Atomic Structure and Subatomic Particles

Atomic Structure

• Atoms are the fundamental building blocks of matter. They consist of a nucleus containing protons and neutrons, surrounded by orbiting electrons.

• The nucleus is incredibly dense, containing most of the atom's mass. Electrons are much lighter and occupy a significantly larger volume around the nucleus.

• Protons have a positive charge, electrons have a negative charge, and neutrons have no charge (are neutral). The number of protons defines the element.

• Atomic number (Z) = number of protons in the nucleus. This determines the element. For a neutral atom, the number of protons equals the number of electrons.

• Mass number (A) = number of protons + number of neutrons. This differs between isotopes of the same element, which have the same number of protons but varying numbers of neutrons.

Subatomic Particles

• Protons: Positively charged particles located within the atomic nucleus. They are relatively massive compared to electrons.

• Neutrons: Neutral particles found in the atomic nucleus. Their mass is very similar to that of protons.

• Electrons: Negatively charged particles orbiting the nucleus in specific energy levels or electron shells. They have a significantly smaller mass than protons or neutrons.

Atomic Models

• Historical models of the atom, like the plum pudding model and the Bohr model, have evolved over time. They have increasingly accurately described the atom's structure.

• The modern model describes electrons as existing in orbitals rather than precise orbits. These orbitals define regions of high probability for finding an electron, and electrons can move between them.

Isotopes

• Isotopes are atoms of the same element with differing numbers of neutrons.

• Isotopes have the same atomic number (Z) but different mass numbers (A).

• Isotopes have nearly identical chemical properties, as the chemical behavior is primarily determined by the number of electrons.

• The existence of isotopes is important for radiometric dating and nuclear chemistry. The different isotopes of an element can have different stabilities and radioactive decay rates.

Atomic Mass and Average Atomic Mass

• Atomic mass refers to the mass of a single atom, typically given in atomic mass units (amu).

• The periodic table lists average atomic mass, which is a weighted average of the masses of all the naturally occurring isotopes of an element, considering their relative abundance.

• The differences between atomic and average atomic mass reflect the variations in isotopic composition in natural samples.

Electron Configuration and Energy Levels

• Electrons occupy specific energy levels or shells around the nucleus. These are often represented by principal quantum numbers (n).

• The innermost energy levels hold fewer electrons than the outermost levels.

• Electrons within each energy level occupy specific sublevels (s, p, d, f). These sublevels have specific shapes and orientations in space, which determine electron behaviour.

• The arrangement of electrons in an atom's energy levels, sublevels, and orbitals is the electron configuration.

Quantum Numbers

• Quantum numbers are a set of four numbers describing the properties of an electron.

• Principle quantum number (n): describes the energy level of the electron.

• Azimuthal quantum number (l): describes the shape of the electron's orbital.

• Magnetic quantum number (ml): describes the orientation of the electron's orbital in space.

• Spin quantum number (ms): describes the intrinsic angular momentum or "spin" of the electron (either +1/2 or -1/2).

Periodic Trends

• The periodic table arranges elements based on atomic structure and chemical properties.

• Trends like electronegativity, ionization energy, and atomic radius correlate with the arrangement of energy levels and the number of protons and electrons in the atoms.

• These trends allow understanding of how elements will interact and react in different chemical processes.