Structure of Atom - Lecture Notes Flashcards

Flashcards covering the fundamental concepts of atomic structure, historical models, electromagnetic radiation, quantum mechanics, and orbital filling rules.

Atom

Derived from the Greek word ‘a-tomio’, meaning ‘un-cutable’ or ‘non-divisible’.

Thomson's Model of Atom

Proposed in 1898, suggesting an atom is a sphere of radius approximately $10^{-10}\,m$ with uniformly distributed positive charge and embedded electrons; also known as the plum pudding, raisin pudding, or watermelon model.

Nucleus

A very small portion of the atom where positive charge and most mass are densely concentrated; discovered by Rutherford with a radius of about $10^{-15}\,m$.

Orbits

Circular paths where electrons move around the nucleus at very high speeds, similar to planets revolving around the sun.

Rutherford Model Stability Issue

According to Maxwell's electromagnetic theory, an accelerating electron should emit radiation and spiral into the nucleus in $10^{-8}\,s$, which does not happen.

Atomic Number (Z)

Equal to the number of protons in the nucleus of an atom or the number of electrons in a neutral atom.

Nucleons

The collective term for protons and neutrons present in the nucleus.

Mass Number (A)

The total number of nucleons (protons and neutrons) in an atom, calculated as $A = Z + n$.

Isobars

Atoms of different elements with different atomic numbers but the same mass number, resulting in different electronic configurations.

Isotopes

Atoms with the same atomic number and position in the periodic table, but different nucleon numbers due to different numbers of neutrons.

Wavelength ($\lambda$)

The physical distance between two consecutive crests or two consecutive troughs.

Frequency ($\nu$)

The number of wave cycles that pass a fixed point per second; measured in Hertz ($Hz$).

Wavenumber

The reciprocal of the wavelength ($\lambda$), indicating the number of waves in a specific length (usually $1\,cm$).

Velocity of Light ($c$)

The speed of light in a vacuum, approximately $2.99792458 \times 10^8\,m/s$, which equals the product of frequency and wavelength ($c = \lambda\nu$).

Planck's Quantum Theory

Proposed that energy is quantized and light has particle character, expressed as $E = h\nu$, where $h = 6.626 \times 10^{-34}\,J \cdot s$.

Photoelectric Effect

The emission of electrons from a metal surface when light of a suitable frequency hits it; first observed by H. Hertz in 1887.

Photons

Described by Einstein as a stream of 'energy packets', where each packet of frequency $\nu$ carries energy $E = h\nu$.

Work Function ($\Phi$)

The minimum energy required to free an electron from a metal surface, represented as $h\nu_0$.

Threshold Frequency ($\nu_0$)

The characteristic minimum frequency for a metal; below this value, no electrons are ejected regardless of light intensity.

Atomic Spectra

The particular pattern of discrete wavelengths absorbed and emitted by an element, used to identify atoms.

Lyman Series

A series of lines in the hydrogen spectrum observed in the UV region when electrons fall to the 1st energy level.

Balmer Series

A series of lines in the hydrogen spectrum observed in the visible region when electrons fall to the 2nd energy level.

Rydberg Equation

An empirical equation relating lines in the hydrogen spectrum, where the Rydberg constant ($R$) is $1.097 \times 10^7\,m^{-1}$.

Ground State

The lowest energy level or orbital where electrons are initially present.

Heisenberg Uncertainty Principle

Proposed in 1927, stating it is impossible to determine simultaneously both the position and momentum of an electron: $\Delta p \Delta x \ge \frac{h}{4\pi}$.

Orbital

A region of space where the probability of finding an electron ($\psi^2$) is maximum; it represents a delocalized wave in the quantum mechanical model.

Principal Quantum Number ($n$)

Specifies the location, energy, and effective volume of the electron cloud, taking values $1, 2, 3, \dots$.

Azimuthal Quantum Number ($l$)

Also called the angular momentum quantum number, it determines the shape of the orbital, with values from $0$ to $n - 1$.

Magnetic Quantum Number ($m$)

Determines the magnetic orientation of an orbital relative to a magnetic field, with values from $-l$ to $+l$ including zero.

Spin Quantum Number

Indicates the direction of electron revolution with two permitted values: $+1/2$ or $-1/2$.

Aufbau Principle

Electrons fill various orbitals in increasing order of their energies; the orbital with the lowest energy is filled first.

Pauli’s Exclusion Principle

States that no two electrons in an atom can have all four quantum numbers the same; an orbital can hold a maximum of two electrons with opposite spins.

Hund’s Rule of Maximum Multiplicity

No electron pairing occurs in $p$, $d$, or $f$ orbitals until each orbital in a sub-shell contains one electron.

Radial Node

A region in an orbital where the probability of finding an electron is zero, calculated as $n - l - 1$.

Exchange Energy

Energy released when electrons with the same spin in degenerate orbitals 'exchange' positions, stabilizing half-filled and fully-filled orbitals.