Chapter 9: Problem 55
Compare the Lewis and molecular orbital treatments of the oxygen molecule.
Short Answer
Expert verified
Lewis structure lacks insight into O₂'s paramagnetism; MO theory explains it through unpaired electrons.
Step by step solution
01
Understanding Lewis Structure
The Lewis structure of the oxygen molecule (O₂) involves drawing the valence electrons of each oxygen atom. Oxygen has six valence electrons, so in O₂, the two oxygen atoms share two pairs of electrons to complete their octets, forming a double bond. This representation suggests that O₂ has no unpaired electrons, which contradicts its paramagnetic nature.
02
Analyzing Molecular Orbital Theory
Molecular Orbital (MO) Theory describes the bonding of O₂ through the combination of atomic orbitals to form molecular orbitals. Oxygen has a total of 16 electrons. In MO theory, these electrons fill the lowest energy orbitals first, following the Aufbau principle. For O₂, these are filled as: \1. \( ext{σ}_{1s}^2\text{σ}_{1s}^{*2}\text{σ}_{2s}^2\text{σ}_{2s}^{*2}\text{σ}_{2p_z}^2\text{π}_{2p_x}^2\text{π}_{2p_y}^2\)2. Then, two electrons fill the antibonding \( ext{π}^{*}_{2p_x}\) and \( ext{π}^{*}_{2p_y}\) orbitals with one electron each.
03
Explaining Paramagnetism of Oxygen
According to the MO theory, the presence of unpaired electrons in the antibonding \( ext{π}^{*}\) orbitals explains why oxygen is paramagnetic. In contrast, the Lewis structure fails to account for this since it depicts all electrons as paired.
04
Comparing Both Treatments
The Lewis structure provides a simple way to represent the bond between two oxygen atoms but does not provide insight into properties such as magnetism. MO theory provides a more comprehensive view, explaining both the bonding and the paramagnetic properties of O₂ due to unpaired electrons in the molecular orbitals.
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Key Concepts
These are the key concepts you need to understand to accurately answer the question.
Oxygen Molecule
The oxygen molecule (\(\text{O}_2\)) consists of two oxygen atoms bonded together. Understanding the behavior of this molecule involves examining both the Lewis structure and Molecular Orbital (MO) Theory.
The Lewis structure focuses on valence electrons and simpler bonding representations. According to it, two oxygen atoms share two pairs of electrons, represented as a double bond, resulting in a total of eight electrons around each oxygen atom—thereby satisfying the octet rule.
Despite its simplicity, this representation misses certain molecular characteristics, such as the paramagnetism of oxygen. The Lewis model implies a non-magnetic state as it shows all electrons paired.
However, MO theory involves deeper analysis through quantum mechanics and uses molecular orbitals to explain electron arrangements more accurately. MO theory reveals that the oxygen molecule has unpaired electrons, accounting for its paramagnetic characteristics.
The Lewis structure focuses on valence electrons and simpler bonding representations. According to it, two oxygen atoms share two pairs of electrons, represented as a double bond, resulting in a total of eight electrons around each oxygen atom—thereby satisfying the octet rule.
Despite its simplicity, this representation misses certain molecular characteristics, such as the paramagnetism of oxygen. The Lewis model implies a non-magnetic state as it shows all electrons paired.
However, MO theory involves deeper analysis through quantum mechanics and uses molecular orbitals to explain electron arrangements more accurately. MO theory reveals that the oxygen molecule has unpaired electrons, accounting for its paramagnetic characteristics.
Paramagnetism
Paramagnetism refers to a property of materials that are attracted to a magnetic field due to the presence of unpaired electrons within their atoms or molecules.
When the oxygen molecule is put into the context of paramagnetism, things become interesting. Lewis structures fail to show \(\text{O}_2\) as paramagnetic because they depict all electrons as paired. However, experimental evidence supports that oxygen is indeed paramagnetic.
Molecular Orbital (MO) Theory offers an explanation: it demonstrates that there are unpaired electrons in the antibonding orbitals of \(\text{O}_2\).
This alignment and the presence of these unpaired electrons mean that oxygen can interact with a magnetic field, thus being characterized as a paramagnetic substance.
When the oxygen molecule is put into the context of paramagnetism, things become interesting. Lewis structures fail to show \(\text{O}_2\) as paramagnetic because they depict all electrons as paired. However, experimental evidence supports that oxygen is indeed paramagnetic.
Molecular Orbital (MO) Theory offers an explanation: it demonstrates that there are unpaired electrons in the antibonding orbitals of \(\text{O}_2\).
This alignment and the presence of these unpaired electrons mean that oxygen can interact with a magnetic field, thus being characterized as a paramagnetic substance.
Bonding Theories
Bonding theories are essential to understand how atoms come together to form molecules and predict their properties. Lewis structure theory and Molecular Orbital Theory are two of the most discussed theories in this context.
Lewis structure provides a simplified model to represent how electrons pair up and form bonds between atoms, focusing on achieving stable electron configurations, often visualized with dots around element symbols.
In contrast, Molecular Orbital (MO) Theory extends this understanding by considering atomic orbitals and how they overlap to form new orbitals—molecular orbitals—where electrons are delocalized across the whole molecule. This theory provides insights into understanding molecular stability, bond order, and magnetism.
With molecules like oxygen, MO theory can explain phenomena that Lewis structures cannot, such as the paramagnetic nature observed due to unpaired electrons, aligning and providing a more comprehensive view of chemical bonding.
Lewis structure provides a simplified model to represent how electrons pair up and form bonds between atoms, focusing on achieving stable electron configurations, often visualized with dots around element symbols.
In contrast, Molecular Orbital (MO) Theory extends this understanding by considering atomic orbitals and how they overlap to form new orbitals—molecular orbitals—where electrons are delocalized across the whole molecule. This theory provides insights into understanding molecular stability, bond order, and magnetism.
With molecules like oxygen, MO theory can explain phenomena that Lewis structures cannot, such as the paramagnetic nature observed due to unpaired electrons, aligning and providing a more comprehensive view of chemical bonding.
Chemical Bonding
Chemical bonding is the attractive force that holds atoms together in a molecule, determining the substance's properties and behaviors. There are various types of chemical bonding, with covalent bonds being one of the most common forms found in molecules like \(\text{O}_2\).
In \(\text{O}_2\), oxygen atoms form a covalent double bond, sharing pairs of electrons to achieve more stable configurations.
Lewis structures often represent this bond clearly. However, chemical bonding complexities extend beyond pair sharing.
Molecular Orbital Theory highlights bonding further by showcasing interactions between electron orbital phases, creating bonding and antibonding molecular orbitals that deeply influence the molecule's magnetic and conductive properties.
This dual perspective helps elucidate why molecules behave a certain way under different circumstances, like magnetic fields, by matching theoretical predictions with empirical observations.
In \(\text{O}_2\), oxygen atoms form a covalent double bond, sharing pairs of electrons to achieve more stable configurations.
Lewis structures often represent this bond clearly. However, chemical bonding complexities extend beyond pair sharing.
Molecular Orbital Theory highlights bonding further by showcasing interactions between electron orbital phases, creating bonding and antibonding molecular orbitals that deeply influence the molecule's magnetic and conductive properties.
This dual perspective helps elucidate why molecules behave a certain way under different circumstances, like magnetic fields, by matching theoretical predictions with empirical observations.
Atomic Orbitals
Atomic orbitals are mathematical functions that describe the regions in an atom where electrons are likely to be found. They are fundamental for understanding the distribution of electrons in atoms and how these contribute to molecular formation.
Each orbital is characterized by a specific energy level and shape, with common orbitals being s, p, d, and f. In atoms, these orbitals lead to the arrangement of electrons into defined shells and subshells.
\(\text{O}_2\)'s electron configuration involves a series of atomic orbitals combining to form molecular orbitals, which in turn influence the molecule's properties.
Molecular orbital theory expands on atomic orbitals by suggesting that when atoms bond, their orbitals merge to form molecular orbitals that can hold shared or unpaired electrons, affecting the molecule's bond order and magnetic properties.
This understanding provides deep insight into how elements interact at an electronic level, forming complex molecular structures with diverse characteristics like those observed in the oxygen molecule.
Each orbital is characterized by a specific energy level and shape, with common orbitals being s, p, d, and f. In atoms, these orbitals lead to the arrangement of electrons into defined shells and subshells.
\(\text{O}_2\)'s electron configuration involves a series of atomic orbitals combining to form molecular orbitals, which in turn influence the molecule's properties.
Molecular orbital theory expands on atomic orbitals by suggesting that when atoms bond, their orbitals merge to form molecular orbitals that can hold shared or unpaired electrons, affecting the molecule's bond order and magnetic properties.
This understanding provides deep insight into how elements interact at an electronic level, forming complex molecular structures with diverse characteristics like those observed in the oxygen molecule.
