2.1.1 Atomic structure and isotopes

What are the three subatomic particles, their positions, relative masses, and charges?

Proton: nucleus, mass 1, charge +1; Neutron: nucleus, mass 1, charge 0; Electron: orbitals, mass 1/1800, charge –1

Define atomic number (Z).

Number of protons in the nucleus.

Define mass number (A).

Total number of protons + neutrons.

How do you calculate the number of neutrons?

Neutrons = A – Z

Define isotope.

Atoms of the same element with the same number of protons but different numbers of neutrons.

Why do isotopes have the same chemical properties?

They have the same electronic structure/ configuration

Why do isotopes have different physical properties?

They have different masses.

Define relative isotopic mass.

Mass of one isotope compared to 1/12 of the mass of a carbon-12 atom.

Define relative atomic mass (Ar).

Weighted mean mass of an atom compared to 1/12 of the mass of a carbon-12 atom.

Define relative molecular mass (Mr).

Average mass of a molecule compared to 1/12 of the mass of a carbon-12 atom.

Define relative formula mass.

Mass of a formula unit compared to 1/12 of the mass of a carbon-12 atom.

Give the equation for relative atomic mass using percentage abundance.

Ar = Σ(isotopic mass × % abundance) ÷ 100

Give the equation for relative atomic mass using relative abundance.

Ar = Σ(isotopic mass × relative abundance) ÷ total relative abundance

What does mass spectrometry allow us to determine?

Relative isotopic masses

relative abundances

relative atomic mass of an element.

How do you work out the probability of molecular ion peaks for diatomic molecules (e.g. Cl2, Br2)?

By multiplying isotope abundances and considering all possible pairings.

Outline the main historical models of the atom.

Dalton → Thomson (plum pudding) → Rutherford (nuclear model) → Bohr (shells) → Quantum model

State differences between isotopes of the same element

• different number of neutrons

• different physical properties

• different atomic/ mass numbers