NZ NCEA Level 3 Chemistry

Energy levels, sublevels and orbitals

Electrons in an atom occupy energy levels numbered $n = 1, 2, 3, \dots$, with the lowest number closest to the nucleus and lowest in energy. Each level is divided into sublevels, named s, p, d and f. An s sublevel holds one orbital, a p sublevel three, a d sublevel five, and each orbital holds at most two electrons, which must have opposite spins. So the sublevels hold 2 (s), 6 (p) and 10 (d) electrons.

An electron configuration lists the occupied sublevels with the number of electrons as a superscript. In $3\mathrm{p}^4$ the 3 is the energy level, the p is the sublevel and the 4 is the number of electrons in it.

The order of filling

Electrons go into the lowest-energy sublevel available. Within the first 36 elements the order is

The 4s sublevel fills before the 3d because it is lower in energy in a neutral atom, so potassium and calcium add their outer electrons to 4s and the 3d sublevel starts filling at scandium. The 4p sublevel then fills from gallium to krypton.

The periodic table follows the same order: groups 1 and 2 are the s block, groups 13 to 18 the p block, and the transition metals the d block.

Writing configurations

The number of electrons in a neutral atom is its atomic number. Fill the sublevels in order until the electrons run out.

For iron the sublevels are filled in the order 4s then 3d, but the configuration is conventionally written with the levels in numerical order, 3d before 4s. Either way the electrons present are the same, and the order written matters when the atom loses electrons, since the 4s electrons are the ones removed first.

A shorthand replaces the core by the symbol of the preceding noble gas in brackets, so sulfur is $[\mathrm{Ne}]3\mathrm{s}^2 3\mathrm{p}^4$ and iron is $[\mathrm{Ar}]3\mathrm{d}^6 4\mathrm{s}^2$. The outer electrons are then easy to see. The number of electrons in the highest energy level of a main-group atom is its valence electron count, which is the group number for groups 1 and 2 and the group number minus 10 for groups 13 to 18.

Chromium and copper

Two elements in the first transition series do not follow the filling order, and their configurations have to be learned:

In each case one electron moves from 4s into 3d. The 3d and 4s sublevels are close in energy in these atoms, so the atom settles into the lower-energy arrangement that results.

Configurations of ions

A cation is formed by removing electrons, and an anion by adding them. Electrons are removed from the highest energy level first. For main-group elements this gives the configuration of a noble gas.

• $\ce{Na+}$ has 10 electrons: $1\mathrm{s}^2 2\mathrm{s}^2 2\mathrm{p}^6$, the same as neon.
• $\ce{Mg^2+}$ has 10 electrons: $1\mathrm{s}^2 2\mathrm{s}^2 2\mathrm{p}^6$.
• $\ce{Cl-}$ has 18 electrons: $1\mathrm{s}^2 2\mathrm{s}^2 2\mathrm{p}^6 3\mathrm{s}^2 3\mathrm{p}^6$, the same as argon.
• $\ce{O^2-}$ has 10 electrons: $1\mathrm{s}^2 2\mathrm{s}^2 2\mathrm{p}^6$.
• $\ce{Ca^2+}$ has 18 electrons: $1\mathrm{s}^2 2\mathrm{s}^2 2\mathrm{p}^6 3\mathrm{s}^2 3\mathrm{p}^6$.

Species with the same electron configuration are isoelectronic. $\ce{N^3-}$, $\ce{O^2-}$, $\ce{F-}$, $\ce{Na+}$, $\ce{Mg^2+}$ and $\ce{Al^3+}$ are all isoelectronic with neon.

Transition metal ions

When a transition metal forms a cation it loses its 4s electrons before its 3d electrons, even though 4s fills first. This is because once 3d has electrons in it, the 3d sublevel lies lower in energy than 4s.

Worked example: the configuration of $\ce{Ni^2+}$. Nickel has atomic number 28, so its atom is $1\mathrm{s}^2 2\mathrm{s}^2 2\mathrm{p}^6 3\mathrm{s}^2 3\mathrm{p}^6 3\mathrm{d}^8 4\mathrm{s}^2$. The ion has lost two electrons, both from 4s, leaving 26 electrons: $1\mathrm{s}^2 2\mathrm{s}^2 2\mathrm{p}^6 3\mathrm{s}^2 3\mathrm{p}^6 3\mathrm{d}^8$.

The zinc ion shows a point worth stating: its 3d sublevel is full, so $\ce{Zn^2+}$ is $3\mathrm{d}^{10}$, and copper(I) and copper(II) differ by one 3d electron.

Interpreting a configuration

A configuration can be used to identify the species. A neutral atom with the configuration $1\mathrm{s}^2 2\mathrm{s}^2 2\mathrm{p}^6 3\mathrm{s}^2 3\mathrm{p}^3$ has 15 electrons, so it is phosphorus, with five valence electrons in the third energy level. An ion with $1\mathrm{s}^2 2\mathrm{s}^2 2\mathrm{p}^6 3\mathrm{s}^2 3\mathrm{p}^6$ could be $\ce{Cl-}$, $\ce{S^2-}$, $\ce{K+}$ or $\ce{Ca^2+}$, which is why a charge or atomic number has to be given to decide. The number of energy levels and the number of valence electrons are what the trends in atomic radius and ionisation energy are explained with.