Atoms are neutral because the number of protons is equal to the number of electrons. Ions are charged
because the number of protons does not equal the number of electrons. You can calculate the number of
neutrons by doing mass number (protons + neutrons) – atomic number (protons)
Isotopes:
Isotopes are different atomic versions of the same element, with the same number of protons, but a
different number of neutrons. Usually only a small number of isotopes are actually stable. Isotopes react in
chemically identical ways as they have the same electron configuration, but they have different atomic
masses due to different number of neutrons
Relative Masses:
Relative Atomic Mass -> Mean weighted average mass of 1 atom of an element, taking into account
isotopes and abundances, relative to 1/12th the mass of a 12C atom
Relative Molecular Mass -> Mean weighted average mass of 1 molecule, taking into account
isotopes and abundances, relative to 1/12th the mass of a 12C atom
Remember to check if the question is asking for RELATIVE masses or not!
Electron Configuration:
Electrons are organised into orbitals (s,p,d,f ). The s orbital can hold 2 electrons, p orbitals can hold 6 and d
orbitals can hold 10. Electrons like to occupy the lowest energy orbital completely by themselves, so
electrons will always fill up an orbital without pairing up initially. As the orbital is filled up, the electrons
then pair up. This is a slightly higher energy arrangement due to electron spin repulsion (between the
similarly charged electrons), reducing the ESA of the electron to the nucleus.
This is the set filling pattern for electron shells. Note how the 4s orbital fills up
before the 3d orbital. This is because a more energetically stable arrangement will
occur. However, the 4s orbital will also empty before the 3d orbital, so ensure you
write out the entire electron configuration before finding the ion!
Copper and Chromium are the exceptions to this rule. They promote a single
electron from their 4s orbital into the 3d orbital as the 3d orbital is filled up either
unpaired or paired. This is a lower energy arrangement however.
Each single headed arrow represents a single electron
Ionisation Energy:
This is the enthalpy change that occurs when 1 mole of electrons is removed from 1 mole of gaseous
atoms, forming 1 mole of gaseous 1+ ions. There are 3 factors affecting the I.E;
Nuclear Charge -> Larger = More ESA = Higher I.E.
Electron shielding -> More = Lower ESA = Lower I.E.
Atomic Radius -> Larger = Less ESA = Lower I.E.
Generally, across period 3 the I.E. increases as nuclear charge increases, atomic radius decreases and
electron shielding is the same. Down groups though the I.E. decreases, although nuclear charge increases,
so to do the atomic radius and electron shielding, giving a lower I.E.
There are 2 exceptions across period 3 however; Mg->Al has a small decrease (electron in new p suborbital, slightly further away .: less ESA) and P->S also has a small decrease (electrons pair up in new p suborbital .: spin repulsion giving less ESA)