Electrons in Atoms and the Periodic Table

Electromagnetic Radiation and Its Properties

Electromagnetic radiation is a form of energy that exhibits wave-like behavior as it travels through space. It includes a wide range of wavelengths and frequencies, from radio waves to cosmic rays.

• Electromagnetic Spectrum: The range of all types of electromagnetic radiation, ordered by wavelength or frequency. It includes (from longest to shortest wavelength): radio waves, microwaves, infrared, visible light, ultraviolet, X-rays, and gamma rays.

• Frequency (\( \nu \)): The number of wave cycles that pass a given point per second. Measured in hertz (Hz).

• Wavelength (\( \lambda \)): The distance between successive crests of a wave. Measured in meters (m), nanometers (nm), etc.

• Relationship between Energy, Frequency, and Wavelength: The energy of light is directly proportional to its frequency and inversely proportional to its wavelength.

Key Equations:

• Relationship between frequency and wavelength: where \( c \) is the speed of light (\( 3.00 \times 10^8 \) m/s).

• Energy of a photon: where \( h \) is Planck's constant (\( 6.626 \times 10^{-34} \) J·s).

• Energy and wavelength:

Visible Light: The portion of the electromagnetic spectrum visible to the human eye, ranging from approximately 400 nm (violet) to 700 nm (red). The order of colors is red, orange, yellow, green, blue, indigo, violet (ROY G BIV).

Atomic Spectra and Energy Levels

Atoms emit or absorb light at specific wavelengths, producing atomic spectra. These spectra are direct evidence of quantized energy levels within atoms.

• Bohr Model: Electrons orbit the nucleus in fixed energy levels (shells). Electrons can move between levels by absorbing or emitting energy equal to the difference between the levels.

• Emission Spectrum: Produced when electrons fall from higher to lower energy levels, emitting photons of specific energies (colors).

• Absorption Spectrum: Produced when electrons absorb energy and move to higher energy levels, resulting in dark lines at specific wavelengths in the spectrum.

Example: The hydrogen atom emits light at specific wavelengths, producing the characteristic Balmer series in the visible region.

Energy Levels, Sublevels, and Orbitals in Atoms

Electrons in atoms occupy regions of space called orbitals, which are organized into energy levels and sublevels.

• Quantized: Electrons can only exist in specific energy states, not between them.

• Principal Shell (n): The main energy level, designated by the principal quantum number (n = 1, 2, 3, ...).

• Subshell: Subdivisions of principal shells, labeled s, p, d, f, corresponding to different orbital shapes.

• Orbital: A region of space where there is a high probability of finding an electron. Each orbital can hold up to two electrons.

• Electron Spin: A property of electrons; each orbital can hold two electrons with opposite spins.

• Valence Electron: Electrons in the outermost shell, responsible for chemical properties.

Subshells and Orbitals:

Subshells in Principal Shells:

• n = 1: s

• n = 2: s, p

• n = 3: s, p, d

• n = 4: s, p, d, f

Orbital Diagrams and Electron Configurations

Electron configurations describe the arrangement of electrons in an atom. Orbital diagrams use arrows to represent electrons and their spins in orbitals.

• Aufbau Principle: Electrons fill orbitals in order of increasing energy (from lowest to highest).

• Order of Filling: 1s < 2s < 2p < 3s < 3p < 4s < 3d < 4p < 5s < 4d < 5p < 6s

• Hund’s Rule: Electrons occupy orbitals singly with parallel spins before pairing up in orbitals of the same energy.

• Pauli Exclusion Principle: Each orbital can hold a maximum of two electrons with opposite spins.

• Blocks on the Periodic Table: The s, p, d, and f blocks correspond to the type of subshell being filled.

• Electron Configuration Notation: Lists the occupied subshells and the number of electrons in each (e.g., 1s2 2s2 2p6).

• Noble Gas Core Notation: Uses the symbol of the previous noble gas in brackets to abbreviate the configuration (e.g., [Ne] 3s2 3p5).

Example: The electron configuration of calcium (Z = 20) is 1s2 2s2 2p6 3s2 3p6 4s2, or [Ar] 4s2.

Identifying Invalid Configurations: Check for violations of the Aufbau principle, Hund’s rule, or the Pauli exclusion principle.

Core and Valence Electrons

Electrons are classified as core or valence based on their location in the atom.

• Core Electrons: Electrons in inner shells, not involved in chemical bonding.

• Valence Electrons: Electrons in the outermost shell, determine chemical reactivity.

• Determining Number: The number of valence electrons can be found from the electron configuration or the group number on the periodic table.

Example: Chlorine (Z = 17) has 7 valence electrons (3s2 3p5).

Chemical Properties and Valence Electrons

The chemical properties of elements are largely determined by the number of valence electrons. Elements in the same group have similar valence electron configurations and similar chemical behavior.

Periodic Trends

Periodic trends are patterns in properties of elements across periods and groups in the periodic table.

• Atomic Radius: The size of an atom. Decreases across a period (left to right), increases down a group.

• Ionization Energy: The energy required to remove an electron from an atom. Increases across a period, decreases down a group.

• Metallic Character: The tendency to lose electrons and form positive ions. Decreases across a period, increases down a group.

Example: Sodium (Na) has a larger atomic radius and lower ionization energy than chlorine (Cl) in the same period.

Additional info: These notes expand on the learning objectives by providing definitions, examples, and tables for clarity and completeness.