Multiple covalent bonds, 11(2)
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Multiple Covalent Bonds. Bonding in acetylene. The shape of a molecule is determined only by the σ-orbitals forming bonds. Hybridization scheme
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In chemistry, hybridisation (or hybridization) is the concept of mixing atomic
orbitals to form new hybrid orbitals suitable for the qualitative description of atomic bonding
properties. Hybridised orbitals are very useful in the explanation of the shape of molecular
orbitals for molecules. It is an integral part of valence bond theory. Although sometimes taught
together with the valence shell electron-pair repulsion (VSEPR) theory, valence bond and
hybridization are in fact not related to the VSEPR model.[1] Contents [hide] 1 Historical
development 2 Types of hybridisation 2.1 sp3 hybrids 2.2 sp2 hybrids 2.3 sp hybrids 3
Hybridisation and molecule shape 3.1 Explanation of the shape of water 3.2 Controversy
regarding d-orbital participation 4 Hybridisation theory vs. MO theory 5 See also 6 External
links 7 References [edit]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 rationalizing the
structures of organic compounds. For quantitative calculations of electronic structure and
molecular properties, hybridisation theory is not as practical as molecular orbital theory.
Problems with hybridisation are especially notable when the d orbitals are involved in bonding,
as in coordination chemistry and organometallic chemistry. Although hybridisation schemes in
transition metal chemistry can be used, they are not generally as accurate. 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 atom for which an exact analytic solution to its Schrödinger equation is
known. In heavier atoms, like carbon, nitrogen, and oxygen, the atomic orbitals used are the 2s
and 2p orbitals, similar to excited state orbitals for hydrogen. Hybridised orbitals are assumed to
be mixtures of these atomic orbitals, superimposed on each other in various proportions. The
theory of hybridisation is most applicable under these assumptions. It gives a simple orbital
picture equivalent to Lewis structures. Hybridisation is not required to describe molecules, but
for molecules made up from carbon, nitrogen and oxygen (and to a lesser extent, sulfur and
phosphorus) the hybridisation theory/model makes the description much easier. The
hybridisation theory finds its use mainly in organic chemistry. Its explanation starts with the way
bonding is organized in methane. [edit]Types of hybridisation [edit]sp3 hybrids 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. The .
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Ch07. structure and syntesys alkenes
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ORGANIC CHEMISTRY.pptx
This document provides an overview of organic chemistry concepts including:
1. Carbon is unique due to its ability to form chains (catenation) and bonds (tetravalency), making it central to organic compounds. Hybridization allows carbon to form different types of bonds to satisfy its valence.
2. Organic compounds can be classified based on their structure as acyclic/aliphatic, cyclic/aromatic, or heterocyclic aromatic. Nomenclature systems like IUPAC provide standardized naming conventions.
3. Key concepts include structural representations showing bonding and 3D orientation, and classification of organic compounds based on functional groups and ring structures. Hybridization explains how carbon satisfies its valence to form
9. Hybrid Orbitals.ppt
The document discusses hybrid orbitals and how they relate to the geometry around carbon atoms. It explains that carbon forms sp, sp2, and sp3 hybrid orbitals for molecules with 2, 3, and 4 electron pairs around carbon respectively. This determines the valence shell electron pair repulsion (VSEPR) geometry and bond angles, leaving 0-2 unhybridized p orbitals available for pi bonding.
chapter-alkene.ppt
1) Alkenes are hydrocarbons containing a carbon-carbon double bond. They include many naturally occurring compounds like flavors and fragrances.
2) This chapter focuses on the general reaction of alkenes, which is electrophilic addition. It examines the consequences of alkene stereoisomerism and how double bonds are present in most organic molecules.
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Valence Bond Theory PPTX
The document discusses theories of covalent bonding including valence shell electron pair repulsion theory, valence bond theory, orbital hybridization, and molecular orbital theory. It focuses on how orbitals overlap to form bonds between atoms and the different types of hybrid orbitals that can form based on electron domain geometry. Examples are provided to illustrate sp, sp2, sp3, and other hybrid orbitals as well as sigma and pi bonding including delocalized pi bonding in benzene.
Topicity.pptx
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Multiple covalent bonds, 11(2)
- 1. Bonding in Alkanes, Alkenes and Alkynes Dr. K. Shahzad Baig Memorial University of Newfoundland (MUN) Canada Petrucci, et al. 2011. General Chemistry: Principles and Modern Applications. Pearson Canada Inc., Toronto, Ontario. Tro, N.J. 2010. Principles of Chemistry. : a molecular approach. Pearson Education, Inc.
- 2. Multiple Covalent Bonds Ethylene has a carbon-to-carbon double bond in its Lewis structure Ethylene is a planar molecule with 120o H – C – H, H – C – C and bond angles The hybridization scheme that produces a set of hybrid orbitals with a trigonal-planar orientation is sp2. The (sp2 + p) orbital set
- 3. The purple orbitals are sp2 hybrid orbitals; the red and blue orbitals are 2p , with the colors indicating their phase. The sp2 hybrid orbitals overlap along the line joining the bonded atoms a σ-bond. The orbitals overlap in a side-to-side fashion and form a π-bond. Notice that the phase of the p-orbitals is retained.
- 4. Bond type # σ bonds # π bonds Single (C-H) 1 0 Double (C=C) 1 1 Triple (C≡C) 1 2
- 5. Bonding the shape of a molecule is determined only by the σ-orbitals forming bonds. the terminal H-atoms can be easily twist or rotate about the s-bonds that join them to a C atom. To twist one –CH2 group out of the plane of the other, however, would reduce the amount of overlap of the p orbitals and weaken the bond. The double bond is rigid, and the molecule is planar. Additionally, in carbon-to-carbon multiple bonds, the σ-bond involves more extensive overlap than does the π=bond. As a result, a carbon-to-carbon double bond (σ + π) is stronger than a single bond (π) but not twice as strong. C - C 347 kJ/mol C = C 611 kJ/mol C Ξ C 837 kJ/mol
- 6. Bonding in C2H2 Bonding in acetylene, C2H2, is somewhat similar to that C2H4, The Lewis structure of features a triple covalent bond, H – C Ξ C – H. The molecule is linear, In the triple bond in C2H2 one of the carbon-to-carbon bonds is a σ-bond and two are π- bonds
Editor's Notes
- One of the bonds between the carbon atoms results from the overlap of sp2 hybrid orbitals from each atom.