Orbital hybridisation

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Not to be confused with s-p mixing in Molecular Orbital theory. See Molecular orbital diagram.

In chemistry, hybridisation (or hybridization) is the concept of mixing atomic orbitals into new hybrid orbitals (with different energies, shapes, etc., than the component atomic orbitals) suitable for the pairing of electrons to form chemical bonds in valence bond theory. Hybrid orbitals are very useful in the explanation of molecular geometry and atomic bonding properties. Although sometimes taught together with the valence shell electron-pair repulsion (VSEPR) theory, valence bond and hybridisation are in fact not related to the VSEPR model.[1]

Historical development

Chemist Linus Pauling first developed the hybridisation theory in order to explain the structure of molecules such as methane (CH4).[2] This concept was developed for such simple chemical systems, but the approach was later applied more widely, and today it is considered an effective heuristic for rationalising the structures of organic compounds.

Orbitals are a model representation of the behaviour of electrons within molecules. In the case of simple hybridisation, this approximation is based on atomic orbitals, similar to those obtained for the hydrogen atom, the only neutral atom for which the Schrödinger equation can be solved exactly. In heavier atoms, such as carbon, nitrogen, and oxygen, the atomic orbitals used are the 2s and 2p orbitals, similar to excited state orbitals for hydrogen. Hybrid orbitals are assumed to be mixtures of these atomic orbitals, superimposed on each other in various proportions. It provides a quantum mechanical insight to Lewis structures. Hybridisation theory finds its use mainly in organic chemistry.

spx and sdx terminology

This terminology describes the weight of the respective components of a hybrid orbital. For example, in methane, the C hybrid orbital which forms each carbon–hydrogen bond consists of 25% s character and 75% p character and is thus described as sp3 (read as s-p-three) hybridised. Quantum mechanics describes this hybrid as an sp3 wavefunction of the form N(s + √3pσ), where N is a normalization constant (here 1/2) and pσ is a p orbital directed along the C-H axis to form a sigma bond. The ratio of coefficients (denoted λ in general) is √3 in this example. Since the electron density associated with an orbital is proportional to the square of the wavefunction, the ratio of p-character to s-character is λ2 = 3. The p character or the weight of the p component is N2λ2 = 3/4.

For atoms forming equivalent hybrids with no lone pairs, there is a correspondence to the number and type of orbitals used. For example, sp3 hybrids are formed from one s and three p orbitals. However, in all other cases, there is no such correspondence. The two bond-forming hybrid orbitals of oxygen in water can be described as sp4, which means that they have 20% s character and 80% p character, but does not imply that they are formed from one s and four p orbitals. As a result, the amount of p-character is not restricted to integer values; i.e., hybridisations like sp2.5 are also readily described. For more information see variable hybridization.

An analogous notation is used to describe sdx hybrids. For example, the permanganate ion (MnO4–) has sd3 hybridisation with orbitals that are 25% s and 75% d.

Types of hybridisation

sp3 hybrids

Four sp3 orbitals.

Hybridisation describes the bonding atoms from an atom's point of view. That is, for a tetrahedrally coordinated carbon (e.g., methane CH4), the carbon should have 4 orbitals with the correct symmetry to bond to the 4 hydrogen atoms.

Carbon's ground state configuration is 1s2 2s2 2px1 2py1 or more easily read:

C↑↓↑↓↑↑ 
1s 2s 2px 2py 2pz

The carbon atom can utilize its two singly occupied p-type orbitals (the designations px py or pz are meaningless at this point, as they do not fill in any particular order), to form two covalent bonds with two hydrogen atoms, yielding the "free radical" methylene CH2, the simplest of the carbenes. The carbon atom can also bond to four hydrogen atoms by an excitation of an electron from the doubly occupied 2s orbital to the empty 2p orbital, so that there are four singly occupied orbitals.

C*↑↓↑↑↑↑
1s 2s 2px 2py 2pz

As the energy released by formation of two additional bonds more than compensates for the excitation energy required, the formation of four C-H bonds is energetically favoured.

Quantum mechanically, the lowest energy is obtained if the four bonds are equivalent which requires that they be formed from equivalent orbitals on the carbon. A set of four equivalent orbitals can be obtained which are linear combinations of the valence-shell (core orbitals are almost never involved in bonding) s and p wave functions[3] which are the four sp3 hybrids.

C*↑↓↑↑↑↑
1s sp3 sp3 sp3 sp3

In CH4, four sp3 hybrid orbitals are overlapped by hydrogen 1s orbitals, yielding four σ (sigma) bonds (that is, four single covalent bonds) of equal length and strength.

translates into

sp2 hybrids

Three sp2 orbitals.
Ethene structure

Other carbon based compounds and other molecules may be explained in a similar way as methane. For example, ethene (C2H4) has a double bond between the carbons.

For this molecule, carbon will sp2 hybridise, because one π (pi) bond is required for the double bond between the carbons, and only three σ bonds are formed per carbon atom. In sp2 hybridisation the 2s orbital is mixed with only two of the three available 2p orbitals:

C*↑↓↑↑↑↑
1s sp2 sp2 sp2 2p

forming a total of three sp2 orbitals with one p orbital remaining. In ethylene (ethene) the two carbon atoms form a σ bond by overlapping two sp2 orbitals and each carbon atom forms two covalent bonds with hydrogen by s–sp2 overlap all with 120° angles. The π bond between the carbon atoms perpendicular to the molecular plane is formed by 2p–2p overlap. The hydrogen–carbon bonds are all of equal strength and length, which agrees with experimental data.

sp hybrids

Two sp orbitals

The chemical bonding in compounds such as alkynes with triple bonds is explained by sp hybridisation.

C*↑↓↑↑↑↑
1s sp sp 2p 2p

In this model, the 2s orbital mixes with only one of the three p orbitals resulting in two sp orbitals and two remaining unchanged p orbitals. The chemical bonding in acetylene (ethyne) (C2H2) consists of sp–sp overlap between the two carbon atoms forming a σ bond and two additional π bonds formed by p–p overlap. Each carbon also bonds to hydrogen in a σ s–sp overlap at 180° angles.

Hybridisation and molecule shape

Hybridisation helps to explain molecule shape since the angles between bonds are (approximately) equal to the angles between hybrid orbitals, as explained above for the tetrahedral geometry of methane. As another example, the three sp2 hybrid orbitals are at angles of 120° to each other, so this hybridisation favours trigonal planar molecular geometry with bond angles of 120°. Other examples are given in the table below.

Classification Main group Transition metal[4]
AX2
  • Linear (180°)
  • sp hybridisation
  • E.g., CO2
  • Bent (90°)
  • sd hybridisation
  • E.g., VO2+
AX3
AX4
AX5 -
AX6 -

Main group compounds with lone pairs

For main group compounds with lone electron pairs, the s orbital lone pair can be hybridised to a certain extent with the bond pairs.[7] This is analogous to s-p mixing in molecular orbital theory, and maximizes energetic stability according to the Walsh diagram for the molecule. This rationalisation is applied to explain deviations from ideal bond angles, such as when only p orbitals are used for bonding, most commonly in second and third period elements.

  • Trigonal pyramidal (AX3E1)
    • The s-orbital can be hybridised with the three p-orbital bonds to give bond angles greater than 90°.
    • Ex. NH3
  • Bent (AX2E1-2)
    • The s-orbital lone pair can be hybridised with the two p-orbital bonds to give bond angles greater than 90°. The out-of-plane p-orbital can either be a lone pair or pi bond. If it is a lone pair, the in-plane and out-of-plane lone pairs are inequivalent, contrary to the common picture depicted by VSEPR theory (see below).
    • Exs. SO2, H2O
  • Monocoordinate (AX1E1-3)
    • The s-orbital lone pair can be hybridised with the p-orbital bond. The two out-of-line p-orbitals can either be lone pairs or pi bonds. The p-orbital lone pairs are not equivalent to the s-rich lone pair.
    • Exs. CO, SO, HF

Hybridisation of hypervalent molecules

Traditional description

In general chemistry courses and mainstream textbooks, hybridisation is often presented for main group AX5 and above, as well as for transition metal complexes, using the hybridisation scheme first proposed by Pauling.

Classification Main group Transition metal
AX2 -
  • Linear (180°)
  • sp hybridisation
  • E.g., Ag(NH3)2+
AX3 -
AX4 -
AX5
AX6
  • Octahedral (90°)
  • d2sp3 hybridisation
  • E.g., Mo(CO)6
AX7
AX8
AX9 -

In this notation, d orbitals of main group atoms are listed after the s and p orbitals since they have the same principal quantum number (n), while d orbitals of transition metals are listed first since the s and p orbitals have a higher n. Thus for AX5 molecules, sp3d hybridisation in the P atom involves 3s, 3p and 3d orbitals, while dsp3 for Fe involves 3d, 4s and 4p orbitals.

However, hybridisation of s, p and d orbitals together is no longer accepted, as more recent calculations based on molecular orbital theory have shown that in main-group molecules the d component is insignificant, while in transition metal complexes the p component is insignificant (see below).

Resonance description

As shown by computational chemistry, hypervalent molecules can only be stable given strongly polar (and weakened) bonds with electronegative ligands such as fluorine or oxygen to reduce the valence electron occupancy of the central atom to a maximum of 8[8] (or 12 for transition metals). This requires an explanation that invokes sigma resonance in addition to hybridisation, which implies that each resonance structure has its own hybridisation scheme. As a guideline, all resonance structures have to obey the octet rule for main group compounds and the dodectet (12) rule for transition metal complexes.

Classification Main group Transition metal
AX2 - Linear (180°)
AX3 - Trigonal planar (120°)
AX4 - Square planar (90°)
AX5 Trigonal bipyramidal (90°, 120°) Trigonal bipyramidal (90°, 120°)
  • Fractional hybridisation (s and d orbitals)
  • E.g., Fe(CO)5
AX6 Octahedral (90°) Octahedral (90°)
AX7 Pentagonal bipyramidal (90°, 72°) Pentagonal bipyramidal (90°, 72°)
  • Fractional hybridisation (s and three d orbitals)
  • E.g., V(CN)74−
AX8 Square antiprismatic Square antiprismatic
  • Fractional hybridisation (s and three p orbitals)
  • E.g., IF8−
  • Fractional hybridisation (s and four d orbitals)
  • E.g., Re(CN)83−
AX9 - Tricapped trigonal prismatic
  • Fractional hybridisation (s and five d orbitals)
  • E.g., ReH92−

Main group compounds with lone pairs

For hypervalent main group compounds with lone electron pairs, the bonding scheme can be split into two components: the "resonant bonding" component and the "regular bonding" component. The "regular bonding" component has the same description (see above), while the "resonant bonding" component consists of resonating bonds utilizing p orbitals. The table below shows how each shape is related to the two components and their respective descriptions.

Regular bonding component (marked in red)
Bent Monocoordinate -
Resonant bonding component Linear axis Seesaw (AX4E1) (90°, 180°, >90°) T-shaped (AX3E2) (90°, 180°) Linear (AX2E3) (180°)
Square planar equator - Square pyramidal (AX5E1) (90°, 90°) Square planar (AX4E2) (90°)
Pentagonal planar equator - Pentagonal pyramidal (AX6E1) (90°, 72°) Pentagonal planar (AX5E2) (72°)

Clarifying misconceptions

VSEPR electron domains and hybrid orbitals are different

The simplistic picture of hybridisation taught in conjunction with VSEPR theory does not agree with high-level theoretical calculations[7] despite its widespread usage in many textbooks. For example, following the guidelines of VSEPR, the hybridization of the oxygen in water is described with two equivalent lone electron-pairs.[9] However, molecular orbital calculations give orbitals that reflect the C2v symmetry of the molecule.[10] One of the two lone pairs is in a pure p-type orbital, with its electron density perpendicular to the H–O–H framework.[11] The other lone pair is in an approximately sp0.8 orbital that is in the same plane as the H–O–H bonding.[11] Photoelectron spectra confirm the presence of two different energies for the nonbonded electrons.[12]

Non-inclusion of d orbitals in main group compounds

Main article: Hypervalent molecule

In 1990, Magnusson published a seminal work definitively excluding the role of d-orbital hybridization in bonding in hypervalent compounds of second-row elements. This had long been a point of contention and confusion in describing these molecules using molecular orbital theory. Part of the confusion here originates from the fact that one must include d-functions in the basis sets used to describe these compounds (or else unreasonably high energies and distorted geometries result), and the contribution of the d-function to the molecular wavefunction is large. These facts were historically interpreted to mean that d-orbitals must be involved in bonding. However, Magnusson concludes in his work that d-orbital involvement is not implicated in hypervalency.[13]

Non-inclusion of p orbitals in transition metal complexes

Similarly, p orbitals have long been thought to be utilized by transition metal centers in bonding with ligands, hence the 18-electron description; however, recent molecular orbital calculations have found that such p orbital participation in bonding is insignificant,[14][15] even though the contribution of the p-function to the molecular wavefunction is calculated to be somewhat larger than that of the d-function in main group compounds.

Hybridization theory vs. Molecular Orbital theory

Hybridisation theory is an integral part of organic chemistry and in general discussed together with molecular orbital theory in advanced organic chemistry textbooks although for different reasons. One textbook notes that for drawing reaction mechanisms sometimes a classical bonding picture is needed with two atoms sharing two electrons.[16] It also comments that predicting bond angles in methane with MO theory is not straightforward. Another textbook treats hybridisation theory when explaining bonding in alkenes[17] and a third[18] uses MO theory to explain bonding in hydrogen but hybridisation theory for methane.

Bonding orbitals formed from hybrid atomic orbitals may be considered as localized molecular orbitals, which can be formed from the delocalized orbitals of molecular orbital theory by an appropriate mathematical transformation. For molecules with a closed electron shell in the ground state, this transformation of the orbitals leaves the total many-electron wave function unchanged. The hybrid orbital description of the ground state is therefore equivalent to the delocalized orbital description for explaining the ground state total energy and electron density, as well as the molecular geometry which corresponds to the minimum value of the total energy.

There is no such equivalence, however, for ionized or excited states with open electron shells. Hybrid orbitals cannot therefore be used to interpret photoelectron spectra, which measure the energies of ionized states, identified with delocalized orbital energies using Koopmans' theorem. Nor can they be used to interpret UV-visible spectra which correspond to electronic transitions between delocalized orbitals. From a pedagogical perspective, the hybridisation approach tends to over-emphasize localisation of bonding electrons and does not effectively embrace molecular symmetry as does MO theory.

See also

References

  1. ↑ Gillespie, R.J. (2004), "Teaching molecular geometry with the VSEPR model", Journal of Chemical Education 81 (3): 298–304, Bibcode:2004JChEd..81..298G, doi:10.1021/ed081p298 
  2. ↑ Pauling, L. (1931), "The nature of the chemical bond. Application of results obtained from the quantum mechanics and from a theory of paramagnetic susceptibility to the structure of molecules", Journal of the American Chemical Society 53 (4): 1367–1400, doi:10.1021/ja01355a027 
  3. ↑ McMurray, J. (1995). Chemistry Annotated Instructors Edition (4th ed.). Prentice Hall. p. 272. ISBN 978-0-131-40221-8
  4. ↑ Weinhold, Frank; Landis, Clark R. (2005). Valency and bonding: A Natural Bond Orbital Donor-Acceptor Perspective. Cambridge: Cambridge University Press. pp. 381–383. ISBN 978-0-521-83128-4. 
  5. ↑ 5.0 5.1 Kaupp, Martin (2001). ""Non-VSEPR" Structures and Bonding in d(0) Systems". Angew Chem Int Ed Engl. 40 (1): 3534–3565. doi:10.1002/1521-3773(20011001)40:19<3534::AID-ANIE3534>3.0.CO;2-#. 
  6. ↑ 6.0 6.1 King, R. Bruce (2000). "Atomic orbitals, symmetry, and coordination polyhedra". Coordination Chemistry Reviews 197: 141–168. 
  7. ↑ 7.0 7.1 Weinhold, Frank. "Rabbit Ears Hybrids, VSEPR Sterics, and Other Orbital Absurdities". University of Wisconsin. Retrieved 2012-11-11. 
  8. ↑ David L. Cooper , Terry P. Cunningham , Joseph Gerratt , Peter B. Karadakov , Mario Raimondi (1994). "Chemical Bonding to Hypercoordinate Second-Row Atoms: d Orbital Participation versus Democracy". Journal of the American Chemical Society 116 (10): 4414–4426. doi:10.1021/ja00089a033. 
  9. ↑ Petrucci R.H., Harwood W.S. and Herring F.G. "General Chemistry. Principles and Modern Applications" (Prentice-Hall 8th edn 2002) p. 441
  10. ↑ Levine I.N. “Quantum chemistry” (4th edn, Prentice-Hall) p. 470–2
  11. ↑ 11.0 11.1 Laing, Michael J. Chem. Educ. (1987) 64, 124–128 "No rabbit ears on water. The structure of the water molecule: What should we tell the students?"
  12. ↑ Levine p. 475
  13. ↑ E. Magnusson. Hypercoordinate molecules of second-row elements: d functions or d orbitals? J. Am. Chem. Soc. 1990, 112, 7940-7951. doi:10.1021/ja00178a014
  14. ↑ C. R. Landis, F. Weinhold (2007). "Valence and extra-valence orbitals in main group and transition metal bonding". Journal of Computational Chemistry 28 (1): 198–203. doi:10.1002/jcc.20492. 
  15. ↑ O’Donnell, Mark (2012). "Investigating P-Orbital Character In Transition Metal-to-Ligand Bonding". Brunswick, ME: Bowdoin College. Retrieved 2012-09-16. 
  16. ↑ Clayden, Jonathan; Greeves, Nick; Warren, Stuart; Wothers, Peter (2001). Organic Chemistry (1st ed.). Oxford University Press. p. 105. ISBN 978-0-19-850346-0. 
  17. ↑ Organic Chemistry, Third Edition Marye Anne Fox James K. Whitesell 2003 ISBN 978-0-7637-3586-9
  18. ↑ Organic Chemistry 3rd Ed. 2001 Paula Yurkanis Bruice ISBN 978-0-130-17858-9

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