Which Of The Following Orbital Diagrams Represents A Paramagnetic Atom

Determining whether an atom is paramagnetic based on its orbital diagram is a fundamental concept in chemistry. Paramagnetism arises from the presence of unpaired electrons within an atom. Understanding how to interpret orbital diagrams allows us to predict the magnetic properties of elements and their ions. This article delves into the principles of paramagnetism, explains how to read orbital diagrams, and provides examples to illustrate which orbital diagrams represent paramagnetic atoms.

Introduction to Paramagnetism

Paramagnetism is a form of magnetism that occurs only in the presence of an externally applied magnetic field. Paramagnetic materials are attracted to magnetic fields, but unlike ferromagnetic materials, they do not retain any permanent magnetism when the field is removed. This behavior is due to the presence of unpaired electrons, which possess a magnetic dipole moment.

Magnetic Dipole Moment: Each electron has an intrinsic property called spin, which creates a magnetic dipole moment. In atoms with paired electrons, these moments cancel each other out because one electron has a spin-up (+1/2) and the other has a spin-down (-1/2).
Unpaired Electrons: When an atom has unpaired electrons, the magnetic dipole moments do not cancel, resulting in a net magnetic moment. These unpaired electrons align with an external magnetic field, causing the atom to be attracted to the field.

Understanding Orbital Diagrams

Orbital diagrams, also known as electronic configurations or energy-level diagrams, visually represent the arrangement of electrons within the orbitals of an atom. These diagrams follow specific rules to ensure accurate representation:

Orbitals and Sublevels: Orbitals are regions around the nucleus where electrons are likely to be found. They are grouped into sublevels (s, p, d, and f), each with a specific number of orbitals:
- s sublevel: 1 orbital
- p sublevel: 3 orbitals
- d sublevel: 5 orbitals
- f sublevel: 7 orbitals
Pauli Exclusion Principle: This principle states that no two electrons in an atom can have the same set of four quantum numbers. In simpler terms, each orbital can hold a maximum of two electrons, and they must have opposite spins (one spin-up and one spin-down).
Hund's Rule: This rule states that electrons will individually occupy each orbital within a sublevel before doubling up in any one orbital. Additionally, electrons in singly occupied orbitals will have the same spin (maximize total spin).
Aufbau Principle: This principle dictates the order in which electrons fill the orbitals, starting with the lowest energy levels. The filling order is typically: 1s, 2s, 2p, 3s, 3p, 4s, 3d, 4p, 5s, 4d, 5p, 6s, 4f, 5d, 6p, 7s, 5f, 6d, 7p.

Identifying Paramagnetic Atoms Using Orbital Diagrams

To determine whether an atom is paramagnetic based on its orbital diagram, follow these steps:

Draw the Orbital Diagram: Based on the atom's electron configuration, draw the orbital diagram showing how electrons are distributed among the orbitals.
Apply Hund's Rule: Fill the orbitals in each sublevel individually before pairing electrons in the same orbital. Ensure that unpaired electrons have the same spin.
Count Unpaired Electrons: Count the number of unpaired electrons in the orbital diagram. If there are one or more unpaired electrons, the atom is paramagnetic. If all electrons are paired, the atom is diamagnetic.

Examples of Paramagnetic Atoms

Let's examine several examples to illustrate how to identify paramagnetic atoms using orbital diagrams:

Example 1: Nitrogen (N)

Nitrogen has an atomic number of 7, meaning it has 7 electrons. Its electron configuration is 1s² 2s² 2p³.

Orbital Diagram:
- 1s: ↑↓
- 2s: ↑↓
- 2p: ↑ ↑ ↑
Analysis: The 1s and 2s orbitals are fully occupied with paired electrons. The 2p sublevel has three orbitals, each occupied by a single electron with the same spin.
Conclusion: Nitrogen has three unpaired electrons, making it a paramagnetic atom.

Example 2: Oxygen (O)

Oxygen has an atomic number of 8, meaning it has 8 electrons. Its electron configuration is 1s² 2s² 2p⁴.

Orbital Diagram:
- 1s: ↑↓
- 2s: ↑↓
- 2p: ↑↓ ↑ ↑
Analysis: The 1s and 2s orbitals are fully occupied with paired electrons. The 2p sublevel has four electrons, with two orbitals having paired electrons and two orbitals having single electrons with the same spin.
Conclusion: Oxygen has two unpaired electrons, making it a paramagnetic atom.

Example 3: Iron (Fe)

Iron has an atomic number of 26, meaning it has 26 electrons. Its electron configuration is 1s² 2s² 2p⁶ 3s² 3p⁶ 4s² 3d⁶.

Orbital Diagram (focus on 3d):
- 3d: ↑↓ ↑ ↑ ↑ ↑
Analysis: The 3d sublevel has six electrons. According to Hund's rule, five orbitals are first filled with single electrons, and then the sixth electron pairs with one of them.
Conclusion: Iron has four unpaired electrons, making it a paramagnetic atom.

Example 4: Nickel (Ni)

Nickel has an atomic number of 28, meaning it has 28 electrons. Its electron configuration is 1s² 2s² 2p⁶ 3s² 3p⁶ 4s² 3d⁸.

Orbital Diagram (focus on 3d):
- 3d: ↑↓ ↑↓ ↑ ↑ ↑
Analysis: The 3d sublevel has eight electrons. Following Hund's rule, each of the five orbitals is first occupied singly, and then the remaining three electrons pair up in three of the orbitals.
Conclusion: Nickel has two unpaired electrons, making it a paramagnetic atom.

Exceptions to Hund's Rule and Paramagnetism

While Hund's rule provides a good guideline for predicting electron configurations and paramagnetism, there are exceptions, particularly among transition metals. These exceptions arise due to the stability associated with having half-filled or fully filled d sublevels.

Chromium (Cr)

Chromium has an atomic number of 24. The expected electron configuration based on the Aufbau principle is 1s² 2s² 2p⁶ 3s² 3p⁶ 4s² 3d⁴. However, the actual electron configuration is 1s² 2s² 2p⁶ 3s² 3p⁶ 4s¹ 3d⁵.

Orbital Diagram (actual):
- 4s: ↑
- 3d: ↑ ↑ ↑ ↑ ↑
Analysis: One electron from the 4s orbital moves to the 3d orbital, resulting in a half-filled 3d sublevel. This configuration is more stable because all six electrons (one in 4s and five in 3d) are unpaired.
Conclusion: Chromium has six unpaired electrons, making it a paramagnetic atom.

Copper (Cu)

Copper has an atomic number of 29. The expected electron configuration is 1s² 2s² 2p⁶ 3s² 3p⁶ 4s² 3d⁹. However, the actual electron configuration is 1s² 2s² 2p⁶ 3s² 3p⁶ 4s¹ 3d¹⁰.

Orbital Diagram (actual):
- 4s: ↑
- 3d: ↑↓ ↑↓ ↑↓ ↑↓ ↑↓
Analysis: One electron from the 4s orbital moves to the 3d orbital, resulting in a fully filled 3d sublevel and a half-filled 4s orbital. This configuration is more stable.
Conclusion: Copper has one unpaired electron, making it a paramagnetic atom.

Paramagnetism of Ions

The concept of paramagnetism also applies to ions. When atoms gain or lose electrons to form ions, their electron configurations change, which can affect their magnetic properties.

Example 1: Iron(II) Ion (Fe²⁺)

Iron(II) ion is formed when iron loses two electrons. Iron has an electron configuration of 1s² 2s² 2p⁶ 3s² 3p⁶ 4s² 3d⁶. When it loses two electrons, it typically loses them from the 4s orbital first. The electron configuration of Fe²⁺ becomes 1s² 2s² 2p⁶ 3s² 3p⁶ 3d⁶.

Orbital Diagram (3d):
- 3d: ↑↓ ↑ ↑ ↑ ↑
Analysis: The 3d sublevel has six electrons, resulting in four unpaired electrons.
Conclusion: Fe²⁺ is paramagnetic.

Example 2: Zinc Ion (Zn²⁺)

Zinc has an electron configuration of 1s² 2s² 2p⁶ 3s² 3p⁶ 4s² 3d¹⁰. When zinc loses two electrons to form Zn²⁺, it loses the 4s electrons. The electron configuration of Zn²⁺ becomes 1s² 2s² 2p⁶ 3s² 3p⁶ 3d¹⁰.

Orbital Diagram (3d):
- 3d: ↑↓ ↑↓ ↑↓ ↑↓ ↑↓
Analysis: The 3d sublevel is fully occupied with paired electrons.
Conclusion: Zn²⁺ is diamagnetic.

Factors Affecting the Strength of Paramagnetism

The strength of paramagnetism depends on several factors:

Number of Unpaired Electrons: The more unpaired electrons an atom or ion has, the stronger its paramagnetic effect.
Temperature: Paramagnetism is temperature-dependent. As temperature increases, the thermal motion of atoms disrupts the alignment of unpaired electrons with the external magnetic field, reducing the magnetic susceptibility.
Magnetic Field Strength: The stronger the external magnetic field, the greater the alignment of unpaired electrons, resulting in a stronger attraction.

Applications of Paramagnetism

Paramagnetic materials have various applications in science and technology:

Magnetic Resonance Imaging (MRI): Paramagnetic contrast agents, such as gadolinium complexes, are used in MRI to enhance the visibility of certain tissues and organs.
Catalysis: Paramagnetic metal ions are used as catalysts in various chemical reactions, facilitating electron transfer processes.
Research: Paramagnetic compounds are used in research to study electron spin resonance (ESR) and to investigate the electronic structure of molecules and materials.

Distinguishing Paramagnetism from Diamagnetism

It's important to distinguish paramagnetism from diamagnetism:

Paramagnetism: Occurs in substances with unpaired electrons, which are attracted to magnetic fields.
Diamagnetism: Occurs in substances where all electrons are paired. Diamagnetic substances are weakly repelled by magnetic fields.

Most substances exhibit either paramagnetic or diamagnetic behavior, depending on their electronic structure.

Conclusion

Identifying paramagnetic atoms using orbital diagrams involves understanding electron configurations, Hund's rule, and the presence of unpaired electrons. By drawing orbital diagrams and analyzing the distribution of electrons, one can determine whether an atom or ion is paramagnetic or diamagnetic. Paramagnetism is a fundamental property that plays a crucial role in various applications, from medical imaging to catalysis. Understanding this concept is essential for students and professionals in chemistry, physics, and materials science. By carefully examining the electron configurations and applying the principles of quantum mechanics, we can accurately predict and utilize the magnetic properties of elements and their compounds.