Electron Configurations

Outline

• Tracking Quantum States and Resulting Orbitals
  • Shells and Configurations
  • Box Notation
• Practice with Concepts
• Discussion Questions
• Optional Website Reading

Tracking Quantum States and Resulting Orbitals

We use a number of terms to describe the location and energy levels of the electrons in a ground state atom. The shell is the n energy level. The subshells are the l level orbitals, which differ by shapes and are usually designated by s, p, d, and f, rather than by numerical values.

Shells and Configurations

We assign electrons to the ground state in the order n+l, lower value to higher value, and not simply in the order n, then l. This is because the actual orbitals overlap somewhat. In practical terms, this means we fill the 4s level (4 + 0) before we fill the 3d level (3 + 2). You notice this is when you look at the periodic table: the third level electrons marking the transition metals occur on row 4, rather than row 3.

As subshells fill, the energy level of the higher subshells is affected by the presence of electrons in the lower subshells. The inner electrons shield the outer eleectrons from some of the the positive attraction of the nucleus. They also repel each other and the outer electrons. So an electron in the 3s orbital of the hydrogen atom would not be shielded; it would have a slightly different energy than an electron in the 3s orbital of a helium atom, which would have another electron in a lower state. The remaining nuclear charge felt by the outer electrons is called the effective nuclear charge.

In a given atom, electrons will occupy as many orbitals as possible. If an atom has three electrons at the p level, they will occupy three different p orbitals (p_x, p_y, and p_z), rather than filling up one and part of another (two in the p_x, one in the p_y orbital). We can thus determine the orbital configuration for an element by knowing the number of electrons that it has.

Fluorine is element 9. It has 9 electrons. Starting from the bottom up, we do the following analysis:

Energy level 1: n = 1, l = 0, there is one s orbital which can hold two electrons.

Energy level 2: n = 2, l = 0 or 1.

We consider first n = 2, l = 0. There is one s orbital which can hold two electrons.

We now consider n = 2, l = 1, m = -1, 0, and +1. Since l = 1, we now have a p orbital; there can be up to three of these (one for each value of m), and each can hold 2 electrons, for a total of 6 at the 2p level. We had 9 to start with; we've put 4 in the 1s and 2s orbitals, so the remaining 5 will fit in the 2p orbitals. We show this by saying 2p⁵, which means 5 electrons in the 2p level (not p to the fifth power!).

The electron configuration for fluorine is thus 1s², 2s², 2p⁵.

Box notation

Box notation is just a way of visualizing spin pairs within electron orbitals, given the 2-electron per orbital limit of the Pauli exclusion principle. When we want to represent a configuration, we draw one box for each pair of electrons in the configuration.

Consider the element silicon, group 4A.

1s

2s

2p

3s

3p

It has full 1s, 2s, 2p and 3s orbitals, so each orbital is shown with two electrons represented by arrows emphasizing the opposite (and cancelling) spins of each pair. The 3p level has only two electrons, which spread out into two different orbitals. We actually have some options for each of these orbitals: the spins could be aligned (paramagnetic) or opposite (cancelling).

We can also use box notation to show excited states. Consider a situation where the silicon is heated enough that one of the 3s electrons jumps into a 3p orbital:

1s

2s

2p

3s

3p

The isolated single electrons in each of the 3s and 3p orbitals could be aligned for maximum magnetic effect, or their spins randomized to reduce magnetic fields or even cancel them out. Excited electrons increase the magnetic field possibilities.

Practice with the Concepts

Orbital configurations

What is the orbital configuration of copper?

Magetism

Discussion Questions

• How does electron shell analysis help us understand energy levels for electrons?
• Are electrons always in the ground state?

Optional Readings

ChemGuide has a good article on ion electron configurations with lots of examples.

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