Chemistry
Bond Hybridization
Bond hybridization refers to the mixing of atomic orbitals to form hybrid orbitals, which have different shapes and energies than the original atomic orbitals. This process occurs when atoms form covalent bonds and helps to explain the geometry and properties of molecules. Hybridization allows for the formation of strong and stable chemical bonds in a wide variety of compounds.
Written by Perlego with AI-assistance
Related key terms
1 of 5
9 Key excerpts on "Bond Hybridization"
eBook - PDF- Paul Flowers, Klaus Theopold, Richard Langley, William R. Robinson(Authors)
- 2019(Publication Date)
- Openstax(Publisher)
8.2 Hybrid Atomic Orbitals LEARNING OBJECTIVES By the end of this section, you will be able to: • Explain the concept of atomic orbital hybridization • Determine the hybrid orbitals associated with various molecular geometries 8.2 • Hybrid Atomic Orbitals 379 Thinking in terms of overlapping atomic orbitals is one way for us to explain how chemical bonds form in diatomic molecules. However, to understand how molecules with more than two atoms form stable bonds, we require a more detailed model. As an example, let us consider the water molecule, in which we have one oxygen atom bonding to two hydrogen atoms. Oxygen has the electron configuration 1s 2 2s 2 2p 4 , with two unpaired electrons (one in each of the two 2p orbitals). Valence bond theory would predict that the two O–H bonds form from the overlap of these two 2p orbitals with the 1s orbitals of the hydrogen atoms. If this were the case, the bond angle would be 90°, as shown in Figure 8.6, because p orbitals are perpendicular to each other. Experimental evidence shows that the bond angle is 104.5°, not 90°. The prediction of the valence bond theory model does not match the real-world observations of a water molecule; a different model is needed. FIGURE 8.6 The hypothetical overlap of two of the 2p orbitals on an oxygen atom (red) with the 1s orbitals of two hydrogen atoms (blue) would produce a bond angle of 90°. This is not consistent with experimental evidence. 1 Quantum-mechanical calculations suggest why the observed bond angles in H 2 O differ from those predicted by the overlap of the 1s orbital of the hydrogen atoms with the 2p orbitals of the oxygen atom. The mathematical expression known as the wave function, ψ, contains information about each orbital and the wavelike properties of electrons in an isolated atom. When atoms are bound together in a molecule, the wave functions combine to produce new mathematical descriptions that have different shapes.
eBook - ePub- M. S. Prasada Rao(Author)
- 2022(Publication Date)
- CRC Press(Publisher)
3 .Further, in case of BeH2 , BF3 and CH4 molecules the number of unpaired electrons present on the central atom is less than the number of covalent bonds formed. In order to explain this, it was presumed that the electrons are excited to the higher orbitals during the formation of covalent bonds. For example, beryllium has the excited state configuration 1s2 , 2s1 , 2px 1 . This should result in two nonequivalent bonds due to the overlap of the hydrogen atoms with 2s and 2px orbitals. However, in BeH2 both the bonds are equivalent and are at an angle of 180°. In order to explain such cases the concept of hybridization was introduced.According to the concept of hybridization, in cases where pure orbitals cannot affect good overlap in the formation of covalent bonds, combination of pure orbitals having same or similar energy takes place, resulting in the formation of equivalent hybrid orbitals. The number of hybrid orbitals formed is equal to the number of combining pure orbitals. If suppose ϕ1 and ϕ2 are the two combining atomic orbitals the resulting hybridized orbitals ψh 1 and ψh 2 areThe values of the coefficients should be such that each of the hybridized orbital is normalized and the two hybrid orbitals should be orthogonal to each other.ψ=h 1c 1ϕ 1+c 2ϕ 2ψ=h 2c 3ϕ 1+c 4ϕ 2.
eBook - PDF- William R. Robinson, Edward J. Neth, Paul Flowers, Klaus Theopold, Richard Langley(Authors)
- 2016(Publication Date)
- Openstax(Publisher)
The atoms that make up the element combine in various ways, ranging from the mostly ionic (NaCl) to the partially ionic (HCl) to what we will call purely covalent. At the most fundamental level, all chemical bonds involve electrons, and a significant percentage of chemical and physical properties can be explained by considering the location and separation of charge in a species. By understanding the structure of matter at the atomic level, we can begin to build an understanding of the behavior of matter at both the microscopic and macroscopic levels. An understanding of dipoles and partial ionic character is fundamental to understanding the interactions between particles, which we will examine in the chapter on liquids and solids. These intermolecular forces become important in the liquid and solid states of matter. 5.2 Hybrid Atomic Orbitals By the end of this section, you will be able to: • Explain the concept of atomic orbital hybridization • Determine the hybrid orbitals associated with various molecular geometries Thinking in terms of overlapping atomic orbitals is one way for us to explain how chemical bonds form in diatomic molecules. However, to understand how molecules with more than two atoms form stable bonds, we require a more detailed model. As an example, let us consider the water molecule, in which we have one oxygen atom bonding to two hydrogen atoms. Oxygen has the electron configuration 1s 2 2s 2 2p 4 , with two unpaired electrons (one in each of the two 2p orbitals). Valence bond theory would predict that the two O–H bonds form from the overlap of these two 2p orbitals with the 1s orbitals of the hydrogen atoms. If this were the case, the bond angle would be 90°, as shown in Figure 5.6, because p orbitals are perpendicular to each other. Experimental evidence shows that the bond angle is 104.5°, not 90°. The prediction of the valence bond theory model does not match the real-world observations of a water molecule; a different model is needed.
eBook - PDF- Edward J. Neth, Paul Flowers, Klaus Theopold, Richard Langley, William R. Robinson(Authors)
- 2019(Publication Date)
- Openstax(Publisher)
The atoms that make up the element combine in various ways, ranging from the mostly ionic (NaCl) to the partially ionic (HCl) to what we will call purely covalent. At the most fundamental level, all chemical bonds involve electrons, and a significant percentage of chemical and physical properties can be explained by considering the location and separation of charge in a species. By understanding the structure of matter at the atomic level, we can begin to build an understanding of the behavior of matter at both the microscopic and macroscopic levels. An understanding of dipoles and partial ionic character is fundamental to understanding the interactions between particles, which we will examine in the chapter on liquids and solids. These intermolecular forces 242 5 • Advanced Theories of Bonding Access for free at openstax.org become important in the liquid and solid states of matter. 5.2 Hybrid Atomic Orbitals LEARNING OBJECTIVES By the end of this section, you will be able to: • Explain the concept of atomic orbital hybridization • Determine the hybrid orbitals associated with various molecular geometries Thinking in terms of overlapping atomic orbitals is one way for us to explain how chemical bonds form in diatomic molecules. However, to understand how molecules with more than two atoms form stable bonds, we require a more detailed model. As an example, let us consider the water molecule, in which we have one oxygen atom bonding to two hydrogen atoms. Oxygen has the electron configuration 1s 2 2s 2 2p 4 , with two unpaired electrons (one in each of the two 2p orbitals). Valence bond theory would predict that the two O–H bonds form from the overlap of these two 2p orbitals with the 1s orbitals of the hydrogen atoms. If this were the case, the bond angle would be 90°, as shown in Figure 5.6, because p orbitals are perpendicular to each other. Experimental evidence shows that the bond angle is 104.5°, not 90°.
- Milton Orchin, Allan R. Pinhas, R. Marshall Wilson, Roger S. Macomber(Authors)
- 2005(Publication Date)
- Wiley-Interscience(Publisher)
Example. As two hydrogen atoms [H A (1) and H B (2)] approach each other, the 1s orbital overlap increases and the electron in each orbital feels the influence of the opposite nucleus and is increasingly associated with it. The molecule is stabilized at the appropriate H–H bond distance by the exchange between two equivalent struc- tures: H A (1) H B (2) and H A (2) H B (1), where A and B refer to the hydrogen atoms and (1) and (2) refer to electrons of opposite spin. The wave function for the occupied 52 BONDS BETWEEN ADJACENT ATOMS orbital of the hydrogen molecule may then be written as ψ [(φ A (1)φ B (2) φ A (2)φ B (1)], where φ A and φ B are the atomic 1s orbitals on hydrogens A and B. This wave function indicates that the two electrons are never associated together at one atom. SUGGESTED READING Bowser, J. R. Inorganic Chemistry. Brooks/Cole: Pacific Grove, CA, 1993. Coulson, C. A. Valence. Oxford University Press: London, 1952. Gray, H. B. Electrons and Chemical Bonding. W.A. Benjamin: New York, 1964. Harris, D. C., and Bertolucci, M. D. Symmetry and Spectroscopy. Oxford University Press: New York, 1978. Jaffé, H. H. and Orchin, M. “Hybridization in Carbon Monoxide.” Tetrahedron 10, 212 (1960). Jean, Y., Volatron, F., and Burdett, J. An Introduction to Molecular Orbitals. Oxford University Press: New York, 1993. Orchin, M., and Jaffé, H. H. Symmetry, Orbitals and Spectra. Wiley-Interscience: New York, 1971. Streitwieser, A. Molecular Orbital Theory for Organic Chemists. John Wiley & Sons: New York, 1961. WAVE FUNCTIONS IN VALENCE BOND (VB) THEORY 53
eBook - PDF- T. W. Graham Solomons, Craig B. Fryhle, Scott A. Snyder(Authors)
- 2017(Publication Date)
- Wiley(Publisher)
7. Hybrid atomic orbitals are obtained by mixing (hybridizing) the wave functions for orbitals of different types (i.e., s and p orbitals) but from the same atom. 8. Hybridizing three p orbitals with one s orbital yields four sp 3 orbitals. Atoms that are sp 3 hybridized direct the axes of their four sp 3 orbitals toward the corners of a tet- rahedron. The carbon of methane is sp 3 hybridized and tetrahedral. 9. Hybridizing two p orbitals with one s orbital yields three sp 2 orbitals. Atoms that are sp 2 hybridized point the axes of their three sp 2 orbitals toward the corners of an equilateral triangle. The carbon atoms of ethene are sp 2 hybridized and trigonal planar. A summary of sp 3 , sp 2 , and sp hybrid orbital geometries. HINT 44 CHAPTER 1 THE BASICS: Bonding and Molecular Structure 10. Hybridizing one p orbital with one s orbital yields two sp orbitals. Atoms that are sp hybridized orient the axes of their two sp orbitals in opposite directions (at an angle of 180°). The carbon atoms of ethyne are sp hybridized and ethyne is a linear molecule. 11. A sigma (σ) bond (a type of single bond) is one in which the electron density has circular symmetry when viewed along the bond axis. In general, the skeletons of organic molecules are constructed of atoms linked by sigma bonds. 12. A pi (π) bond, part of double and triple carbon–carbon bonds, is one in which the electron densities of two adjacent parallel p orbitals overlap sideways to form a bonding pi molecular orbital. 1.16 HOW TO PREDICT MOLECULAR GEOMETRY: THE VALENCE SHELL ELECTRON PAIR REPULSION MODEL We can predict the arrangement of atoms in molecules and ions on the basis of a relatively simple idea called the valence shell electron pair repulsion (VSEPR) model. We apply the VSEPR model in the following way: 1. We consider molecules (or ions) in which the central atom is covalently bonded to two or more atoms or groups.
eBook - PDFQuantum Nanochemistry, Volume Three
Quantum Molecules and Reactivity
- Mihai V. Putz(Author)
- 2016(Publication Date)
- Apple Academic Press(Publisher)
22 Quantum Nanochemistry—Volume III: Quantum Molecules and Reactivity • The number of involved functions ϕ i immediately becomes very large and therefore has, also without involving the excited orbital (states), about 4 millions for the simple ion as SO 4 − , for example; • Using truncating method (as in Heitler and London method – to be amended in the Section 1.4) and then applying the perturbation approx-imations (viz. adiabatic coupling), leads to uncertain procedures. For all these reasons, soon after the mesomerism, Hund and Mulliken had been developed a more consistent theory with the physical signifi -cance of molecular orbitals, which confers the atoms as the rooting role of molecular orbitals’ formation, from where the molecular orbitals (MO) method was as such nominated. 1.3 MOLECULAR ORBITAL THEORY OF BONDING The basic idea is simple: next to each nucleus, the most representative terms are those who correspond to the smallest nucleus-electron distances, so the total Hamiltonian is a little different respecting that of the isolated atom. Therefore the wave-functions which describe the electrons around each atom of bonding, called atomic orbitals , represent approximate local solutions of the Schrödinger’s equation in molecule; accordingly, one will search for the molecular function, called molecular orbital , as a linear combination of the atomic orbitals, or hybridized, of various atoms get-ting into the molecule (Hückel, 1934; Coulson, 1938, 1952; Hall, 1950; Griffith & Orgel, 1957; Jensen, 1999; Licker, 2004). Resuming, the chemical bond which may be formed between two atoms A and B is considered as the resulted MO/the wave-function by “composition” of the two atomic orbitals ϕ A and ϕ B which by overlap-ping/superposition constitute the bond Ψ AB A B c c = + 1 2 ϕ ϕ (1.62) without other approximations.
eBook - PDFBasic Physical Chemistry
The Route to Understanding
- E Brian Smith(Author)
- 2012(Publication Date)
- ICP(Publisher)
The 2p x and 2p y orbitals of the fluorine have zero overlap with the 1s electrons of the hydrogen and are termed non-bonding orbitals, as is the fluorine 2s orbital. Molecular orbital theory can also be extended to polyatomic molecules, but now we might expect to construct orbitals from a combination of the orbitals of the valence electrons of all atoms in the molecule. Such orbitals can be employed when studying aromatic compounds, but, for most other types of molecule, it is simpler to construct localised two-centre molecular orbitals that can represent individual chemical bonds. For example, in describing water, we could regard the bonds between the hydrogen and the oxygen atoms as involving the 1s orbitals of the hydrogen and two of the 2p orbitals of the oxygen. The fact that the bond angle in water, 104.5 ◦ , is little more than 90 ◦ indicates that the geometry of the water molecule is, indeed, largely determined by the shape of the p orbitals. We assume, to a good approximation, that bonds are independent of each other and, if we replace one of the hydrogen atoms by another group, the remaining bond will be substantially unchanged. Thus, each bond has its own characteristic features, such as energy and length, which are determined by the wave function of the bond. 5.6 Hybridisation The treatment of polyatomic molecules in terms of directed orbitals runs into a very serious problem. It is very hard to understand how carbon, with a ground state configuration 1s 2 2s 2 2p 2 , can have a valence of four and have all four bonds equivalent. With two unpaired p electrons, we would expect carbon to be divalent 98 | Basic Physical Chemistry and form compounds like water with a bond angle of approximately 90 ◦ . To form four bonds, one of the 2s electrons would need to be promoted to the remaining empty p orbital, giving four unpaired electrons to take part in the formation of chemical bonds.
eBook - PDFChemistry
The Molecular Nature of Matter
- Neil D. Jespersen, Alison Hyslop(Authors)
- 2021(Publication Date)
- Wiley(Publisher)
TOOLS Sigma and pi bonds and hybridization 9.7 Molecular Orbital Theory Basics 457 PRACTICE EXERCISE 9.21 Consider the molecule below. What kind of hybrid orbitals are used by atoms 1, 2, and 3? How many sigma bonds and pi bonds are in the molecule? (Hint: Study the brief summary above.) O H H H H 2 3 1 H C C C C H PRACTICE EXERCISE 9.22 Consider the molecule below. What kind of hybrid orbitals are used by atoms 1, 2, and 3? How many σ bonds and π bonds are in the molecule? H C C C C N H 3 H H H 2 1 9.7 Molecular Orbital Theory Basics Molecular orbital theory takes the view that molecules and atoms are alike in one important respect: Both have energy levels that correspond to various orbitals that can be populated by electrons. In atoms, these orbitals are called atomic orbitals; in molecules, they are called molecular orbitals. (We shall frequently call them MOs.) In most cases, the actual shapes and energies of molecular orbitals cannot be determined exactly. Nevertheless, reasonably good estimates of their shapes and energies can be obtained by combining the electron waves corresponding to the atomic orbitals of the atoms that make up the molecule. In forming molecular orbitals, these waves interact by constructive and destructive interference just like other waves we’ve seen in Chapter 7. Their intensities are either added or subtracted when the atomic orbitals overlap. Formation of Molecular Orbitals from Atomic Orbitals In Chapter 7 you learned that an atomic orbital is represented mathematically by a wave func- tion, ψ, and that the square of the wave function, ψ 2 , describes the distribution of electron density around the nucleus. In MO theory, molecular orbitals are also represented by wave functions, ψ MO , that when squared describe the distribution of electron density around the entire set of nuclei that make up the molecule.
Index pages curate the most relevant extracts from our library of academic textbooks. They’ve been created using an in-house natural language model (NLM), each adding context and meaning to key research topics.
