# 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. Commented Feb 18, 2015 at 4:22

• Electronic configuration of molybdenum: $$[\ce{Kr}] \ce{4d^{5} 5s^1}$$, instead of $$[\ce{Kr}] \ce{5s^2 4d^{4}}$$ 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 $$[\ce{Kr}]\ce{ 5s^2 4d^{8}}$$ 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.