# Why does the electron configuration for some elements not follow the diagonal rule?

I'm doing a high-school assignment, and I came across a question that I didn't quite understand.

Explain how the electron configurations for the following elements do not follow the diagonal rule:

• Gold
• Curium
• Thorium
• Molybdenum

From what I understand, all of these elements do follow the diagonal rule. How exactly do they not follow this rule?

• Look them up—they're not what you think. For example, gold is [Xe] 4f14 5d10 6s1, not [Xe] 4f14 5d9 6s2. – Michael DM Dryden Feb 18 '15 at 4:22

• Electronic configuration of molybdenum: $[\ce{Kr}] \ce{4d^{5} 5s^1}$, instead of $[\ce{Kr}] \ce{5s^24d^{4}}$ according to the diagonal rule, because a half-full $\ce{4d}$ subshell and a half full $\ce{5s}$ subshell are more stable than $\ce{4d}$ filled with four electrons and a full $\ce{5s}$ subshell.
• Electronic configuration of Gold: $[\ce{Xe}] \ce{4f^14 5d^{10} 6s^1}$, instead of $[\ce{Xe}] \ce{4f^14 5d^{9} 6s^2}$ according to the diagonal rule, because a full $\ce{5d}$ and half full $\ce{6s}$ subshell is more stable than $\ce{5d}$ filled with 9 electrons and a full $\ce{6s}$ subshell.
• Electronic configuration of Palladium: $[\ce{Kr}] \ce{4d^{10}}$, instead of $[\ce{Kr}]\ce{ 5s^24d^{8}}$ according to the diagonal rule, because a full $\ce{4d}$ orbital is more stable than $\ce{4d}$ filled with eight electrons.
• Electronic configuration of Curium: $[\ce{Rn}] \ce{7s^2 5f^7 6d^{1}}$, instead of $[\ce{Rn}] \ce{7s^2 5f^8}$ according to the diagonal rule, because a half full $\ce{5f}$ orbital is more stable than $\ce{5f}$ filled with eight electrons.