Understanding the Quantum Mechanical Model of Atoms

Bohr Model Limitations

The Bohr model of the atom could only explain the behavior of hydrogen atoms and failed to explain the spectra of atoms with more than one electron.

Quantum Mechanical Model

Describes the behavior of electrons in atoms using principles of quantum mechanics, treating electrons as probability clouds.

Heisenberg Uncertainty Principle

States that the position and momentum of an electron cannot both be precisely determined at the same time.

Wave-Particle Duality

Electrons exhibit both wave-like and particle-like properties.

Orbital

A region of space where there is a high probability of finding an electron.

Quantum Numbers

Describes the properties of orbitals and the electrons within them.

Principal Quantum Number (n)

Describes the energy level (shell) and size of the orbital.

Angular Momentum Quantum Number (l)

Describes the shape of the orbital (s, p, d, f).

Magnetic Quantum Number (m_l)

Describes the orientation of the orbital in space.

Spin Quantum Number (m_s)

Describes the direction of electron spin (either +1/2 or -1/2).

Electron Configuration

The arrangement of electrons in an atom, following specific rules.

Aufbau Principle

Electrons fill orbitals starting with the lowest energy level.

Pauli Exclusion Principle

No two electrons can have the same set of quantum numbers.

Hund's Rule

Electrons will fill degenerate orbitals singly before pairing up.

Electron Configuration for Oxygen

Oxygen (O) has 8 electrons: 1s² 2s² 2p⁴.

Ionization Energy

Energy required to remove an electron from an atom, increases across a period and decreases down a group.

Atomic Radius

The size of an atom, decreases across a period and increases down a group.

Electron Affinity

Measures the attraction of an atom for electrons.

Electronegativity

Measures the tendency of an atom to attract a bonding pair of electrons.

Chemical Bonding

Quantum mechanics helps explain how atoms bond to form molecules.

Covalent Bonding

The sharing of electron pairs between atoms, typically between nonmetals.

Ionic Bonding

The transfer of electrons from one atom to another, resulting in oppositely charged ions attracting each other.

Hybridization

The mixing of atomic orbitals to form new hybrid orbitals that can explain the geometry of molecules.

Max Planck

Proposed that energy is quantized, leading to the development of quantum theory.

Albert Einstein

Expanded on the idea of light as quantized particles (photons), contributing to the wave-particle duality.

Niels Bohr

Developed the Bohr model, introducing quantized energy levels for electrons.

Werner Heisenberg

Formulated the uncertainty principle, which states that the position and momentum of an electron cannot both be precisely measured simultaneously.

Erwin Schrödinger

Developed the Schrödinger equation, which describes how the quantum state of a system changes over time.

Emission Spectrum

When electrons drop from a higher energy level to a lower one, they emit photons at specific wavelengths, producing bright lines on a dark background.

Absorption Spectrum

When electrons absorb energy and move to higher levels, they take in photons, leading to dark lines on a continuous spectrum.

Quantized Energy Levels

Atoms have discrete energy levels (or shells) where electrons reside.

Spectral Lines

Each element has a unique set of energy levels, leading to a characteristic set of spectral lines that can be used to identify the element.

Quantum Numbers (set)

A set of four numbers used to describe the energy, shape, orientation, and spin of an electron in an atom.

Azimuthal Quantum Number (l)

Defines the shape of the orbital, with values from 0 to (n-1).

Magnetic Quantum Number (m_l)

Indicates the orientation of the orbital in space, with values from -l to +l.

Spin Quantum Number (m_s)

Describes the spin of the electron, with values +1/2 or -1/2.