Atomic Structure and Periodicity

Comprehensive vocabulary flashcards covering atomic structure, isotopes, relative masses, mass spectrometry, ionisation energy, electron configurations, and periodicity based on the lecture notes.

Rutherford scattering experiment

A 1911 experiment that discovered the current model of the atom, consisting of a small, dense central nucleus surrounded by orbiting electrons in shells.

Proton

A fundamental particle found in the nucleus with a relative charge of $+1$ and a relative mass of $1$.

Neutron

A fundamental particle found in the nucleus with a relative charge of $0$ and a relative mass of $1$.

Electron

A fundamental particle that orbits the nucleus in shells with a relative charge of $-1$ and a relative mass of $\frac{1}{1840}$.

$$2n^2$$

The formula used to calculate the maximum number of orbiting electrons that can be held by any single shell, where $n$ is the number of the shell.

Mass number (A)

The sum of protons and neutrons in an atom.

Atomic number (Z)

The number of protons in an atom, also referred to as the proton number.

Isotopes

Atoms of the same element with the same atomic number but a different number of neutrons, resulting in different mass numbers and physical properties.

Deuterium

An isotope of Hydrogen containing $1$ proton and $2$ neutrons.

Tritium

An isotope of Hydrogen containing $1$ proton and $3$ neutrons.

Relative atomic mass ($Ar$)

The mean mass of an atom of an element, relative to one twelfth of the mean mass of an atom of the carbon-12 isotope.

Relative isotopic mass

The isotopic mass of an isotope relative to one twelfth of the mean mass of an atom of the carbon-12 isotope.

Relative molecular mass ($Mr$)

The mean mass of a molecule of a compound, relative to one twelfth of the mean mass of an atom of the carbon-12 isotope.

Relative formula mass

A value similar to $Mr$ but specifically used for compounds with giant structures.

Mass Spectrometry

An analytical technique used to identify different isotopes and find the overall relative atomic mass of an element.

Ionisation (TOF Stage 1)

The process where a sample is vapourised and a high voltage is applied, removing electrons to produce $+1$ charged ions.

Acceleration (TOF Stage 2)

The process where positively charged ions are accelerated towards a negatively charged detection plate.

Ion Drift (TOF Stage 3)

The stage where ions are deflected by a magnetic field into a curved path, the radius of which depends on the charge and mass of the ion.

Detection (TOF Stage 4)

The stage where positive ions hit a negatively charged plate and gain an electron, producing a current proportional to their abundance.

Mass to charge ratio (m/z)

The ratio of an ion's mass to its charge; if a $2+$ ion is produced, this value is halved, appearing at half the expected trace on a spectra.

Molecular ion peak

The tallest peak on a mass spectrum which corresponds to the relative molecular mass of the molecule, formed from the $M^+$ species.

Chlorine Spectra Ratios

A characteristic pattern of $3:1$ for $Cl^+$ ions and $3:6:9$ for $Cl_2^+$ ions due to isotope combinations.

Ionisation energy

The minimum energy required to remove one mole of electrons from one mole of atoms in a gaseous state, measured in $kJ\,mol^{-1}$.

First ionisation energy trend (Period)

Increases along a Period due to decreasing atomic radius and greater electrostatic forces of attraction.

First ionisation energy trend (Group)

Decreases down a Group due to increasing atomic radius and electron shielding reducing electrostatic attraction.

Successive ionisation energies

The energy required to remove further electrons; large increases indicate an electron is being removed from an orbital closer to the nucleus (a different energy level).

Orbitals

Clouds of negative charge ($s$, $p$, $d$, and $f$) that can hold up to two electrons with opposite spins.

Subshell capacities

The $s$-subshell holds $2$ electrons, $p$-subshell holds $6$ electrons, and $d$-subshell holds $10$ electrons.

Spin

A property where electrons pair up with opposite arrows in an orbital to make the atom as stable as possible.

Periodicity

The study of regularly repeating patterns of physical, atomic, and chemical properties within the Periodic Table.

Atomic Radius Trend

Decreases along a period due to increased nuclear charge pulling electrons closer; increases down a group as extra electron shells are added.

Electron shielding

The effect where inner electron shells create a 'barrier' that blocks the attractive forces of the nucleus on outer electrons.

Metallic bonding

The bonding found in elements like Lithium, Beryllium, Sodium, Magnesium, and Aluminium, where melting points increase with greater positive ionic charge and more free electrons.

Giant covalent lattices (Macromolecular)

Structures with very high melting points like Boron, Carbon, and Silicon, held by strong covalent bonds in up to three dimensions.

Van der Waals forces

Weak intermolecular forces holding small, simple covalent molecules (like $N_2$, $O_2$, $P_4$, $S_8$) or noble gases (Argon) together, resulting in lower melting points.