Atomic & Electronic Structure and Chemical Bonding Practice Flashcards

A comprehensive set of vocabulary flashcards covering basic atomic theory, historical models, the electronic structure of the atom, ionisation trends, and the various forms of chemical bonding (ionic, covalent, and metallic) based on chapter one and two notes.

Mass number ($A$)

The total number of protons and neutrons in an atom.

Atomic number ($Z$)

The number of protons in an atom; also called the proton number.

Isotope

Atoms of the same element with different mass numbers OR atoms of the same element with the same atomic number but different mass numbers.

Ion

A species in which there is a surplus or lack of electrons.

Proton

A sub-atomic particle with a relative mass of $1$ and a relative charge of $+1$.

Neutron

A sub-atomic particle with a relative mass of $1$ and a relative charge of $0$.

Electron

A sub-atomic particle with a relative mass of $1/2000$ and a relative charge of $-1$.

First ionisation energy

The amount of energy required to remove the outermost electron from a gaseous atom in its ground state.

Valence electrons

The electrons in the outer energy level of an atom.

Solid sphere model

John Dalton's 1803 model stating that atoms are indivisible, those of a given element are identical, and compounds are combinations of different types of atoms.

Plum pudding model

J.J. Thomson's 1904 model showing the atom as composed of electrons scattered throughout a spherical cloud of positive charge.

Nuclear model

Ernest Rutherford's 1911 model where positive charge is concentrated in the centre (nucleus) and the atom is mostly empty space.

Planetary model

Niels Bohr's 1913 model stating that electrons move around the nucleus in orbits of fixed sizes and energies; electron energy is quantised.

Quantum model

Erwin Schrödinger's 1926 model where electrons move in waves and exist in 'clouds of probability' called orbitals rather than set paths.

Cation

A positive ion created by a lack of electrons.

Anion

A negative ion created by a surplus of electrons.

Orbitals

Areas in which it is statistically likely to find electrons; all types can hold a maximum of $2$ electrons each.

Pauli’s exclusion principle

States that each orbital can only contain $2$ electrons, represented as arrows pointing in opposite directions to show opposite rotation.

Hund's rule

Electrons will fill an unoccupied orbital within the same energy level before doubling up in a sub-orbital.

Aufbau diagram

A diagram using boxes to represent $s$ and $p$ orbitals for each energy level, filled from the lowest energy level upwards.

Ionic bond

A transfer of electrons followed by electrostatic attraction between the resulting oppositely charged ions.

Lattice

An orderly, fixed-ratio three-dimensional arrangement where positive and negative ions pack together.

Stock notation

The use of Roman numerals in parentheses to distinguish between different oxidation states (charges) of transition metals in compounds.

Polyatomic ions

Ions comprised of two or more different types of atoms bonded together in a specific ratio with an overall non-zero charge.

Covalent bond

The sharing of at least one pair of electrons between two non-metal atoms.

Electronegativity

A measure of the tendency of an atom to attract a bonding pair of electrons.

Non-polar covalent bond

A bond where electrons are shared equally, occurring when the difference in electronegativity is zero.

Polar covalent bond

A bond where electrons are shared unequally due to a non-zero difference in electronegativity, leading to the formation of a dipole.

Intramolecular bond

A bond which occurs between atoms within molecules, such as a covalent bond.

Intermolecular force

A force of attraction between molecules, ions, or atoms of noble gases.

Allotrope

Different structural forms of the same element, such as diamond and graphite for carbon.

Giant covalent substance

A substance containing a large network of covalent bonds, also known as a macromolecular structure or network solid.

Metallic bond

A bond between a positive kernel and a sea of delocalised electrons.

Sea of delocalised electrons

Free-moving electrons created by the overlapping of valence orbitals in metal atoms.

Law of conservation of matter

States that matter cannot be created or destroyed; it can only be transferred from one place to another.

Stoichiometric ratio

The whole-number ratio of particles in a balanced chemical equation that ensures no matter is lost or created.

Lewis structure

A representation of atoms and their valence electrons using dots or crosses to show bonding.