How many unpaired electron(s) are there in the nitrogen sub-levels?

1. Analyze the question

The question asks for the number of unpaired electrons in a neutral nitrogen atom. To find this, we need to determine the electronic configuration of nitrogen and apply Hund's Rule of Maximum Multiplicity.

2. Explain the concept

• Atomic Number: Nitrogen (N) has an atomic number of 7. This means a neutral nitrogen atom has 7 electrons.
• Electronic Configuration: Electrons fill orbitals in order of increasing energy: $1s^2, 2s^2, 2p^3$.
• Hund’s Rule: This rule states that for a given electron sub-shell (like the $p$-subshell), electrons fill empty orbitals singly before they start pairing up. The $p$-subshell has three orbitals ($p_x, p_y, p_z$).

3. Show the solution

• Step 1: Write the full electronic configuration of Nitrogen ($Z=7$): $1s^2, 2s^2, 2p^3$
• Step 2: Focus on the outermost sub-level, which is $2p^3$.
• Step 3: Distribute the 3 electrons in the $2p$ sub-shell into the three available $p$-orbitals ($p_x, p_y, p_z$): • $p_x$ gets 1 electron • $p_y$ gets 1 electron • $p_z$ gets 1 electron
• Step 4: Count the unpaired electrons. Since each of the three $p$-orbitals contains exactly one electron, there are 3 unpaired electrons.

4. Identify the answer

Based on the configuration, there are 3 unpaired electrons in the $2p$ sub-level. This matches option A.

5. Tip

Memory Trick: Remember the "Bus Seat Rule." Electrons are like people getting on a bus; they prefer to sit in empty seats (orbitals) individually before they are forced to share a seat (pair up). Since nitrogen has 3 electrons in the $p$-subshell and there are 3 "seats" available, each electron gets its own seat, leaving 3 unpaired!