Question

1. ne 10 10. fe 26 create box diagrams for each of the following elements/ions. only use your periodic table and your knowledge of the three rules to complete this task. 1. ne 2. na⁺¹ 3. al 4. p 5. fe

Explanation:

1. Ne (Neon)

• Step 1: Determine electron configuration

Neon has an atomic number of 10, so its electron configuration is \(1s^2 2s^2 2p^6\).

• Step 2: Draw box diagram for each subshell
• \(1s\): 1 box, 2 electrons (paired): \(\boxed{\uparrow\downarrow}\)
• \(2s\): 1 box, 2 electrons (paired): \(\boxed{\uparrow\downarrow}\)
• \(2p\): 3 boxes, each with 2 electrons (all paired): \(\boxed{\uparrow\downarrow}\ \boxed{\uparrow\downarrow}\ \boxed{\uparrow\downarrow}\)

2. \(\boldsymbol{Na^{+1}}\) (Sodium ion)

• Step 1: Determine electron configuration

Sodium (\(Na\)) has atomic number 11 (\(1s^2 2s^2 2p^6 3s^1\)). The \(Na^+\) ion loses 1 electron (from \(3s\)), so its configuration is \(1s^2 2s^2 2p^6\) (same as Ne).

• Step 2: Draw box diagram
• \(1s\): \(\boxed{\uparrow\downarrow}\)
• \(2s\): \(\boxed{\uparrow\downarrow}\)
• \(2p\): \(\boxed{\uparrow\downarrow}\ \boxed{\uparrow\downarrow}\ \boxed{\uparrow\downarrow}\)

3. Al (Aluminum)

• Step 1: Determine electron configuration

Aluminum has atomic number 13: \(1s^2 2s^2 2p^6 3s^2 3p^1\).

• Step 2: Draw box diagram
• \(1s\): \(\boxed{\uparrow\downarrow}\)
• \(2s\): \(\boxed{\uparrow\downarrow}\)
• \(2p\): \(\boxed{\uparrow\downarrow}\ \boxed{\uparrow\downarrow}\ \boxed{\uparrow\downarrow}\)
• \(3s\): \(\boxed{\uparrow\downarrow}\)
• \(3p\): 3 boxes, 1 electron (unpaired in first box): \(\boxed{\uparrow}\ \boxed{}\ \boxed{}\)

4. P (Phosphorus)

Answer:

• Step 1: Determine electron configuration

Iron has atomic number 26. The electron configuration is \(1s^2 2s^2 2p^6 3s^2 3p^6 4s^2 3d^6\) (note: \(4s\) fills before \(3d\), then \(3d\) fills).

• Step 2: Draw box diagram
• \(1s\): \(\boxed{\uparrow\downarrow}\)
• \(2s\): \(\boxed{\uparrow\downarrow}\)
• \(2p\): \(\boxed{\uparrow\downarrow}\ \boxed{\uparrow\downarrow}\ \boxed{\uparrow\downarrow}\)
• \(3s\): \(\boxed{\uparrow\downarrow}\)
• \(3p\): \(\boxed{\uparrow\downarrow}\ \boxed{\uparrow\downarrow}\ \boxed{\uparrow\downarrow}\)
• \(4s\): \(\boxed{\uparrow\downarrow}\)
• \(3d\): 5 boxes. Fill 6 electrons (Hund’s rule: 5 unpaired first, then pair one):

\(\boxed{\uparrow\downarrow}\ \boxed{\uparrow}\ \boxed{\uparrow}\ \boxed{\uparrow}\ \boxed{\uparrow}\)

Key Notes:

• Pauli Exclusion Principle: Each orbital (box) holds 2 electrons with opposite spins (\(\uparrow\downarrow\)).
• Aufbau Principle: Electrons fill lower-energy orbitals first (\(1s \to 2s \to 2p \to 3s \to 3p \to 4s \to 3d \to \dots\)).
• Hund’s Rule: Electrons fill degenerate (same-energy) orbitals singly before pairing.