Question

Write the electron configuration of the metal in each of the following: (a) \(\underline{V}\) (b) \(\mathrm{Co}^{2+}\) (c) \(\mathrm{MnO}_{2}\) (d) \(\mathrm{CuCl}_{4}^{2-}\)

Short answer

(a) V: 1s^2 2s^2 2p^6 3s^2 3p^6 4s^2 3d^3; (b) Co^2+: 1s^2 2s^2 2p^6 3s^2 3p^6 3d^7; (c) MnO_2: 1s^2 2s^2 2p^6 3s^2 3p^6 3d^3; (d) CuCl_4^{2-}: 1s^2 2s^2 2p^6 3s^2 3p^6 3d^9.

Step-by-step solution

Determine the Atomic Number

Identify the atomic number for each element. For vanadium (V) the atomic number is 23, for cobalt (Co) it is 27, for manganese (Mn) it is 25, and for copper (Cu) it is 29.

Write Ground State Electron Configuration

For each element, start by writing the electron configuration for the neutral atom:(a) Vanadium (V): The configuration is \[ 1s^2 \, 2s^2 \, 2p^6 \, 3s^2 \, 3p^6 \, 4s^2 \, 3d^3. \](b) Cobalt (Co): The configuration is \[ 1s^2 \, 2s^2 \, 2p^6 \, 3s^2 \, 3p^6 \, 4s^2 \, 3d^7. \](c & d) Manganese for MnOeffectively acts like a neutral Mn, similarly for copper in CuCl_4^{2-}, so consider electron configurations accordingly.

Adjust for Ion Charge (if applicable)

For ions, adjust the electron configuration by removing electrons from the highest energy orbitals first:(b) For \(\mathrm{Co}^{2+}\), remove 2 electrons from 4s and 3d orbitals. New configuration is \[ 1s^2 \, 2s^2 \, 2p^6 \, 3s^2 \, 3p^6 \, 3d^7. \](d) Copper in \(\mathrm{CuCl}_{4}^{2-}\) takes a similar approach: Copper normally has 4s^1 3d^{10}, but as Cu^{2+}, it becomes \[ 1s^2 \, 2s^2 \, 2p^6 \, 3s^2 \, 3p^6 \, 3d^9. \]

Consider Compounds Electron Configuration

When in complex like compounds, verify the oxidation states:(c) Mn in \(\mathrm{MnO}_{2}\) is +4, implying 3d^3 configuration due to loss of 2 4s electrons and 2 3d electrons: \[ 1s^2 \, 2s^2 \, 2p^6 \, 3s^2 \, 3p^6 \, 3d^3. \](d) For \(\mathrm{CuCl}_{4}^{2-}\), the influence as a complex might stabilize its configuration, as previously mentioned.

Key concepts

Transition Metals

Transition metals are a group of elements found in the central block of the periodic table. These elements include the d-block elements, which are located in groups 3 through 12.
Transition metals are characterized by having partially filled d orbitals, which means they can exhibit a wide range of oxidation states.Some notable properties of transition metals include:

• They often form colored compounds. For example, copper compounds are frequently blue or green.
• They are known for their ability to form stable complexes. A common example is copper's ability to coordinate with chlorine in the complex \(\mathrm{CuCl}_{4}^{2-}\).
• These metals can also display catalytic properties, which makes them valuable in industrial processes.

In terms of electron configuration, transition metals fill their 3d sublevel after filling the 4s sublevel. This unique electron arrangement allows for the diverse chemical behavior observed in these elements.

Oxidation States

Oxidation states refer to the effective charge of an atom in a compound. This concept is essential for understanding how atoms share or transfer electrons in chemical reactions.
For transition metals, the trick lies in their multiple oxidation states due to electrons in both s and d orbitals being available for bonding.Key points about oxidation states in transition metals include:

• Transition metals can lose varying numbers of electrons, giving rise to different oxidation states. For example, manganese in \(\mathrm{MnO}_{2}\) has an oxidation state of +4.
• The oxidation state is important for determining electron configurations. For example, cobalt \(\mathrm{Co}^{2+}\) loses two electrons resulting in the electron configuration \( \big[ \mathrm{Ar} \big] 3d^7 \).
• Understanding oxidation states helps predict and explain the reactivity and formation of specific compounds.

For instance, in \(\mathrm{CuCl}_{4}^{2-}\), copper usually exhibits a +2 oxidation state, which shows how it loses electrons to form stable ionic compounds.

Quantum Numbers

Quantum numbers are used to describe the unique quantum state of an electron in an atom. Each electron in an atom is described by four quantum numbers.

• The principal quantum number \(n\) signifies the electron's energy level or shell.
• The azimuthal quantum number \(l\) describes the subshell or orbital type (s, p, d, f).
• The magnetic quantum number \(m_l\) indicates the orientation of the orbital in space.
• The spin quantum number \(m_s\) describes the electron's spin direction, either +1/2 or -1/2.

Quantum numbers become especially helpful in determining electron configurations. For instance, the electron configuration of vanadium \( \big[ \mathrm{Ar} \big] 3d^3 \) can be broken down using quantum numbers to show specific electron positions within orbitals:
- Electrons in the 3d orbital of vanadium have \(n = 3\), \(l = 2\).- Each have different \(m_l\) values ranging between -2 to +2.Mastering these concepts allows for a better understanding of chemical bonding, electron arrangement, and properties of elements, particularly transition metals.