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Oliver Wright
Oliver Wright

Limitations of Valence Bond Theory and How Molecular Orbital Theory Can Overcome Them - PDF Download



# Limitations of Valence Bond Theory PDF Download ## Introduction - Explain what valence bond theory is and how it describes covalent bond formation and molecular structure - Mention some applications of valence bond theory in chemistry - State the main purpose of the article: to discuss the limitations of valence bond theory and provide a PDF download link for further reading ## Limitation 1: Inability to Explain Color and Magnetic Properties - Explain how valence bond theory fails to account for the color and magnetic properties of some coordination compounds - Give examples of such compounds and their observed properties - Explain how molecular orbital theory can better explain these properties by considering the splitting of d-orbitals and the formation of molecular orbitals ## Limitation 2: Inability to Explain Resonance Structures - Explain how valence bond theory fails to account for the resonance structures of some molecules - Give examples of such molecules and their resonance structures - Explain how molecular orbital theory can better explain these structures by considering the delocalization of electrons and the formation of molecular orbitals ## Limitation 3: Inability to Explain Bond Energies and Bond Lengths - Explain how valence bond theory fails to account for the variation in bond energies and bond lengths of some molecules - Give examples of such molecules and their bond energies and bond lengths - Explain how molecular orbital theory can better explain these variations by considering the overlap of atomic orbitals and the formation of bonding and antibonding molecular orbitals ## Limitation 4: Inability to Explain Hypervalency - Explain how valence bond theory fails to account for the hypervalency of some atoms (having more than eight electrons in their valence shell) - Give examples of such atoms and their hypervalent compounds - Explain how molecular orbital theory can better explain hypervalency by considering the involvement of d-orbitals and the formation of molecular orbitals ## Limitation 5: Inability to Explain Aromaticity - Explain how valence bond theory fails to account for the aromaticity of some molecules (having special stability due to cyclic delocalization of electrons) - Give examples of such molecules and their aromatic properties [^2 I'll try to create that. Here is the first section of the article. # Limitations of Valence Bond Theory PDF Download ## Introduction Valence bond theory is one of the basic theories that were developed to use the methods of quantum mechanics to explain chemical bonding. It focuses on how the atomic orbitals of the dissociated atoms combine to give individual chemical bonds when a molecule is formed . Valence bond theory can describe covalent bond formation as well as the electronic structure of molecules . Valence bond theory has many applications in chemistry, such as predicting the geometry, hybridization, and magnetic behavior of molecules . It can also explain some aspects of coordination chemistry, such as the formation of complexes between transition metals and ligands . However, valence bond theory has some limitations that prevent it from explaining all the observed properties and behaviors of molecules and complexes. In this article, we will discuss five major limitations of valence bond theory and provide a PDF download link for further reading. The limitations are: - Inability to explain color and magnetic properties of some coordination compounds - Inability to explain resonance structures of some molecules - Inability to explain bond energies and bond lengths of some molecules - Inability to explain hypervalency of some atoms - Inability to explain aromaticity of some molecules ## Limitation 1: Inability to Explain Color and Magnetic Properties One of the limitations of valence bond theory is that it cannot account for the color and magnetic properties of some coordination compounds. Coordination compounds are formed by the interaction of a central metal atom or ion with a set of ligands, which are molecules or ions that donate electrons to the metal . According to valence bond theory, the metal-ligand bond is essentially covalent in nature, and the electrons donated by the ligands occupy the atomic orbitals of the metal . However, this does not explain why some coordination compounds are colored and others are not, or why some are paramagnetic (attracted by a magnetic field) and others are diamagnetic (repelled by a magnetic field). For example, the complexes containing iron (II) can exhibit green and pale green colors, but the coordination compounds containing iron (III) have a brown or yellowish-brown color . Similarly, the complexes containing cobalt (II) can be pink, blue, or green depending on the ligands, but the complexes containing cobalt (III) are usually yellow . These color differences cannot be explained by valence bond theory. Moreover, valence bond theory predicts that all coordination compounds should be paramagnetic because they have unpaired electrons in their valence shells . However, this is not always true. For example, the complex [Ni(CN)4]2- is diamagnetic even though nickel (II) has two unpaired electrons in its 3d orbitals . This magnetic behavior cannot be explained by valence bond theory. A better way to explain the color and magnetic properties of coordination compounds is to use molecular orbital theory. Molecular orbital theory considers that the atomic orbitals of the metal and the ligands combine to form molecular orbitals that span the whole complex . These molecular orbitals can be classified into bonding, nonbonding, and antibonding orbitals depending on their energy and stability. The color of a coordination compound depends on the energy difference between two sets of molecular orbitals: the t2g set and the eg set. These sets are derived from the splitting of the five d orbitals of the metal due to the electrostatic interaction with the ligands . The t2g set consists of three lower-energy orbitals (dxy, dyz, dxz), and the eg set consists of two higher-energy orbitals (dx2-y2, dz2). When a coordination compound absorbs light, an electron can be promoted from a t2g orbital to an eg orbital. The energy difference between these two sets corresponds to a specific wavelength of light in the visible region. The color we see is the complementary color of the absorbed light [^1 I'll try to create that. Here is the second section of the article. ## Limitation 2: Inability to Explain Resonance Structures Another limitation of valence bond theory is that it cannot account for the resonance structures of some molecules. Resonance structures are a set of two or more Lewis structures that collectively describe the electronic bonding of a single polyatomic species, including fractional bonds and fractional charges . Resonance structures describe the delocalization of electrons within molecules that have a lone pair or a double or triple bond on the atom next to a double bond . The actual structure is an average of the resonance structures, called a resonance hybrid . Resonance is not the same as isomerism, which involves different arrangements of atomic nuclei in space . Resonance structures have the same positions of atoms, but different positions of electrons . The double-headed arrow indicates that the actual electronic structure is an average of those shown, not that the molecule oscillates between the two structures . For example, ozone (O3) has two resonance structures, as shown below . ![Ozone resonance structures](https://chem.libretexts.org/@api/deki/files/2231/Resonance_O3.png?revision=1) Both predict one OO single bond and one O=O double bond. However, the actual structure of ozone has two identical OO bonds that are shorter than a single bond and longer than a double bond . The bond order (the number of bonds between two atoms) is 1.5, which is a fractional value . The actual structure of ozone is a resonance hybrid of the two resonance structures. Valence bond theory cannot explain resonance structures because it assumes that electrons are localized in specific atomic orbitals or bonds . It does not account for the delocalization or spreading out of electrons over several atoms or bonds . Valence bond theory also does not allow for fractional bond orders or charges . A better way to explain resonance structures is to use molecular orbital theory. Molecular orbital theory considers that the atomic orbitals of the atoms combine to form molecular orbitals that span the whole molecule . These molecular orbitals can be classified into bonding, nonbonding, and antibonding orbitals depending on their energy and stability. The delocalization of electrons in resonance structures can be explained by the formation of molecular orbitals that extend over more than two atoms [^1 I'll try to create that. Here is the third section of the article. ## Limitation 3: Inability to Explain Bond Energies and Bond Lengths Another limitation of valence bond theory is that it cannot account for the variation in bond energies and bond lengths of some molecules. Bond energy or bond dissociation energy is the energy needed to break a covalent bond between two atoms of a diatomic covalent compound in its gaseous state . Bond length or bond distance is the equilibrium distance between the nuclei of two bonded atoms . According to valence bond theory, the bond energy and bond length depend on the type and number of bonds between two atoms . For example, a single bond has lower bond energy and longer bond length than a double bond, which has lower bond energy and longer bond length than a triple bond . However, this does not explain why some molecules have different bond energies and bond lengths for the same type and number of bonds. For example, the CC single bond in ethane (C2H6) has a bond energy of 347 kJ/mol and a bond length of 154 pm, while the CC single bond in ethene (C2H4) has a bond energy of 612 kJ/mol and a bond length of 134 pm . Similarly, the OO single bond in oxygen (O2) has a bond energy of 498 kJ/mol and a bond length of 121 pm, while the OO single bond in ozone (O3) has a bond energy of 355 kJ/mol and a bond length of 127 pm . These variations cannot be explained by valence bond theory. A better way to explain the variation in bond energies and bond lengths is to use molecular orbital theory. Molecular orbital theory considers that the atomic orbitals of the atoms combine to form molecular orbitals that span the whole molecule . These molecular orbitals can be classified into bonding, nonbonding, and antibonding orbitals depending on their energy and stability. The variation in bond energies and bond lengths can be explained by the overlap of atomic orbitals and the formation of bonding and antibonding molecular orbitals . The greater the overlap of atomic orbitals, the stronger and shorter the bonding molecular orbital. The greater the number of electrons in bonding molecular orbitals than in antibonding molecular orbitals, the higher the net bonding effect. For example, in ethane, each CC single bond is formed by the overlap of one sp3 hybrid orbital from each carbon atom, resulting in a sigma (σ) bonding molecular orbital. In ethene, each CC double bond is formed by the overlap of one sp2 hybrid orbital from each carbon atom, resulting in a sigma (σ) bonding molecular orbital, and by the overlap of one p orbital from each carbon atom, resulting in a pi (π) bonding molecular orbital. The double overlap in ethene results in a stronger and shorter CC bond than in ethane. Similarly, in oxygen, each OO single bond is formed by two electrons occupying a sigma (σ) bonding molecular orbital and two electrons occupying a sigma (σ*) antibonding molecular orbital, resulting in no net bonding effect. In ozone, each OO single bond is formed by one electron occupying a sigma (σ) bonding molecular orbital and one electron occupying a sigma (σ*) antibonding molecular orbital, resulting in a weak net bonding effect. The weak overlap in ozone results in a weaker and longer OO bond than in oxygen. I'll try to create that. Here is the fifth section of the article. ## Limitation 5: Inability to Explain Aromaticity Another limitation of valence bond theory is that it cannot account for the aromaticity of some molecules. Aromaticity is a property of conjugated cycloalkenes in which the stabilization of the molecule is enhanced due to the ability of the electrons in the π orbitals to delocalize . Aromatic molecules are very stable and do not break apart easily. They also have special chemical and physical properties, such as distinct ring currents and characteristic absorption spectra . Aromaticity is not the same as unsaturation, which involves the presence of double or triple bonds in a molecule . Aromatic molecules have a specific set of criteria that must be met in order to be considered aromatic. These criteria are: - The molecule must be cyclic (a ring of atoms) - The molecule must be planar (all atoms in the molecule lie in the same plane) - The molecule must be fully conjugated (p orbitals at every atom in the ring) - The molecule must have 4n+2 π electrons (n=0 or any positive integer), according to Hückel's rule For example, benzene (C6H6) is an aromatic molecule because it meets all these criteria. It has a six-membered ring of carbon atoms, each with one hydrogen atom attached. It is planar and fully conjugated, with one p orbital on each carbon atom overlapping with its neighbors to form a continuous ring of π electrons. It has 6 π electrons, which satisfies Hückel's rule when n=1 . Valence bond theory cannot explain aromaticity because it assumes that electrons are localized in specific atomic orbitals or bonds . It does not account for the delocalization or spreading out of electrons over several atoms or bonds . Valence bond theory also does not allow for fractional bond orders or charges . A better way to explain aromaticity is to use molecular orbital theory. Molecular orbital theory considers that the atomic orbitals of the atoms combine to form molecular orbitals that span the whole molecule . These molecular orbitals can be classified into bonding, nonbonding, and antibonding orbitals depending on their energy and stability. The delocalization of electrons in aromatic molecules can be explained by the formation of molecular orbitals that extend over more than two atoms [^1 I'll try to create that. Here is the conclusion and FAQs of the article. ## Conclusion In this article, we have discussed the limitations of valence bond theory and how molecular orbital theory can overcome them. Valence bond theory is a useful approach to explain covalent bond formation and molecular structure, but it fails to account for some important phenomena, such as: - Color and magnetic properties of some coordination compounds - Resonance structures of some molecules - Bond energies and bond lengths of some molecules - Hypervalency of some atoms - Aromaticity of some molecules Molecular orbital theory is a more advanced and accurate approach that considers the formation of molecular orbitals from atomic orbitals and their energy and stability. Molecular orbital theory can explain these phenomena by considering: - The splitting of d-orbitals and the formation of molecular orbitals in coordination compounds - The delocalization of electrons and the formation of molecular orbitals in resonance structures - The overlap of atomic orbitals and the formation of bonding and antibonding molecular orbitals in molecules - The involvement of d-orbitals and the formation of molecular orbitals in hypervalent compounds - The formation of molecular orbitals that extend over more than two atoms in aromatic molecules If you want to learn more about valence bond theory and molecular orbital theory, you can download this PDF file that contains more details and examples. ## FAQs Q1. What are the advantages of valence bond theory? A1. Valence bond theory has some advantages, such as: - It is simple and easy to understand - It can predict the geometry, hybridization, and magnetic behavior of many molecules - It can explain some aspects of coordination chemistry, such as the formation of complexes between transition metals and ligands Q2. What are the disadvantages of molecular orbital theory? A2. Molecular orbital theory has some disadvantages, such as: - It is complex and difficult to apply - It requires a lot of mathematical calculations and approximations - It does not provide a clear picture of the shape and size of molecules Q3. What is the difference between valence bond theory and Lewis theory? A3. Valence bond theory is an extension of Lewis theory that incorporates quantum mechanics. Lewis theory is based on the concept of electron pairs and octets, whereas valence bond theory is based on the concept of atomic orbitals and hybridization. Q4. What is the difference between bonding molecular orbitals and antibonding molecular orbitals? A4. Bonding molecular orbitals are formed by the constructive interference of atomic orbitals, which means that their wave functions add up to give a higher electron density between the nuclei. Antibonding molecular orbitals are formed by the destructive interference of atomic orbitals, which means that their wave functions cancel out to give a lower electron density between the nuclei. Q5. What is Hückel's rule? A5. Hückel's rule is a rule that helps determine if a cyclic, planar molecule with a ring of resonance bonds is aromatic or not. The rule states that if a molecule has 4n+2 π electrons, where n is any positive integer or zero, then it is aromatic. For example, benzene has 6 π electrons, which satisfies Hückel's rule when n=1.




limitations of valence bond theory pdf download

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