; it rises across a period due to a stronger nuclear charge and falls down a group due to enhanced shielding effects.
Electronegativity: Reflects an atom’s ability to attract electrons; rises across a period and drops down a group.
Electron Affinity: Represents the energy change accompanying the addition of an electron; often increases across a period.
Elements in the same group exhibit comparable chemical properties due to their analogous valence electron configurations.