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Structure 1.3 Flashcards

Active Recall for Electron Configurations

1.3.1 Shells, Subshells & Orbitals: read all 6 questions and answers
  1. What is the maximum number of electrons in each main energy level (n)?
    2n². So: n=1 holds 2, n=2 holds 8, n=3 holds 18, n=4 holds 32.
  2. What are the four types of subshell?
    s (holds 2e), p (holds 6e), d (holds 10e), f (holds 14e).
  3. What is an orbital?
    A region of space around the nucleus where there is a high probability of finding an electron. Each orbital holds a maximum of 2 electrons with opposite spins.
  4. How many orbitals in each subshell type?
    s: 1 orbital, p: 3 orbitals, d: 5 orbitals, f: 7 orbitals.
  5. Describe the shape of an s orbital.
    Spherical, centred on the nucleus.
  6. Describe the shape of a p orbital.
    Dumbbell-shaped (two lobes). The three p orbitals are oriented along x, y, and z axes at 90 degrees to each other.
1.3.2 Aufbau, Hund & Pauli: read all 6 questions and answers
  1. State the Aufbau Principle.
    Electrons fill the lowest available energy orbital first. Order: 1s, 2s, 2p, 3s, 3p, 4s, 3d, 4p...
  2. State Hund's Rule.
    Orbitals of equal energy (degenerate) are each singly occupied before any orbital is doubly occupied, and all singly occupied orbitals have the same spin.
  3. State the Pauli Exclusion Principle.
    No two electrons in an atom can have the same set of four quantum numbers. In practice, a maximum of 2 electrons per orbital with opposite spins.
  4. Why is the electron configuration of Chromium anomalous?
    [Ar] 3d5 4s1 (not [Ar] 3d4 4s2). A half-filled d subshell provides extra stability due to symmetrical electron distribution.
  5. Why is the electron configuration of Copper anomalous?
    [Ar] 3d10 4s1 (not [Ar] 3d9 4s2). A fully filled d subshell provides extra stability.
  6. Why does 4s fill before 3d?
    In neutral atoms, the 4s orbital has slightly lower energy than 3d. However, once the atom is ionised, 3d becomes lower in energy than 4s.
1.3.3 Ion Configurations: read all 5 questions and answers
  1. Which electrons are removed first when forming transition metal cations?
    The 4s electrons are removed before 3d electrons. E.g. Fe2+ is [Ar] 3d6 (not [Ar] 3d4 4s2).
  2. Write the electron configuration of Fe2+.
    [Ar] 3d6. Iron loses its two 4s electrons first.
  3. Write the electron configuration of Fe3+.
    [Ar] 3d5. Iron loses two 4s electrons and one 3d electron.
  4. How do negative ions form?
    Atoms gain electrons to achieve a full outer shell. E.g. O2- is 1s2 2s2 2p6 (same as Ne).
  5. What are isoelectronic species?
    Species with the same electron configuration. E.g. Na+, Mg2+, F-, O2-, and Ne all have the configuration 1s2 2s2 2p6.
1.3.4 Ionisation Energy Trends: read all 6 questions and answers
  1. Define first ionisation energy.
    The energy required to remove one mole of electrons from one mole of gaseous atoms to form one mole of gaseous 1+ ions.
  2. How does IE trend across a period?
    Generally increases across a period due to increasing nuclear charge with the same shielding, so greater attraction for outer electrons.
  3. How does IE trend down a group?
    Decreases down a group because outer electrons are further from the nucleus with more shielding, so less attraction.
  4. Why does IE drop between Groups 2 and 13?
    Group 13 elements lose a p electron, which is higher in energy and easier to remove than the s electron lost by Group 2.
  5. Why does IE drop between Groups 15 and 16?
    In Group 16, the electron being removed is paired in a p orbital. Electron-electron repulsion in the pair makes it easier to remove.
  6. How do successive ionisation energies provide evidence for electron shells?
    Large jumps in IE occur when electrons are removed from a new (inner) shell, indicating distinct energy levels.
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