A-Level Chemistry Flashcards: Atomic Structure

A-Level Chemistry: Atomic Structure — Flashcard Deck

20 flashcards with spaced repetition. Press Space to flip, then rate your recall (1—4).

What These Flashcards Cover

These flashcards test your knowledge of atomic structure, electron configurations, and periodic trends. You should be able to write electron configurations, interpret ionisation energy data, and explain patterns in the periodic table.

Key areas:

• Electron Configuration: Fill orbitals in order: $1s$, $2s$, $2p$, $3s$, $3p$, $4s$, $3d$, $4p$, $5s$, $4d$, $5p$, $6s$, $4f$, $5d$, $6p$. Each $s$ orbital holds 2 electrons, $p$ holds 6, $d$ holds 10, $f$ holds 14. Use the noble gas shorthand (e.g. Fe: $[\text{Ar}]3d^{6}4s^{2}$).
• Sub-shells and Orbitals: An orbital is a region of space with a high probability of finding an electron. $s$ orbitals are spherical; $p$ orbitals are dumbbell-shaped (three orientations: $p_x$, $p_y$, $p_z$). $d$ orbitals have more complex shapes (five orientations).
• Ionisation Energy: First ionisation energy is the energy to remove one electron from each atom in one mole of gaseous atoms: $X(g) \rightarrow X^{+}(g) + e^{-}$. Successive ionisation energies increase. Large jumps indicate a new electron shell.
• Periodic Trends: Ionisation energy generally increases across a period (increasing nuclear charge, similar shielding) and decreases down a group (increasing shielding and atomic radius). Anomalies in period 2: Be > B (full $2s$ sub-shell) and N > O (half-filled $2p$ sub-shell).
• Atomic and Ionic Radii: Atomic radius decreases across a period (stronger nuclear attraction) and increases down a group (more shells). Cations are smaller than their parent atoms; anions are larger.
• Quantum Numbers: Principal quantum number $n$ (shell), angular momentum quantum number $l$ (sub-shell: $0 = s$, $1 = p$, $2 = d$, $3 = f$), magnetic quantum number $m_l$ (orbital orientation: $-l$ to $+l$), spin quantum number $m_s$ ($+\frac{1}{2}$ or $-\frac{1}{2}$). No two electrons can have the same set of four quantum numbers (Pauli exclusion principle).

Intuition

Think of the atom as a building. The shells ($n$) are floors, the sub-shells ($l$) are rooms on each floor, and the orbitals ($m_l$) are individual desks within each room. Each desk can hold at most two electrons (with opposite spins). Electrons fill the lowest floors first, like people filling a building from the ground up.

Ionisation energy is like the cost of evicting an electron. Electrons closer to the nucleus (lower $n$) or in filled sub-shells are harder to remove, so more energy is required.

Common Pitfalls

1. Filling order mistakes. The $4s$ orbital fills before $3d$, but when writing the configuration for transition metals, the $3d$ electrons are sometimes written before $4s$ (by convention). Be careful about which order the question expects.
2. Confusing ionisation energy with electron affinity. Ionisation energy is the energy to remove an electron. Electron affinity is the energy change when an electron is added. These are different processes with different signs.
3. Miscounting electrons in ions. For transition metal ions, electrons are removed from the $4s$ orbital before the $3d$. Fe$^{2+}$ is $[\text{Ar}]3d^{6}$, not $[\text{Ar}]3d^{4}4s^{2}$.
4. Forgetting that the $p$ sub-shell has three orbitals. This means it can hold a maximum of 6 electrons, not 3. Similarly, $d$ holds 10 and $f$ holds 14.

Cross-References

• Physical Chemistry: Atomic structure underpins physical chemistry concepts
• Organic Chemistry: Carbon’s atomic structure enables organic chemistry
• Equilibrium: Atomic structure influences chemical equilibrium positions