Structure 1.3—Electron configurations

Structure 1.3.1—Emission spectra are produced by atoms emitting photons when electrons in
excited states return to lower energy levels.

Structure 1.3.2—The line emission spectrum of hydrogen provides evidence for the existence of
electrons in discrete energy levels, which converge at higher energies

Structure 1.3.3—The main energy level is given an integer number, n, and can hold a maximum of
2n2 electrons

Structure 1.3.4—A more detailed model of the atom describes the division of the main energy level
into s, p, d and f sublevels of successively higher energies.

Structure 1.3.5—Each orbital has a defined energy state for a given electron configuration and
chemical environment, and can hold two electrons of opposite spin.
Sublevels contain a fixed number of orbitals, regions of space where there is a high probability of
finding an electron.

Structure 1.3.6—In an emission spectrum, the limit of convergence at higher frequency corresponds
to ionization. AHL

Structure 1.3.7—Successive ionization energy (IE) data for an element give information about its
electron configuration. AHL

What You’ll Learn:

  • Qualitatively describe the relationship between colour, wavelength, frequency, and energy across the electromagnetic spectrum.
  • Distinguish between a continuous 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.
  • The names of the different series in the hydrogen emission spectrum will not be assessed.
  • Deduce the maximum number of electrons that can occupy each energy level.
  • Recognize the shape and orientation of an s atomic orbital and the three p 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
  • Full electron configurations and condensed electron configurations using the noble gas core should be covered.
  • Orbital diagrams, i.e. arrow-in-box diagrams, should be used to represent the filling and relative energy of orbitals.
  • The electron configurations of Cr and Cu as exceptions should be covered.

AHL Only

  • Explain the trends and discontinuities in first ionization energy (IE) across a period and down a group.
  • Calculate the value of the first IE from spectral data that gives the wavelength or frequency of the convergence limit.
  • The value of the Planck constant h and the equations E = h f and c = λ f are given in the data booklet.
  • Deduce the group of an element from its successive ionization data.
  • Databases are useful for compiling graphs of trends in IEs


Syllabus Links

Inquiry 2—In the study of emission spectra from gaseous elements and of light, what qualitative and quantitative data can be collected from instruments such as gas discharge tubes and prisms?
Nature of science, Structure 1.2—How do emission spectra provide evidence for the existence of different elements?

Structure 3.1—How does an element’s highest main energy level relate to its period number in the periodic table?

Structure 3.1—What is the relationship between energy sublevels and the block nature of the periodic table?

Structure 3.1—How does the trend in IE values across a period and down a group explain the trends in properties of metals and non-metals?
Nature of science, Tool 3, Reactivity 3.1—Why are log scales useful when discussing [H+] and IEs?

HL Structure 3.1—How do patterns of successive IEs of transition elements help to explain the variable oxidation states of these elements?

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