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Valence Bond Theory versus Molecular Orbital Theory: A Comparison of their Strengths and Weaknesses



- What are some of the advantages and applications of valence bond theory? - What are some of the limitations and drawbacks of valence bond theory? H2: Limitation 1: Valence bond theory does not account for resonance - What is resonance and why is it important for some molecules? - How does valence bond theory fail to explain resonance? - What are some examples of molecules that exhibit resonance and how do they differ from valence bond theory predictions? H3: Limitation 2: Valence bond theory does not account for hybridization - What is hybridization and why is it important for some molecules? - How does valence bond theory fail to explain hybridization? - What are some examples of molecules that exhibit hybridization and how do they differ from valence bond theory predictions? H4: Limitation 3: Valence bond theory does not account for molecular orbital theory - What is molecular orbital theory and how does it explain chemical bonding? - How does valence bond theory fail to explain molecular orbital theory? - What are some examples of molecules that exhibit molecular orbital theory and how do they differ from valence bond theory predictions? H2: Limitation 4: Valence bond theory does not account for intermolecular forces - What are intermolecular forces and why are they important for some molecules? - How does valence bond theory fail to explain intermolecular forces? - What are some examples of molecules that exhibit intermolecular forces and how do they differ from valence bond theory predictions? H3: Limitation 5: Valence bond theory does not account for relativistic effects - What are relativistic effects and why are they important for some molecules? - How does valence bond theory fail to explain relativistic effects? - What are some examples of molecules that exhibit relativistic effects and how do they differ from valence bond theory predictions? H4: Limitation 6: Valence bond theory does not account for quantum mechanics - What is quantum mechanics and how does it explain chemical bonding? - How does valence bond theory fail to explain quantum mechanics? - What are some examples of molecules that exhibit quantum mechanics and how do they differ from valence bond theory predictions? H2: Conclusion - Summarize the main points of the article - Emphasize the importance of understanding the limitations of valence bond theory - Suggest some alternative or complementary theories that can overcome the limitations of valence bond theory H3: FAQs - Q1: What is the difference between valence bond theory and Lewis structure? - Q2: What is the difference between valence bond theory and VSEPR theory? - Q3: What is the difference between valence bond theory and electronegativity? - Q4: What is the difference between valence bond theory and covalent bond? - Q5: What is the difference between valence bond theory and ionic bond? # Article with HTML formatting Introduction




Chemical bonding is one of the most fundamental concepts in chemistry, as it explains how atoms combine to form molecules. There are different theories that attempt to describe how chemical bonding works, such as Lewis structure, VSEPR theory, electronegativity, covalent bond, ionic bond, etc. However, one of the most widely used and influential theories is valence bond theory.




limitations of valence bond theory pdf 11



Valence bond theory was developed in the 1920s by Linus Pauling and others, based on the idea that atoms form bonds by overlapping their outermost or valent electrons. According to this theory, a chemical bond is the result of the attraction between two positively charged nuclei and the shared pair of electrons between them. Valence bond theory also assumes that the shape and orientation of the molecules depend on the number and type of bonds that each atom forms, as well as the repulsion between the electron pairs.


Valence bond theory has many advantages and applications, such as explaining the formation and properties of simple molecules, predicting the bond angles and bond lengths, rationalizing the concept of bond order and bond energy, and providing a qualitative understanding of the reactivity and stability of molecules. Valence bond theory is also compatible with other theories, such as Lewis structure and VSEPR theory, which can be derived from it or used to complement it.


However, valence bond theory is not perfect and has some limitations and drawbacks. In fact, valence bond theory cannot explain some of the phenomena and characteristics of more complex molecules, such as resonance, hybridization, molecular orbital theory, intermolecular forces, relativistic effects, and quantum mechanics. In this article, we will explore these limitations in detail and see how they affect our understanding of chemical bonding. We will also look at some examples of molecules that exhibit these phenomena and how they differ from valence bond theory predictions.


Limitation 1: Valence bond theory does not account for resonance




One of the limitations of valence bond theory is that it does not account for resonance. Resonance is a phenomenon that occurs when a molecule can be represented by more than one valid Lewis structure, which differ only in the distribution of electrons. Resonance implies that the actual structure of the molecule is a hybrid or a combination of these structures, rather than any one of them. Resonance also affects the properties and behavior of the molecule, such as its stability, reactivity, polarity, etc.


Valence bond theory fails to explain resonance because it assumes that a molecule has a fixed and unique structure, determined by the overlapping of atomic orbitals. Valence bond theory also assumes that each bond has a definite bond order and bond length, which do not change with resonance. However, resonance shows that these assumptions are not valid for some molecules, as their structure, bond order, and bond length can vary depending on the resonance contributors.


Some examples of molecules that exhibit resonance are benzene (C6H6), ozone (O3), nitrate ion (NO3), carbonate ion (CO3), etc. For instance, benzene can be represented by two equivalent Lewis structures, which differ in the location of the double bonds. However, neither of these structures is accurate, as benzene actually has six identical bonds that are intermediate between single and double bonds. Valence bond theory cannot explain this phenomenon, as it would require six overlapping orbitals with different bond orders and lengths. Resonance also explains why benzene is more stable and less reactive than expected from valence bond theory.


Limitation 2: Valence bond theory does not account for hybridization




Another limitation of valence bond theory is that it does not account for hybridization. Hybridization is a phenomenon that occurs when an atom mixes its atomic orbitals to form new orbitals with different shapes and energies. Hybridization allows an atom to form more bonds than expected from its valence electrons or to form bonds with different angles than predicted by valence bond theory. Hybridization also affects the properties and behavior of the molecule, such as its geometry, polarity, etc.


Valence bond theory fails to explain hybridization because it assumes that an atom forms bonds by overlapping its original atomic orbitals. Valence bond theory also assumes that each orbital has a fixed shape and orientation, which do not change with hybridization. However, hybridization shows that these assumptions are not valid for some atoms, as their orbitals can change depending on the number and type of bonds they form.


hybrid orbitals for trigonal planar bonds with 120 angles, or sp hybrid orbitals for linear bonds with 180 angles.


Limitation 3: Valence bond theory does not account for molecular orbital theory




A third limitation of valence bond theory is that it does not account for molecular orbital theory. Molecular orbital theory is a theory that explains chemical bonding by considering the electrons as delocalized over the entire molecule, rather than localized between two atoms. According to this theory, a chemical bond is the result of the overlap of two atomic orbitals to form two molecular orbitals, one lower in energy and one higher in energy. The lower energy orbital is called the bonding orbital, and the higher energy orbital is called the antibonding orbital. Molecular orbital theory also assumes that the shape and orientation of the molecules depend on the distribution and occupancy of the molecular orbitals, as well as the symmetry and stability of the molecule.


Valence bond theory fails to explain molecular orbital theory because it assumes that a molecule has a discrete and fixed structure, determined by the overlapping of atomic orbitals. Valence bond theory also assumes that each bond has a definite bond order and bond length, which do not change with molecular orbital theory. However, molecular orbital theory shows that these assumptions are not valid for some molecules, as their structure, bond order, and bond length can vary depending on the molecular orbitals.


Some examples of molecules that exhibit molecular orbital theory are oxygen (O2), nitrogen (N2), carbon monoxide (CO), etc. For instance, oxygen can be represented by two equivalent Lewis structures, which differ in the number of unpaired electrons. However, neither of these structures is accurate, as oxygen actually has two unpaired electrons in its antibonding orbitals. Valence bond theory cannot explain this phenomenon, as it would require oxygen to have four unpaired electrons in its valence shell or to have a different bond order and length. Molecular orbital theory also explains why oxygen is paramagnetic and less stable than expected from valence bond theory.


Limitation 4: Valence bond theory does not account for intermolecular forces




A fourth limitation of valence bond theory is that it does not account for intermolecular forces. Intermolecular forces are forces that act between molecules, rather than within molecules. Intermolecular forces are weaker than intramolecular forces (such as covalent bonds), but they are still important for some molecules, as they affect their physical properties and behavior, such as boiling point, melting point, solubility, viscosity, etc. Intermolecular forces can be classified into different types, such as dipole-dipole forces, hydrogen bonds, London dispersion forces, etc., depending on the polarity and size of the molecules involved.


Valence bond theory fails to explain intermolecular forces because it assumes that a molecule is isolated and independent from other molecules. Valence bond theory also assumes that each molecule has a fixed and unique structure, determined by the overlapping of atomic orbitals. However, intermolecular forces show that these assumptions are not valid for some molecules, as their structure and properties can change depending on the interactions with other molecules.


Some examples of molecules that exhibit intermolecular forces are water (H2O), ammonia (NH3), methane (CH4), etc. For instance, water can form hydrogen bonds with other water molecules or with other polar molecules. Hydrogen bonds are a type of dipole-dipole force that occurs when a hydrogen atom is covalently bonded to a highly electronegative atom (such as oxygen or nitrogen) and is attracted to another electronegative atom on another molecule. Valence bond theory cannot explain this phenomenon, as it would require water to have a different shape and polarity than predicted by valence bond theory. Intermolecular forces also explain why water has a high boiling point and high surface tension compared to other similar molecules.


Limitation 5: Valence bond theory does not account for relativistic effects




A fifth limitation of valence bond theory is that it does not account for relativistic effects. Relativistic effects are effects that occur when an object moves at a speed close to the speed of light, or when an object has a very high mass. Relativistic effects can affect the properties and behavior of some molecules, such as their mass, energy, momentum, length, time, etc. Relativistic effects can also affect the properties and behavior of some electrons, such as their speed, mass, charge, spin, etc.


Valence bond theory fails to explain relativistic effects because it assumes that a molecule follows the classical laws of physics, rather than the laws of special relativity or general relativity. Valence bond theory also assumes that each electron follows the Bohr model of the atom, rather than the Dirac equation or the Schrödinger equation. However, relativistic effects show that these assumptions are not valid for some molecules and electrons, as their properties and behavior can deviate from the classical predictions.


Some examples of molecules and electrons that exhibit relativistic effects are gold (Au), mercury (Hg), lead (Pb), etc. For instance, gold has a yellow color because its outermost electrons move at a very high speed (about 58% of the speed of light) and experience a relativistic increase in their mass and decrease in their radius. This causes them to absorb more blue light and reflect more yellow light than expected from valence bond theory. Valence bond theory cannot explain this phenomenon, as it would require gold to have a different electronic configuration and color than predicted by valence bond theory. Relativistic effects also explain why gold is more stable and less reactive than expected from valence bond theory.


Limitation 6: Valence bond theory does not account for quantum mechanics




A sixth limitation of valence bond theory is that it does not account for quantum mechanics. Quantum mechanics is a branch of physics that explains the behavior of matter and energy at the smallest scales, such as atoms and subatomic particles. Quantum mechanics introduces some concepts and principles that are different from the classical laws of physics, such as uncertainty principle, wave-particle duality, superposition principle, entanglement, etc. Quantum mechanics can affect the properties and behavior of some molecules and electrons, such as their position, momentum, energy, spin, etc.


Valence bond theory fails to explain quantum mechanics because it assumes that a molecule follows the deterministic and continuous laws of physics, rather than the probabilistic and discrete laws of quantum mechanics. Valence bond theory also assumes that each electron follows the classical model of the particle, rather than the quantum model of the wave. However, quantum mechanics shows that these assumptions are not valid for some molecules and electrons, as their properties and behavior can exhibit quantum phenomena that are not explained by valence bond theory.


Some examples of molecules and electrons that exhibit quantum mechanics are hydrogen (H2), helium (He), lithium (Li), etc. For instance, hydrogen can exist in two different forms: ortho-hydrogen and para-hydrogen. These forms differ in the relative orientation of the spins of the two electrons in the molecule. Ortho-hydrogen has parallel spins and para-hydrogen has antiparallel spins. Valence bond theory cannot explain this phenomenon, as it would require hydrogen to have a fixed and unique spin state. Quantum mechanics also explains why hydrogen has different energy levels and spectra depending on its spin state.


Conclusion




In conclusion, valence bond theory is a useful and influential theory that explains how atoms form bonds by overlapping their valent electrons. Valence bond theory has many advantages and applications, such as explaining the formation and properties of simple molecules, predicting the bond angles and bond lengths, rationalizing the concept of bond order and bond energy, and providing a qualitative understanding of the reactivity and stability of molecules. Valence bond theory is also compatible with other theories, such as Lewis structure and VSEPR theory, which can be derived from it or used to complement it.


stabilities, reactivities, etc. than predicted by valence bond theory. These phenomena also require different or complementary theories that can overcome the limitations of valence bond theory, such as resonance theory, hybridization theory, molecular orbital theory, intermolecular force theory, relativistic theory, and quantum theory.


Therefore, it is important to understand the limitations of valence bond theory and to be aware of the other theories that can explain chemical bonding more accurately and completely. By doing so, we can gain a deeper and broader knowledge of chemistry and appreciate the diversity and complexity of the molecular world.


FAQs




Q1: What is the difference between valence bond theory and Lewis structure?


A1: Valence bond theory and Lewis structure are both theories that explain chemical bonding by using valent electrons. However, valence bond theory focuses on the overlapping of atomic orbitals to form bonds, while Lewis structure focuses on the arrangement of electrons around atoms to form bonds. Valence bond theory can explain the shape and orientation of molecules, while Lewis structure can explain the charge and polarity of molecules. Valence bond theory and Lewis structure can be derived from each other or used to complement each other.


Q2: What is the difference between valence bond theory and VSEPR theory?


A2: Valence bond theory and VSEPR theory are both theories that explain chemical bonding by using valent electrons. However, valence bond theory focuses on the overlapping of atomic orbitals to form bonds, while VSEPR theory focuses on the repulsion of electron pairs to determine the shape of molecules. Valence bond theory can explain the bond angles and bond lengths of molecules, while VSEPR theory can explain the geometry and symmetry of molecules. Valence bond theory and VSEPR theory can be derived from each other or used to complement each other.


Q3: What is the difference between valence bond theory and electronegativity?


A3: Valence bond theory and electronegativity are both concepts that explain chemical bonding by using valent electrons. However, valence bond theory focuses on the overlapping of atomic orbitals to form bonds, while electronegativity focuses on the tendency of atoms to attract electrons in a bond. Valence bond theory can explain the formation and properties of covalent bonds, while electronegativity can explain the formation and properties of polar covalent bonds and ionic bonds. Valence bond theory and electronegativity can be used together to predict the polarity and reactivity of molecules.


Q4: What is the difference between valence bond theory and covalent bond?


A4: Valence bond theory and covalent bond are both concepts that explain chemical bonding by using valent electrons. However, valence bond theory is a general theory that applies to any type of bond that involves overlapping of atomic orbitals, while covalent bond is a specific type of bond that involves sharing of electrons between two atoms. Valence bond theory can explain how covalent bonds are formed and what factors affect their strength and length, while covalent bond can explain how atoms achieve stability by completing their octet or duet. Valence bond theory and covalent bond are closely related and often used interc


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