Electronic configurations using spectroscopic notation

Using spectroscopic notation, write down the electronic configurations of the following species (note that some are ions) ... 

Consider the electronic configuration of carbon again Is 2s 2pl Remember, there are three different p orbitals in the 2p subshell the p orbital lies on the x-axis the p orbital lies on the y-axis and the p orbital lies on the z-axis. The different p orbitals are degenerate. To obey Hund s rule, these degenerate orbitals must be filled singly before spin pairing occurs. To obey the Pauli exclusion principle, when an orbital is full with two electrons, these electrons must have opposite spins. This is not shown using spectroscopic notation, but is seen when orbital box notation is used. 

The table shows the electronic configuration in spectroscopic and orbital box notation for the elements from scandium to zinc. [Ar] represents the electronic configuration of argon, which is Is 2s 2p 3s 3p . It is okay to use this shorthand here instead of writing out the full electron shells up to 3p. However, in the exam you should write out the spectroscopic notation for each element in full. 

The d block transition metals are metals with an incomplete d subshell in at least one of their ions. Try to explain why Sc and Zn are often considered not to be transition metals. Consider the electronic configurations of the Fe + and Fe ions in both spectroscopic and orbital box notations. Use these notations to explain why Fe(lll) compounds are more stable than Fe(ll) compounds. 

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