Atomic Structure

Electromagnetic Radiation

The entire range of wavelengths or frequencies of electromagnetic waves, including gamma rays, X-rays, ultraviolet radiation, infrared radiation, microwaves, and radio waves.

Gamma Rays

High-energy electromagnetic radiation emitted from the atomic nucleus, originating from natural and artificial sources such as radioactive materials and nuclear explosions. They have high penetration power, capable of passing through three meters of concrete.

X-rays

High-energy electromagnetic radiation that can easily pass through materials, including soft tissues of animals. They are used in medicine for imaging and therapy purposes, but excessive exposure can lead to cell damage and mutations.

Ultraviolet Radiation

High-energy light waves that can ionize atoms in the Earth's atmosphere, creating the ionosphere. While small doses can promote vitamin D production and minimal tanning, excessive exposure can cause sunburn, retinal damage, and skin cancer.

Infrared Radiation

Weak heat radiation emitted by objects above absolute zero temperature. Is used in various applications such as night vision scopes, electronic detectors, sensors, and household heating systems.

Microwaves

Highest frequency radio waves emitted by large objects like the Earth, buildings, and airplanes. They are also present in space and used for RADAR, satellite communication, and astronomical studies.

Radio Waves

Low-energy waves that transfer energy in a continuous fashion. They are used in radio and television transmission, magnetic resonance imaging (MRI), and can be detected by radiotelescopes.

Visible Light

The small portion of the electromagnetic spectrum that is visible to the human eye, ranging from deep red to deep violet. It represents only about 2.5% of the entire spectrum.

Spectroscopy

The study of electromagnetic radiation absorbed or emitted by atoms or molecules. It includes emission spectroscopy, which analyzes the wavelengths of light emitted by excited atoms, and absorption spectroscopy, which identifies elements by analyzing the wavelengths of light absorbed by a sample.

Emission Spectroscopy

A technique where light of various wavelengths is emitted by excited atoms, producing an emission spectrum. It is commonly used to identify elements and study their energy transitions.

Absorption Spectroscopy

The study of electromagnetic radiation absorbed by atoms or molecules, resulting in dark lines within a continuous spectrum. It is used to identify elements in stars and determine the concentration of chemical compounds in samples.

Fireworks

The process where atoms or molecules in fireworks are vaporized and excited, emitting light of characteristic wavelengths. The emitted light produces the vibrant colors seen in fireworks displays.

Dalton's Atomic Theory

A theory stating that all elements are composed of tiny particles called atoms, atoms of the same element are identical, atoms of different elements can be mixed or combined in simple ratios to form compounds, and atoms are neither created nor destroyed in chemical reactions.

Thomson's Plum Pudding Model

A model proposed by—-, suggesting that atoms consist of negatively charged electrons embedded in a positively charged material, similar to raisins in a dough.

Rutherford's Nuclear Model

A model proposed by —-, stating that atoms are mostly empty space, with a positive charge concentrated in the nucleus and electrons orbiting around it. It explained the results of the gold foil experiment and the deflection of alpha particles.

Bohr's Planetary Model

A model proposed by ——, where electrons are arranged in fixed energy levels or orbits around the nucleus. Electrons absorb and emit energy in the form of light when transitioning between energy levels, resulting in line spectra.

Quantum Theory

A theory developed by Max Planck and applied by Bohr, stating that energy is quantized and can only occur in discrete packets or quanta. It explained the stability of electrons in specific energy levels and the emission of characteristic frequencies of light.

Wave-Particle Duality

The concept proposed by Louis de Broglie, suggesting that particles like electrons can exhibit wave-like behavior. It challenged the idea of well-defined electron orbits and led to the development of quantum mechanics.

Quantum Mechanical Model

A model developed by Erwin Schrödinger, describing the location and energy of electrons in terms of probability distributions called orbitals. It uses wave functions and complex differential calculus to estimate the statistical likelihood of finding an electron in a particular region of an atom.

Orbitals

Regions in space where electrons with a specific energy level are most likely to be found.

Principal quantum number (n)

Indicates the energy level in which an electron is located. The higher the value, the further the orbital is from the nucleus and the higher its energy.

Quantum number (l)

Identifies the sublevel within an energy level where an electron can be found.

Quantum number (m)

Represents the specific orbital within the sublevel where the electron is located.

Electronegativity

The tendency of an atom to attract electrons towards itself. Increases from left to right across a period and decreases as you go down a group for representative elements.

Atomic Radius

Half the distance between adjacent nuclei in a crystal for metals, and half the distance between chemically bonded identical atoms for non-metals. Decreases from left to right across the periodic table and increases down a group.

Ionic Radii

The size of both positive and negative ions. Decreases across a period and increases down a group due to higher principal energy levels.

Ionization Energy

The energy required to remove an electron from an atom. High electronegativity leads to high ionization energy. Decreases down a group and increases from left to right across the periodic table.