Question

Which element in the transition-metal series Sc through \(\mathrm{Zn}\) has five unpaired electrons in its \(+3\) ion?

Short answer

No transition metal among Sc-Zn forms "+3" ions with 5 unpaired electrons.

Step-by-step solution

Identify the Configuration of Transition Metals

Transition metals, from scandium (Sc) to zinc (Zn), fill their d-orbitals according to the 3d energy level. Remember that the atomic number increases from Sc (21) to Zn (30), filling the 3d orbitals with varying numbers of electrons up to 10 because each orbital can hold 2 electrons.

Understand the Ion Formation

When a transition metal forms a +3 ion, it loses three electrons, typically starting from the 4s orbital and then the 3d orbitals if necessary. This is because the 4s orbital is filled before the 3d orbitals but is higher in energy once filled.

Calculate Electron Configuration for +3 Ion

Determine which element in its neutral state will result in exactly five unpaired electrons after losing three electrons. Generally, you fill the d-orbital, and losing electrons from it will impact on the electron pairing unfavorably.

Determine the Element Required

Manganese (Mn) has an atomic number of 25. Its neutral electron configuration is \[ \text{[Ar]} 4s^2 3d^5 \]. When it forms a +3 ion, it loses 2 electrons from 4s and 1 from 3d resulting in \[ 3d^4 \]. In this configuration, there are four unpaired electrons, not five. Look for configuration \[ 3d^5 \].

Verification

Chromium (Cr) with an atomic number 24, has a neutral-state configuration of \[ \text{[Ar]} 4s^1 3d^5 \]. When it loses three electrons forming Cr^3+, it results in \[ 3d^3 \] instead of \[ 3d^2 \], suggesting possible confusion. Recalculate V^(+3) due to similar numbers in its configuration to check errors. \[ \text{[Ar]} 4s^2 3d^3 \] for vanadium with similar calculation.

Confirm the Correct Element

Double-check calculations for each element within this range for \[Mn^+3 \]. Correct mistake, exactly determine that manganese (Mn) is the only element having 5 unpaired electrons in a neutral state to look at the implications in loss. Special case noted with electron configuration loss corrected.

Key concepts

Electron Configuration

Transition metals, situated in the middle of the periodic table, exhibit some unique behaviors due to their electron configuration. These elements usually involve the filling of 3d and 4s orbitals. Electrons are added to 3d orbitals following the Aufbau principle, which states that electrons occupy the lowest energy orbitals first.
In general, the typical electron configuration for these metals can be written as

• Before filling the 3d orbital, the 4s orbital is usually filled first, as illustrated by the configuration
• The electron configuration of chronium (Cr): \[ [Ar] 4s^1 3d^5 \]
  

When these elements form ions, electron loss occurs typically from 4s before 3d due to their positioning in energy.
Understanding electron configuration helps predict the chemical properties and reactivity of transition metals.

Unpaired Electrons

Unpaired electrons greatly influence an atom’s magnetic properties and reactivity. Transition metals often have unpaired electrons due to their d-orbitals not being completely filled. When a metal loses electrons to form a cation, the configuration can change, impacting electron pairing.
For instance, neutral manganese (Mn) has the configuration \[ [Ar] 4s^2 3d^5 \], boasting five unpaired electrons in its neutral state. However, when it forms a \( 3+ \) cation and loses electrons, it's configuration changes to \[ 3d^4 \], resulting in four unpaired electrons .Identifying elements with specific numbers of unpaired electrons is crucial for understanding their oxidation states and magnetic behavior, such as paramagnetism, which arises when substances have unpaired electrons that align with an external magnetic field.

Oxidation States

Transition metals can exhibit various oxidation states, which differ by the number of electrons lost or shared during reactions. This flexibility is due to their similar energy levels in the 3d and 4s orbitals. The most common oxidation states for a metal often involve the loss of its 4s and, occasionally, 3d electrons.
For instance, manganese (Mn) can exhibit multiple oxidation states, with \( +2, +3, +4, \) and \( +7 \) being quite common .
These oxidation states are essential for comprehending the role of transition metals in redox reactions, catalysis, and coordination chemistry. Knowing the possible oxidation states helps anticipate how a metal might react with different ligands or elements.

d-Orbitals

Understanding d-orbitals is crucial when studying transition metals. These orbitals have distinct shapes and spatial orientations, which significantly impact the physical and chemical properties of an element. The 3d orbitals set transition metals apart due to their ability to host electrons even after the 4s orbital is filled.
The arrangement within these d-orbitals determines aspects such as electron configuration and oxidation states.

• They influence the possible states in which a metal can exist due to similar energy levels of 3d and 4s orbitals
• d-Orbitals' involvement in bonding and their role in absorbing visible light lead to the typically vibrant coloration of transition metals.

Examining d-orbitals gives insight into concepts like ligand field theory, which explains color and magnetism based on electron transitions within d-orbitals.