BCH124 Unit-1 Atomic Structure

Practice flashcards covering the fundamental concepts of atomic structure including historical models, quantum numbers, and electron configuration principles.

Atomic Number ($Z$)

The number of protons found in the nucleus of an atom.

Mass Number ($A$)

The sum of the number of protons ($Z$) and the number of neutrons ($n$) in an atom.

Nucleons

The collective name for protons and neutrons, which are particles of about the same mass found within the nucleus.

Femtometer

A unit of length equal to $10^{-15}\text{ meters}$, reflecting the approximate size of a nucleus.

Law of Conservation of Mass

States that mass can neither be created nor destroyed but is transformed from one form to another.

Law of Constant Composition

States that chemical compounds are composed of elements present in a fixed ratio by mass, such as $H_2O$.

Law of Multiple Proportion

The principle that if two elements form more than one compound, the ratio of the masses of the second element that combine with a fixed mass of the first is a small, whole number.

Isotopes

Atoms of the same element that have varying masses, which was a concept Dalton's atomic theory failed to account for.

Plum Pudding Model

J.J. Thomson's atomic model where negatively charged electrons were envisioned as being contained within a positively charged soup.

Rutherford Gold Foil Experiment

An experiment involving the bombardment of a thin gold foil with positively charged alpha particles, demonstrating that atoms consist mostly of empty space with a central nucleus.

Bohr Model

A model proposed in 1915 where electrons move in fixed orbitals (shells) around the nucleus, each having a fixed energy.

Ground State

The condition of an atom when an electron occupies the lowest available energy level (designated as shell $n=1$).

Zeeman Effect

The effect of a magnetic field on the spectra of atoms, which the Bohr model failed to explain.

Stark Effect

The effect of an electric field on the spectra of atoms, which the Bohr model could not explain.

Heisenberg Uncertainty Principle

The principle stating that it is impossible to precisely measure or calculate both an object's position and its momentum simultaneously.

Quantum-Wave Model

A model describing electron movement as standing waves within probability zones known as atomic orbitals rather than fixed paths.

Atomic Orbitals

Probability areas or regions where there is a likelihood of locating an electron, characterized by different energy levels and sub-levels.

Principal Quantum Number ($n$)

Specifies the principal electron shell and the probable distance of the electron from the nucleus, with values $n=1, 2, 3...$ corresponding to shells $K, L, M, N...$

Azimuthal (Angular Momentum) Quantum Number ($l$)

Defines the shape of an orbital ($s, p, d, f$) and is determined by the formula $n-1$, with values $0, 1, 2, 3$ respectively.

Magnetic Quantum Number ($m_l$)

Specifies the orientation of an orbital in space and can have values ranging from $-l$ to $+l$, totaling $(2l + 1)$ states.

Spin Quantum Number ($m_s$)

Specifies the orientation of electron spin within an orbital, either as spin up ($+1/2$) or spin down ($-1/2$).

Pauli Exclusion Principle

States that no two electrons in an atom can have the same set of all four quantum numbers.

Aufbau Principle

Derived from the German word for 'building up,' it states that electrons occupy atomic orbitals of lower energy before filling higher energy levels.

Hund’s Rule

States that every orbital in a subshell is singly occupied with one electron before any orbital is doubly occupied, and all singly occupied electrons have the same spin.

Octet Rule

The tendency of atoms to prefer having eight electrons in their valence shell ($s^2p^6$ configuration) to achieve greater stability.