Unit 2: Structure of Atom Lecture Notes

Comprehensive flashcards covering the discovery of sub-atomic particles, historical atomic models, wave-particle duality, quantum mechanics, and electronic configuration based on the chemistry lecture notes.

Atom

Derived from the Greek word 'a-tomio' which means 'uncut-able' or 'non-divisible', these were viewed by early philosophers as the fundamental building blocks of matter.

Dalton’s Atomic Theory

Proposed by John Dalton in 1808, it regarded the atom as the ultimate particle of matter and successfully explained laws of chemical combination but failed to explain electrical charging by rubbing.

Cathode Ray Discharge Tubes

Glass tubes containing two thin metal electrodes used to study electrical discharge through gases at very low pressures and high voltages.

Electrons

Negatively charged sub-atomic particles that constitute cathode rays and are basic constituents of all atoms, discovered through experiments in discharge tubes.

Charge to Mass Ratio of Electron ($e/m_e$)

Determined by J.J. Thomson in 1897 using perpendicular electrical and magnetic fields, given as $1.758820 \times 10^{11} \, C \, kg^{-1}$.

Millikan’s Oil Drop Experiment

A method devised by R.A. Millikan to determine the charge on an electron, which he found to be $-1.6 \times 10^{-19} \, C$ (present value: $-1.602176 \times 10^{-19} \, C$).

Proton

The smallest and lightest positive ion obtained from hydrogen, characterized in 1919 as a fundamental positively charged particle.

Neutron

Electrically neutral sub-atomic particles with a mass slightly greater than that of protons, discovered by James Chadwick in 1932 by bombarding beryllium with $\alpha$-particles.

Thomson Model of Atom

Also known as the plum pudding or watermelon model, it proposed that an atom is a sphere of uniform positive charge with electrons embedded into it.

Rutherford’s Nuclear Model

A model proposing that the positive charge and most of the atom's mass are concentrated in an extremely small region called the nucleus, with electrons moving in circular orbits.

Atomic Number ($Z$)

The number of protons in the nucleus of an atom, which is equal to the number of electrons in a neutral atom.

Mass Number ($A$)

The total number of nucleons (protons and neutrons) present in the nucleus of an atom.

Isobars

Atoms with the same mass number ($A$) but different atomic numbers ($Z$), such as $_{6}^{14}C$ and $_{7}^{14}N$.

Isotopes

Atoms with the same atomic number but different mass numbers due to different numbers of neutrons, such as protium ($_{1}^{1}H$), deuterium ($_{1}^{2}D$), and tritium ($_{1}^{3}T$).

Electromagnetic Radiation

Radiation produced by accelerated charged particles that consists of oscillating electric and magnetic fields transmitted in the form of waves.

Frequency ($\nu$)

The number of waves that pass a given point in one second, with the SI unit hertz ($Hz$ or $s^{-1}$).

Wavelength ($\lambda$)

The distance between consecutive points of the same phase in a wave, related to frequency and the speed of light by $c = \nu \lambda$.

Black Body Radiation

The radiation emitted by an ideal body that absorbs and emits all frequencies of electromagnetic radiation uniformly.

Planck’s Quantum Theory

The theory that atoms and molecules emit or absorb energy only in discrete quantities called quanta, where $E = h \nu$.

Planck’s Constant ($h$)

The proportionality constant in the energy equation $E = h \nu$, with a value of $6.626 \times 10^{-34} \, J \, s$.

Photoelectric Effect

The phenomenon where electrons are ejected from a metal surface when it is exposed to light of at least a characteristic threshold frequency ($\nu_0$).

Line Spectrum

As opposed to a continuous spectrum, this consists of light emitted at specific wavelengths with dark spaces between them, also known as atomic spectra.

Balmer Series

A series of lines in the hydrogen spectrum that appear in the visible region, obeying the wavenumber formula $\bar{\nu} = 109,677 \, (\frac{1}{2^2} - \frac{1}{n^2}) \, cm^{-1}$ where $n \ge 3$.

Rydberg Constant ($R_H$)

A constant used for expressing wavelengths in atomic spectra; for hydrogen, its energy value is $2.18 \times 10^{-18} \, J$ and its wavenumber value is $109,677 \, cm^{-1}$.

Bohr Orbit

The fixed circular path of radius $a_0 = 52.9 \, pm$ where an electron in a hydrogen atom moves in the ground state ($n = 1$).

Heisenberg’s Uncertainty Principle

The principle stating it is impossible to determine simultaneously the exact position and exact momentum of a sub-atomic particle like an electron.

de Broglie Relation

The equation $\lambda = \frac{h}{mv} = \frac{h}{p}$ showing that matter exhibits both particle and wave-like properties.

Quantum Mechanics

A theoretical science that deals with the motion of microscopic objects that have both observable wave-like and particle-like properties.

Schrödinger Equation

The fundamental equation of quantum mechanics, $\hat{H} ψ = E ψ$, which describes the behavior of electrons in atoms.

Atomic Orbital

A one-electron wave function ($ψ$) representing a region in space where the probability of finding an electron is high.

Principal Quantum Number ($n$)

A positive integer ($1, 2, 3...$) that determines the size and largely the energy of the orbital and identifies the main shell.

Azimuthal Quantum Number ($l$)

Also called the orbital angular momentum or subsidiary quantum number, it defines the three-dimensional shape of the orbital ($s, p, d, f$).

Magnetic Orbital Quantum Number ($m_l$)

Determines the spatial orientation of the orbital with respect to a standard coordinate axis; it has $(2l + 1)$ values for a given subshell.

Electron Spin Quantum Number ($m_s$)

The fourth quantum number that accounts for the intrinsic spin of the electron, taking values of $+1/2$ or $-1/2$.

Aufbau Principle

The rule that electrons occupy the lowest energy orbital available to them before entering higher energy orbitals ($1s, 2s, 2p, 3s...$).

Pauli Exclusion Principle

The rule that no two electrons in an atom can have the same set of four quantum numbers, meaning an orbital can hold a maximum of two electrons with opposite spins.

Hund’s Rule of Maximum Multiplicity

The rule that pairing of electrons in degenerate orbitals belonging to the same subshell does not occur until each orbital is singly occupied.

Effective Nuclear Charge ($Z_{eff}$)

The net positive charge experienced by outer-shell electrons due to the partial screening (shielding) of the nucleus by inner-shell electrons.