One common misunderstanding among Year 10 students about electron configuration is the belief that all electrons fill the outer shell first. This isn’t true! While it seems natural to think that the outer shells get filled before the inner shells, electrons actually fill different energy levels according to something called the Aufbau principle.

Another frequent mistake is mixing up the order of filling subshells. Many students accidentally switch the $3d$ and $4s$ subshells. The correct order is: $1s$, $2s$, $2p$, $3s$, $3p$, $4s$, $3d$, and then $4p$. It’s important to remember that the $4s$ subshell fills up before the $3d$ because it has less energy.

Students can also forget about how electrons pair up. When electrons start filling the same orbital, they tend to pair together. This can cause confusion about how many electrons can fit in a shell or subshell. For example, $p$ orbitals can hold up to 6 electrons, but that’s only true when they are spread out across the three available orbitals before any pairing happens.

Lastly, students might think every element has the same pattern for electron configuration, but that’s not correct. Each element has its own special electron configuration based on its atomic number. This uniqueness helps determine the properties and behaviors of the elements.

Understanding these differences is really important for getting a better grasp of chemistry.

By clearing up these misunderstandings, students can strengthen their knowledge of atomic structure and electron configuration!