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Electron configuration describes how electrons are arranged around an atom in shells, subshells, and orbitals. This arrangement helps explain an element's chemical behavior, including bonding, reactivity, and periodic trends. By learning the order in which electrons fill available energy levels, students can predict properties of atoms and ions. It is one of the key links between atomic structure and the periodic table.
Electrons occupy orbitals according to the Aufbau principle, which states that lower energy orbitals fill before higher energy ones. This process also follows the Pauli exclusion principle, which limits each orbital to two electrons with opposite spins, and Hund's rule, which favors single occupancy in equal energy orbitals before pairing. Shells are labeled by principal energy level n, while subshells are labeled s, p, d, and f. Together, these ideas create a map for writing configurations such as 1s2 2s2 2p6 and for understanding why some elements have similar chemistry.
Key Facts
- Maximum electrons in shell n: 2n^2
- Subshell capacities are s = 2, p = 6, d = 10, f = 14
- Orbital filling order commonly follows: 1s, 2s, 2p, 3s, 3p, 4s, 3d, 4p, 5s, 4d, 5p, 6s, 4f, 5d, 6p, 7s
- Each orbital holds at most 2 electrons with opposite spins
- Number of orbitals in each subshell: s = 1, p = 3, d = 5, f = 7
- For a neutral atom, total electrons = atomic number
Vocabulary
- Electron configuration
- The arrangement of an atom's electrons among its shells, subshells, and orbitals.
- Shell
- A main energy level of an atom identified by the principal quantum number n.
- Subshell
- A division of a shell labeled s, p, d, or f that contains orbitals of similar energy.
- Orbital
- A region around the nucleus where there is a high probability of finding an electron.
- Aufbau principle
- The rule that electrons fill the lowest available energy orbitals before occupying higher ones.
Common Mistakes to Avoid
- Putting electrons into higher energy orbitals too early, which breaks the Aufbau principle and gives the wrong configuration. Always fill the lowest available orbital first according to the standard order.
- Forgetting that a p subshell has three orbitals, which leads students to pair electrons too soon. Apply Hund's rule by placing one electron in each equal energy orbital before pairing.
- Assuming shells and subshells are the same thing, which causes confusion about labels like 3d and 4s. A shell is the main energy level, while a subshell is a part of that shell.
- Using the atomic mass instead of the atomic number to count electrons, which gives the wrong total for a neutral atom. The number of electrons in a neutral atom equals the atomic number.
Practice Questions
- 1 Write the full electron configuration for oxygen, which has atomic number 8.
- 2 How many electrons can fit in the n = 3 shell, and how many electrons can fit in the 3d subshell?
- 3 Explain why nitrogen has three unpaired electrons in its 2p subshell instead of one paired orbital and one single electron.