Questions: Complete the partial orbital diagrams for the ions indicated. Fill orbitals from left to right and do not leave any boxes unfilled. Ion 4s 3d Cu^2+(Z=29)

Complete the partial orbital diagrams for the ions indicated. Fill orbitals from left to right and do not leave any boxes unfilled.

Ion 4s 3d
Cu^2+(Z=29)

Transcript text: Complete the partial orbital diagrams for the ions indicated. Fill orbitals from left to right and do not leave any boxes unfilled. Ion 4s 3d $\square$ $\mathrm{Cu}^{2+}(\mathrm{Z}=29)$

Solution

Solution Steps

Step 1: Cobalt(III) (Co³⁺)

Cobalt has an atomic number of 27, meaning a neutral cobalt atom has 27 electrons. Co³⁺ has lost 3 electrons, leaving it with 24. The electron configuration of Co is [Ar] 4s² 3d⁷. When forming the ion, electrons are first removed from the 4s orbital and then from the 3d orbitals. This leaves Co³⁺ with the configuration [Ar] 3d⁶. Following Hund's rule, the electrons in the 3d orbitals are filled individually before pairing.

Step 2: Copper(II) (Cu²⁺)

Copper has an atomic number of 29. Cu²⁺ has lost 2 electrons, resulting in 27 electrons. The electron configuration of Cu is [Ar] 4s¹ 3d¹⁰ (note the exception to the typical filling order). When forming Cu²⁺, the electrons are removed from the 4s orbital first, leaving Cu²⁺ with the configuration [Ar] 3d⁹.

Step 3: Selenide (Se²⁻)

Selenium has an atomic number of 34. Se²⁻ has gained 2 electrons, giving it a total of 36 electrons. The electron configuration of Se is [Ar] 4s² 3d¹⁰ 4p⁴. Adding two electrons fills the 4p orbital, resulting in the configuration [Ar] 4s² 3d¹⁰ 4p⁶ for Se²⁻.

Final Answer

Co³⁺: \[\boxed{4s: \quad \boxed{\phantom{\uparrow\downarrow}}\qquad 3d: \quad \boxed{\uparrow}\boxed{\uparrow}\boxed{\uparrow}\boxed{\uparrow}\boxed{\uparrow}\boxed{\phantom{\uparrow}}}\] Cu²⁺: \[\boxed{4s: \quad \boxed{\phantom{\uparrow\downarrow}}\qquad 3d: \quad \boxed{\uparrow\downarrow}\boxed{\uparrow\downarrow}\boxed{\uparrow\downarrow}\boxed{\uparrow\downarrow}\boxed{\uparrow}}\] Se²⁻: \[\boxed{4s: \quad \boxed{\uparrow\downarrow}\qquad 3d: \quad \boxed{\uparrow\downarrow}\boxed{\uparrow\downarrow}\boxed{\uparrow\downarrow}\boxed{\uparrow\downarrow}\boxed{\uparrow\downarrow}}\]

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