Atomic Structure Practice Flashcards

These vocabulary flashcards cover the fundamental concepts of atomic structure from Class 11 Chemistry, including subatomic particle discovery, historical atomic models, quantum theory, and electron configurations.

Subatomic Particles

The fundamental particles that make up atoms, specifically electrons, protons, and neutrons.

Cathode Rays

A stream of negatively charged particles traveling from the cathode to the anode, which travel in straight lines and can produce fluorescence.

Anode Rays (Canal Rays)

A stream of positively charged gas ions traveling from the region between electrodes toward the cathode; their properties depend on the nature of the gas used.

Phosphorescence

Delayed light emission after the energy source is removed, such as observed on zinc sulfide screens.

Charge-to-Mass Ratio of the Electron

Measured by J.J. Thomson as $-1.75 \times 10^{11}\,C/kg$, a value independent of the gas or cathode material used.

Fundamental Charge of an Electron

Determined by Millikan as $-1.60 \times 10^{-19}\,C$.

Mass of an Electron

$9.11 \times 10^{-31}\,kg$, calculated using the charge and charge-to-mass ratio.

Proton

The lightest and smallest positive ion, named by Rutherford in 1919 after ejecting hydrogen nuclei from nitrogen through alpha particle bombardment.

Neutron

A neutral heavy particle discovered by James Chadwick in 1932 by bombarding beryllium with alpha particles.

Thomson Model (Plum Pudding Model)

An atomic model proposing a positively charged sphere with electrons embedded uniformly throughout it like seeds in a watermelon.

Alpha Particles

Particles consisting of 2 protons and 2 neutrons (helium nucleus) with a charge of $+2$ and a mass number of $4$.

Rutherford Model (Nuclear Model)

An atomic model stating that the entire mass and positive charge of an atom is concentrated in a tiny, dense center called the nucleus, which electrons orbit.

Atomic Number ($Z$)

The number of protons present inside the nucleus of an atom.

Mass Number ($A$)

The total number of protons and neutrons collectively known as nucleons present inside the nucleus; calculated as $A = Z + N$.

Isotopes

Atoms of the same element having the same atomic number but different mass numbers due to a different number of neutrons.

Isobars

Atoms of different elements which have different atomic numbers but the same mass number.

Electromagnetic Waves

Waves consisting of alternating electric and magnetic fields perpendicular to each other and the direction of propagation, traveling at the speed of light ($3 \times 10^{8}\,m/s$).

Frequency ($\nu$)

The number of waves that pass through a given point in one second, measured in Hertz ($Hz$).

Wave Number ($\bar{\nu}$)

The number of waves present in a unit length, related to wavelength by the formula $\bar{\nu} = \frac{1}{\lambda}$.

Black Body

An ideal body that absorbs and emits radiation of all frequencies without reflection.

Planck's Quantum Theory

The theory that energy is emitted or absorbed in discrete packets called quanta, where energy is defined as $E = h \times \nu$.

Planck's Constant ($h$)

$6.62 \times 10^{-34}\,J\cdot s$.

Threshold Frequency ($\nu_0$)

The minimum frequency of incident radiation required to eject electrons from a metal surface in the photoelectric effect.

Work Function ($\phi$)

The minimum energy required to eject an electron from a metal surface, calculated as $\phi = h\nu_0$.

Emission Spectrum

A spectrum produced when radiation is emitted from a source as excited electrons return to lower energy states, appearing as bright colored lines on a dark background.

Absorption Spectrum

A spectrum appearing as dark lines on a continuous band, obtained when white light passes through a substance before being analyzed.

Rydberg Constant ($R_H$)

$109677\,cm^{-1}$, used in the Rydberg equation to calculate the wavelength of spectral lines in the hydrogen spectrum.

Balmer Series

The spectral series of hydrogen where electrons jump to the $n=2$ level, resulting in radiation in the visible region.

Bohr's Orbit (Stationary State)

Discrete orbits where electrons revolve around the nucleus with fixed energy and radius without losing or gaining energy.

de Broglie Hypothesis

The principle stating that all moving matter particles, including electrons, exhibit wave-like properties with a wavelength $\lambda = \frac{h}{mv}$.

Heisenberg's Uncertainty Principle

The principle that it is impossible to simultaneously determine the exact position and momentum of a microscopic particle, expressed as $\Delta x \times \Delta p \ge \frac{h}{4\pi}$.

Schrödinger's Wave Equation

The fundamental quantum mechanical equation $\hat{H}\psi = E\psi$ which describes the motion of electrons as waves.

Orbital

A three-dimensional space around the nucleus where the probability of finding an electron is maximum ($90\%$ probability).

Principal Quantum Number ($n$)

Indicates the main energy level or shell of an electron ($1, 2, 3, \dots$) and determines the size and energy of an orbital.

Azimuthal Quantum Number ($l$)

Defines the shape of the subshell and orbital ($0, 1, 2, \dots$ up to $n-1$), where $l=0$ is s, $l=1$ is p, and $l=2$ is d.

Magnetic Quantum Number ($m$)

Describes the orientation of orbitals in space, with values ranging from $-l$ to $+l$.

Spin Quantum Number ($s$)

Refers to the intrinsic spin of the electron, with possible values of $+\frac{1}{2}$ or $-\frac{1}{2}$.

Node

A region within an atom where the probability of finding an electron is zero.

Aufbau Principle

The rule that electrons fill orbitals in order of increasing energy, starting from the lowest available energy level.

Pauli Exclusion Principle

States that no two electrons in an atom can have the same set of four quantum numbers; an orbital can hold a maximum of 2 electrons with opposite spins.

Degenerate Orbitals

Orbitals belonging to the same subshell that have exactly the same energy level.

Hund's Rule of Maximum Multiplicity

When filling degenerate orbitals, electrons occupy each orbital singly with parallel spins before pairing begins.

Exchange Energy

Energy associated with the stability of half-filled or fully-filled subshells due to electrons with the same spin exchanging positions.