Chemistry Grade 11 - Unit 1: Atomic Structure and Periodic Properties of the Elements

Vocabulary flashcards covering the historical development of atomic theory, subatomic particle discovery, quantum mechanics, electronic configurations, and periodic trends as outlined in the Grade 11 Chemistry textbook.

Democritus

The ancient Greek philosopher (460–370 BC) who suggested that matter eventually ends up as tiny, indestructible particles called atomos, meaning "indivisible."

Dalton’s Atomic Theory

A theory developed by John Dalton in 1808 based on the laws of conservation of mass and definite proportions, emphasizing how atoms combine to form compounds.

Modern Atomic Theory

A framework beginning with the work of Dalton and evolving with the discovery of subatomic particles like the electron by J.J. Thomson in 1897.

Law of Conservation of Mass

One of the three fundamental laws of chemistry that served as a basis for Dalton’s atomic theory, stating that mass is neither created nor destroyed in a chemical reaction.

Law of Multiple Proportions

A law stating that when two elements form multiple compounds, the ratios of the masses of the second element that combine with a fixed mass of the first are small whole numbers.

Cathode Rays

Rays originating from the negative electrode (cathode) in a vacuum tube that J.J. Thomson concluded consist of negatively charged particles or electrons.

Electron ($e^-$)

A constituent of all matter discovered by J.J. Thomson with a mass of $9.109 \times 10^{-31}\,kg$ and a charge of $-1.602 \times 10^{-19}\,C$.

Radioactivity

The spontaneous emission of particles and/or radiation from the unstable nuclei of certain atoms such as uranium and radium.

Alpha ($\alpha$) rays

Radioactive emissions consisting of positively charged particles with a mass about four times that of a hydrogen atom, identical to helium nuclei.

Beta ($\beta$) rays

Radioactive emissions consisting of electrons coming from inside the nucleus that are deflected by negatively charged plates.

Gamma ($\gamma$) rays

High-energy radioactive radiation with no charge that is not affected by external electric or magnetic fields.

Plum-pudding model

An atomic model proposed by J.J. Thomson where electrons and protons were randomly distributed in a positively charged cloud.

Ernest Rutherford

A scientist who overturned the plum-pudding model through his gold foil experiment, leading to the discovery of the atomic nucleus.

Neutron ($n$)

A neutral nuclear particle discovered by James Chadwick in 1932 with a mass of $1.67493 \times 10^{-27}\,kg$.

Proton ($p$)

A nuclear particle having a positive charge equal in magnitude to the electron and a mass of $1.67262 \times 10^{-27}\,kg$, discovered by Rutherford.

Atomic mass unit (amu)

A mass unit equal to $\frac{1}{12}$ the mass of a carbon-12 atom.

Atomic number ($Z$)

The number of protons in the nucleus of each atom of an element, which is unique to each element.

Mass number ($A$)

The total number of protons and neutrons in the nucleus of an atom.

Isotopes

Atoms of an element that have different numbers of neutrons and therefore different mass numbers but identical atomic numbers.

Electromagnetic Radiation (EMR)

The emission and transmission of energy in the form of electromagnetic waves, which consist of electric and magnetic field components.

Wavelength ($\lambda$)

The distance a wave travels during one cycle, typically expressed in meters ($m$), nanometers ($nm$), or angstrom ($\text{Å}$).

Frequency ($\nu$)

The number of cycles a wave undergoes per second, expressed in units of $1/\text{second}$ or hertz ($Hz$).

Quantum

The name given to the discrete quantities of energy that Max Planck proposed atoms and molecules can emit or absorb ($E = h\nu$).

Planck’s constant ($h$)

A physical constant used to calculate the energy of a quantum, with a value of $6.63 \times 10^{-34}\,J \cdot s$.

Photoelectric effect

The phenomenon where electrons are ejected from the surface of metals exposed to light of at least a certain threshold frequency ($\nu_o$).

Photon

A unit or particle of light energy defined as a packet of electromagnetic radiation.

Atomic spectrum

Also called an emission spectrum, it is a series of individual colored lines characteristic of a specific element, produced as electrons return to lower energy states.

Niels Bohr

A Danish physicist who proposed that electrons travel in circular orbits and that absorption or emission of light mirrors quantized energy changes within the atom.

Ground state

The lowest energy state or ground level of an atom, corresponding to $n = 1$ in the Bohr model.

Excited state

Any energy state of an atom higher than its ground state, reached when an electron absorbs energy.

Rydberg constant ($R_H$)

A constant used to calculate the energy levels of a hydrogen atom, valued at $2.18 \times 10^{-18}\,J$.

Wave-particle duality

The character of matter and energy exhibiting both wave and particulate properties, first proposed by Louis de Broglie.

Heisenberg uncertainty principle

A principle stating it is not possible to know both an electron's position and its momentum ($p = mv$) with total certainty at the same time.

Orbital

A region of space around the nucleus where there is a high probability (about 90%) of finding an electron.

Principal quantum number ($n$)

A quantum number describing the main energy level or shell an electron occupies, taking values of $1, 2, 3, 4, \dots$.

Angular momentum quantum number ($\ell$)

Also known as the azimuthal quantum number, it designates the shape of atomic orbitals and takes values from $0$ to $n-1$.

Magnetic quantum number ($m_\ell$)

A quantum number related to the orientation of an orbital in space, with values between $-\ell$ and $\ell$.

Electron spin quantum number ($m_s$)

A quantum number describing the magnetic property or spin of an electron, which can be either $+\frac{1}{2}$ or $-\frac{1}{2}$.

Aufbau Principle

A building-up scheme where electrons occupy the lowest-energy orbital available before entering higher energy orbitals.

Hund’s Principle

A rule stating that degenerate orbitals are each occupied by a single electron before the second electron of opposite spin enters.

Pauli’s Exclusion Principle

The rule that no two electrons in the same atom can have the same four quantum numbers.

Valence electrons

Electrons in the outermost principal quantum level of an atom which are important in chemical bonding.

Periodic Law

The modern law stating that physical and chemical properties of elements repeat periodically when elements are arranged by increasing atomic number.

Metalloid

An element that has physical and chemical properties intermediate between those of metals and non-metals.

Atomic radius

The estimated size of an atom, often determined as half the distance between the nuclei of two adjacent touching atoms.

Effective nuclear charge ($Z_{eff}$)

The net nuclear charge an electron actually experiences after accounting for the shielding caused by inner electrons.