Atoms and atomic structure

What the exam asks

Expect data rather than recall: a table of successive ionisation energies to place an element in a group, a mass spectrum to convert into a relative atomic mass, or two species to rank by radius. The trap that catches most candidates is the isoelectronic ranking. Told that Na+ and F- both have ten electrons, they call the radii equal, or make the sodium bigger because sodium is the bigger atom. Both are wrong, and the fix is the same each time: when the electron count is fixed, only the proton count is still voting, so more protons means smaller. The second trap is averaging isotope masses instead of weighting them, which turns 35.5 into 36.

This topic covers where electrons sit, how tightly the nucleus holds them, how big the resulting atom or ion is, and how isotopes show up in a mass spectrum. Nearly all of it runs off one idea: an electron feels the nuclear charge minus whatever is shielding it, and every trend on the periodic table is that quantity changing.

The exam does not ask you to recite configurations. It gives you a table of successive ionisation energies, or a mass spectrum, or two radii, and asks what the numbers mean. So the useful thing to carry in is not the order of the orbitals but the reasoning: a jump in ionisation energy means a new shell has been broken into, a smaller radius across a period means the same shell pulled tighter, and a peak in a mass spectrum is one isotope, not one element.

Mass spectrometry is examined as arithmetic with a concept attached. The relative atomic mass is a weighted average over isotopes, which is why it is rarely a whole number, and why chlorine sits at 35.5 without any chlorine atom weighing 35.5.

What to hold

• Effective nuclear charge is the nuclear charge an outer electron actually feels once inner electrons have shielded it, and it rises across a period because protons are added while the shielding shell stays the same.
• Atomic radius falls across a period and rises down a group: across, the same shell is pulled in harder; down, a whole new shell is added and that beats the extra charge.
• The 4s orbital fills before 3d but is emptied first, so Fe2+ is [Ar]3d6 and not [Ar]3d4 4s2.
• First ionisation energy rises across a period and falls down a group, with two dips in each period: at group 13, where the electron leaving is a p rather than an s, and at group 16, where it leaves a paired p orbital and repulsion helps it go.
• Every successive ionisation energy is larger than the one before it, because the electron is being pulled off a species that is already more positive.
• A large jump between successive ionisation energies marks the point where the outer shell is exhausted and a core shell is being broken into, so the number of electrons removed before the jump is the group's outer electron count.
• A cation is smaller than its parent atom and an anion is larger, because the electron count changed while the proton count did not.
• Within an isoelectronic series, every species has the same electron count, so radius falls as nuclear charge rises: N3- > O2- > F- > Na+ > Mg2+ > Al3+.
• Isotopes differ in neutron number, so they differ in mass and in nuclear stability but behave the same chemically, because chemistry is done by the electrons.
• Relative atomic mass is the abundance-weighted average of the isotope masses, not the arithmetic mean of them, so it always lies nearer the more abundant peak.
• A mass spectrometer separates ions by mass-to-charge ratio, and with the usual singly charged ions the m/z value reads directly as a mass.
• An M and M+2 pair in a roughly 3:1 ratio signals one chlorine; a roughly 1:1 pair signals one bromine.