Quick questions on Atomic structure, isotopes and relative atomic mass (QCE Chemistry Unit 1)

The proton number (atomic number, Z) defines the element. The mass number (A) is the total count of protons plus neutrons. A neutral atom has equal numbers of protons and electrons.

where X is the element symbol, A is the mass number (top), and Z is the atomic number (bottom). Examples:

Isotopes of an element have the same number of protons but different numbers of neutrons, so the same atomic number Z but different mass number A.

A mass spectrometer ionises a sample (usually by electron impact, knocking out one electron to form a singly charged cation), accelerates the ions through an electric field, separates them by mass-to-charge ratio (m/z) in a magnetic field, and detects each beam. The output is a mass spectrum: a plot of relative abundance against m/z.

The relative atomic mass (Ar) is the weighted mean of the isotopic masses by their natural abundance:

If the question gives Ar and the two isotopic masses, solve for the abundances using x for one fraction and (1 - x) for the other:

Atomic structure feeds directly into electron configuration (the next dot point) and then into bonding (Topic 2) and the mole concept (Topic 3). The molar mass used in stoichiometry is numerically equal to Ar in g/mol, so accurate weighted-mean reasoning here transfers to mass-to-mole calculations later.

Cl-35 (mass 34.97, abundance 75.78 percent), Cl-37 (mass 36.97, abundance 24.22 percent).

Peaks at 24, 25, 26 with heights 79, 10, 11.

Mass number is an integer for a specific isotope (A). Relative atomic mass is the weighted average across isotopes, usually a non-integer.

(35 + 37) / 2 = 36 is wrong because the isotopes are not equally abundant. Always weight by fractional abundance.

Al^3+ has 10 electrons, not 13. The 3+ means three have been lost.

Singly charged atomic ions give m/z equal to the mass number, but multiply charged species (rare for QCAA Unit 1) would not.