DP Chemistry: Electron configuration
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Electron configuration

2.2 Electron configuration (3 hours)

Pause for thought

Practically the diagram on the left can be used to determine the order in which orbitals are filled by electrons. It is known either as the Madelung rule or as the Klechkowski rule although in practice many people know it simply as the n+1 rule. The rule does not actually work for chromium and copper. The order can be more logically deduced by looking at graphs of ionization energies or from the position of elements in the periodic table but what is the theory behind it?

Under 'Applications and skills' in sub-topic 2.2 it states “Application of the Aufbau principle, Hund's rules and the Pauli exclusion principle to write electron configurations for atoms and ions up to Z = 36.” Technically knowledge of quantum numbers is required to do this although quantum numbers are not on the IB programme.

Pauli’s exclusion principle was formulated by Wolfgang Pauli in 1925. It applies to all particles collectively known as fermions which include protons and neutrons as well as electrons. Fermions are particles with half spin and the value of the spin quantum number can either be + ½ or – ½. At this level we restrict it to just electrons and Pauli’s exclusion principle states that no two electrons within an atom can possess the same four quantum numbers. The principle quantum number (n) refers to the main energy level and can have values 1, 2, 3….. The azimuthal quantum number (l) is related to angular momentum and has the values 0 to n-1. It determines the type of orbital or sub energy level; s (l=0), p (l=1), d (l=2) and f (l=3). Thus when n = 1 there are only s orbitals as l can only have the value of zero whereas when n = 3 there are s, p and d orbitals. The magnetic quantum number (m) has values –l, -(l-1)...0...(l+1), l and determines the number of orbitals within a sub-level. For example, when l = 1 there are three values -1, 0, and +1 which explains why there are three p orbitals (px, py and pz). Finally there is the spin quantum number (s) which can have the values of + ½ or – ½ which explains why each orbital can contain only two electrons. Pauli’s exclusion principle is therefore the basis of the Aufbau principle. Hund’s rules are required to state that electrons fill up orbitals of the same type singly with the same spin before they spin pair. This explains why there is a drop in first ionization energies between nitrogen (1s22s22px12py12pz1) and oxygen (1s22s22px22py12pz1) as it is energetically easier to remove the spin paired electron in oxygen. Hund's rules can also be used to explain the chromium and copper exceptions.

Nature of Science

Another example of the improvement in apparatus leading to developments in science is the use of electric and magnetic fields in Thomson's cathode ray tubes.
Quantum mechanics provides the basis for current models of the atom superseding the Bohr model.
Natural phenomena can be explained by theories e.g. line spectra can be explained by the Bohr model of the atom.

Learning outcomes

After studying this topic students should be able to:


Understand:

  • When electrons that have been excited by gaining energy return to a lower energy level in an atom they emit photons and produce an emission spectrum.
  • Evidence for the existence of electrons in discrete energy levels comes from the line emission spectrum of hydrogen, in which the lines converge at higher energies.
  • The main energy level (or shell) is identified by an integer number, n, and can hold a maximum number of electrons equal to 2n2.
  • The main energy levels can be split into s, p, d and f sub-levels with successively higher energies.
  • Sub-levels contain a fixed number of orbitals. An orbital is a volume of space where there is a high probability of finding an electron.
  • Each orbital has a defined energy state for a given electronic configuration and chemical environment.
  • Each orbital can contain a maximum of two electrons each with opposite spin.


Apply their knowledge to:

  • Describe the relationship between colour, wavelength, frequency and energy across the electromagnetic spectrum.
  • Distinguish between a continuous spectrum and a line spectrum.
  • Describe the emission spectrum of the hydrogen atom, including the relationships between the lines and energy transitions to the first, second and third energy levels.
  • Recognize the shape of an s atomic orbital and the px, py and pz atomic orbitals.
  • Apply the Aufbau principle, Hund’s rule and the Pauli exclusion principle to deduce electron configurations for atoms and ions up to Z = 36.

Clarification notes

The electromagnetic spectrum is given in section 3 of the data booklet.

The names of the different series in the hydrogen line emission spectrum are not required.

Full electron configurations (e.g. 1s22s22p63s23p64s23d3) and condensed electron configurations (e.g. [Ar] 4s23d3) should be taught.

Orbital diagrams should be used to represent the character and relative energy of orbitals. Orbital diagrams refer to arrow-in-box diagrams, e.g. for carbon.

The electron configurations of Cr and Cu as exceptions should be covered.

International-mindedness

CERN (The European Organization for Nuclear Research) is run by 20 European member states and involves scientists from many other countries.

It operates the world’s largest particle accelerator and its detectors are used to study the fundamental constituents of matter.

Teaching tips

This is quite a large topic and you will need to tread carefully as students are asked to apply a lot of knowledge without really being required to understand the theories and principles upon which it is based.

Using flame tests and/or gas discharge tubes linked to a spectrometer you can use the emission spectra to provide evidence that electrons are in fixed energy levels.

It is good that students now need to understand the concept of orbitals and the splitting of the main energy levels into s, p, d and f sub-levels.

Even though it is not on the Standard Level programme for this topic (it is on the Higher Level under 12.1 and is mentioned in Topic 3.2) I still get students to plot the graph of first ionization energy against atomic number for the first twenty elements and then show them the graph for the first eighty or so elements (see Topics 2 & 12). The shape of this graph not only shows periodicity but can easily be related to the existence of sub-levels and how many electrons can occupy each sub-level. Using this knowledge the electron configuration for each element can then be deduced. I feel giving them this evidence, which they can easily comprehend, is better than just telling them about electron configurations with no real evidence at all. Ultimately students should understand how they can use the s, p, d & f blocks in the periodic table to deduce the electron configuration (see Topic 3).

Stress the importance of energy levels and 'going into the lowest available level' and this can then be used to explain the two main 'exceptions'. Chromium is [Ar]4s13d5 as it has all the 4s and 3d electrons spinning in the same direction which is more energetically favourable than [Ar]4s23d4. Similarly copper is [Ar]4s13d10.

It is also worth explaining that when electrons are removed from transition metals the ion formed has one 4s electron less than the neutral atom. This is because the protons in the nucleus of the positive ion 'pull' the energy levels closer to the nucleus so 3d becomes lower in energy than 4s. See the separate page on The electron configuration of scandium.


For the shapes of the orbitals I think they should understand a little about the wave nature of the electron (you could perhaps mention de Broglie) and Heisenberg's Uncertainty Principle so that they realise that the shapes of the different orbitals represent a volume of space where there is a high probability the electron (or pair of electrons) will be located.

Study guide

Pages 11 - 12

Questions

For ten 'quiz' multiple choice questions with the answers explained see MC test: Electron configuration.

For short-answer questions which can be set as an assignment for a test, homework or given for self study together with model answers see Electron configuration questions.

Vocabulary list

emission spectrum
discrete
continuous
s, p, d & f orbitals
Aufbau principle
Hund's rules
Pauli exclusion principle
Electron configuration
Heisenberg's uncertainty principle

IM, TOK, 'Utilization' etc.

See separate page which covers all of Topics 2

Practical work

Flame tests on metal ions using a nichrome wire. Note that you can also use a scent spray to spray the ions into the flame but do not use methanol as the solvent as there is a danger of a flash fire which can cause serious burns.

Teaching slides

Teachers may wish to share these slides with students for learning or for reviewing key concepts.

  

Other resources

1. A really nice video showing how all the orbitals build up around the nucleus to give the configuration for transition metals. It actually says it is for scandium but since no electrons are added it could be for any element between scandium and zinc.

  Arrangement of orbitals

2. Maybe instead you prefer your orbitals to be shown using balloons? If so take a look at this Brightstone video on atomic orbitals.

  Atomic orbitals

3. From a Nature of Science perspective it is well worth reading this interesting article on "What is an electron' by an astrophysicist.
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