Atomic structure flashcards

Atomic structure underpins all of chemistry. Students must understand that atoms consist of a tiny, dense nucleus containing protons and neutrons, surrounded by electrons arranged in energy levels and sublevels. The atomic number defines the element, while the mass number gives the total of protons and neutrons. Isotopes are atoms of the same element with different numbers of neutrons, and relative atomic mass is the weighted average mass of all naturally occurring isotopes. Electrons occupy orbitals within sublevels (s, p, d, f), and the arrangement of electrons, called the electron configuration, determines chemical behaviour. Evidence for discrete energy levels comes from atomic emission spectra, where electrons release photons of specific frequencies when falling to lower energy levels. Understanding ionisation energy trends also reinforces the shell and sublevel model.

1. Card 1

   What is the atomic number of an element?

   The atomic number is the number of protons in the nucleus of an atom. It uniquely identifies the element and equals the number of electrons in a neutral atom.

2. Card 2

   How is mass number defined and calculated?

   The mass number is the total number of protons and neutrons in the nucleus. It is calculated as: mass number = protons + neutrons.

3. Card 3

   What are isotopes?

   Isotopes are atoms of the same element that have the same number of protons but different numbers of neutrons, giving them different mass numbers.

4. Card 4

   How is relative atomic mass calculated from isotope data?

   Relative atomic mass is the weighted average mass of all naturally occurring isotopes relative to 1/12 the mass of carbon-12, calculated using isotope masses and their percentage abundances.

5. Card 5

   What is the difference between a shell and a sublevel?

   A shell is a main energy level numbered 1, 2, 3 etc. Each shell contains one or more sublevels (s, p, d, f) of slightly different energies.

6. Card 6

   How many electrons can s, p, and d sublevels each hold?

   An s sublevel holds up to 2 electrons, a p sublevel holds up to 6, and a d sublevel holds up to 10, because they contain 1, 3, and 5 orbitals respectively.

7. Card 7

   What is an atomic orbital?

   An atomic orbital is a region of space around the nucleus where there is a high probability of finding an electron. Each orbital holds a maximum of two electrons.

8. Card 8

   Write the full electron configuration of iron (Fe, Z = 26).

   The electron configuration of iron is 1s2 2s2 2p6 3s2 3p6 3d6 4s2, showing how 26 electrons are distributed across sublevels in order of increasing energy.

9. Card 9

   What is the first ionisation energy and what trend does it show across period 3?

   First ionisation energy is the energy to remove one mole of electrons from one mole of gaseous atoms. Across period 3 it generally increases due to greater nuclear charge with similar shielding.

10. Card 10

    Why does atomic emission spectroscopy provide evidence for discrete energy levels?

    Atomic emission spectra show distinct lines of specific frequencies, proving electrons can only occupy fixed energy levels and release photons of precise energy when dropping between them.

11. Card 11

    What causes the convergence of lines in the hydrogen emission spectrum?

    Lines converge at higher frequencies because the energy levels become closer together as they approach the ionisation limit, where the electron is completely removed from the atom.

12. Card 12

    Why do transition metals often lose 4s electrons before 3d electrons?

    Once the 3d sublevel is occupied, it becomes lower in energy than 4s, so 4s electrons are removed first during ionisation, explaining the electron configurations of transition metal ions.

• Students often confuse mass number with relative atomic mass, forgetting that relative atomic mass is a weighted average and is rarely a whole number.
• A common error is writing the electron configuration of transition metal ions by removing 3d electrons first rather than 4s electrons, which are actually at higher energy once the 3d sublevel is occupied.
• Many students incorrectly state that electrons travel in fixed circular orbits rather than occupying orbitals, which are probability regions with no defined path.

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