Chemical Bonding & Structure Further aspects of covalent bonding and structure Hybridization Ms. Thompson - HL Chemistry Wooster High School
Topic 14.2 Hybridization A hybrid orbital results from the mixing of different types of atomic orbitals on the same atom.
Nature of science The need to regard theories as uncertain hybridization in valence bond theory can help explain molecular geometries, but is limited. Quantum mechanics involves several theories explaining the same phenomena, depending on specific requirements.
Models, theories, assumptions, and deductions Treat models and theories in chemistry with a high degree of critical perspective Lewis dot structure and VESPR theories can help deduce the electron domains but do not explain covalent bonds or electronic structure of molecules - need quantum mechanics for that. Schrodinger s wave equation paved the way for quantum mechanics Helps us better understand chemical bonding and provides an understanding of shapes of molecules We ll further discuss: Valence bond theory (VBT) Molecular orbital theory (MOT)
Valence bond theory (VBT) Looks at how atoms approach each other to form a molecule Bonds form from overlap of atomic orbitals but each atoms retains their own respective atomic orbitals For example H2: + HA HB H H The above figure shows the interaction of two hydrogen atoms HA and HB, each having one electron, to form the diatomic molecule H2. There is a decrease in the energy (stabilization) of the system due to the interaction, a chemical bond is formed.
Valence bond theory (VBT) Let s look at potential energy versus distance: + 0 Y X At point X the potential energy is essentially zero, as the H atoms are too far to interact. At point Y the H atoms approach each other. Electron in atom A, ea, is attracted to nucleus in atom B, Hb. Simultaneously, - Z ea and eb repel each other and nuclei HA and HB repel each other. There is a decrease in potential energy going from X to Y as the attraction is greater than the repulsion At point Z the minimum potential energy is achieved. This represents the most stable state of H2.
Valence bond theory (VBT) The greater the degree of atomic orbital overlap, the stronger the bond will be. Can be applied to homonuclear diatomic molecules such as F2 and to heteronuclear diatomic molecules such as HCl or HF. Lewis structure shows all covalent bonds as the same and does not explain inherent differences between covalent bonds. VBT considers these changes in energy that occur on formation of a chemical bond Can be applied to electronic structure of polyatomic molecules. VBT has its own limitations when discussing hybridization so we must consider molecular orbital theory
Molecular orbital theory (MOT) VBT maintains that atomic orbitals do not change when there s overlap. MOT involves atomic orbitals overlapping but the overlap results in the formation of new orbitals called molecular orbitals. Electrons are assigned to these molecular orbitals and associated with the whole molecule and not the individual atom. MOT strives to explain this phenomenon by describing the electron density surrounding the entire molecule.
Molecular orbital theory (MOT) The two 1s orbitals on the hydrogen atom combine to form two new molecular orbitals One is additive and results in a bonding molecular orbital sigma (σ) and is lower in energy The other is difference and results in an anti-bonding molecular orbital sigma star (σ*) and is higher in energy
Molecular orbital theory (MOT) Can use a molecular orbital diagram (next slide) For molecular orbitals to be formed two conditions must be met: 1. The atomic orbitals must be relatively close in energy for effective overlap 2. The symmetry of the atomic orbitals (sign of the wavefunction, Ψ) must be identical. (i) i.e. both Ψ+ could be positive and sum to form the σ bonding molecular orbital. in contrast the subtractive combination would have one Ψ+ positive and one Ψ- negative resulting in σ* anti-bonding orbital
Molecular orbital theory (MOT) In MOT x atomic orbitals combine to form x new molecular orbitals i.e. two 1s 1 atomic orbitals combine to form two new molecular orbitals, σ and σ*. Many theories have advantages and their constraints. Quantum mechanics provides several theories that can explain the same phenomena, depending on specific requirements
Hybridization Hybridization is a term used to describe the mixing of atomic orbitals to generate a set of new hybrid orbitals that are equivalent. A hybrid orbital results from the mixing of different types of atomic orbital on the same atom VBT uses hybridization to provide an electronic description of polyatomic molecules such as CH4 and NH3 but it can also account for geometries of molecules. Three different types of hybridization: sp 3 as seen in methane sp 2 as seen in ethene sp as seen in ethyne
The formation of sp 3 hybrid orbitals in methane Methane is a hydrocarbon and is one the major components of natural gas Full electron configuration Carbon: 1s 1 2s 2 2p 2 Condensed electron configuration Carbon: [He]2s 2 2p 2 Ground-state orbital diagram: [He] 2s 2 2px 1 2py 1 2pz 0 Hydrogen has the electron configuration 1s 1 and some expect carbon two form two bonds with hydrogen in its ground-state forming CH2 instead of CH4 but CH2 is very unstable (but does exist) The tetrahedral nature of methane involves the hybridization of its one 2s and three 2p orbitals on the C atom.
The formation of sp 3 hybrid orbitals in methane STEP ONE One of the electrons in the 2s orbital of the ground-state configuration of carbon is promoted to the vacant 2pz orbital to form an excited state: [He] 2s 2 2px 1 2py 1 2pz 0 Gives four potential C H bonds but we know methane has a tetrahedral molecular geometry with H C H bond angles of 109.5º. Geometry is incompatible because the three 2p orbitals are 90º to each other
The formation of sp 3 hybrid orbitals in methane STEP TWO The four atomic orbitals 2s, 2px, 2py, and 2pz combine to form a set of four new hybrid orbitals The four new sp 3 hybrid orbitals are entirely equivalent to each other and each one consists of 25% s character and 75% p character, since they were formed from one s orbital and three p orbitals. The shape of each sp 3 hybrid orbital will have 75% p orbital characteristics combined with 25% s orbital characteristics The four sp 3 hybrid orbitals point to the corners of a tetrahedron + = 109.5º Central carbon
The formation of sp 3 hybrid orbitals in methane STEP THREE The final step involves overlap of each carbon sp 3 orbital with a hydrogen 1s 1 atomic orbital. Hydrogen has 1s 1 electron configuration and the s orbital is spherically symmetrical Each 1s 1 atomic orbital on hydrogen and each sp 3 hybrid orbital on central carbon contains one electron. Any molecule with a tetrahedral electron domain geometry on its interior central atom (based on four electrons domains) would be predicted to have sp 3 hybridization!!
The formation of sp 2 hybrid orbitals in ethene Ethene, C2H4 is another hydrocarbon and is an alkene with a C=C double bond. There are three electron domains around each carbon atom so the electron domain geometry is trigonal planar. The H C=C bond angle is 121.3º and H C H bond is 117º 117º To deduce the hybridization scheme of o the carbon atom we will use the orbital diagram as shown previously for methane [He] 2s 2 2px 1 2py 1 2pz 0
The formation of sp 2 hybrid orbitals in ethene STEP ONE One of the electrons in the 2s orbital of the ground state configuration is promoted to the vacant 2pz orbital to form an excited-state [He] 2s 2 2px 1 2py 1 2pz 0
The formation of sp 2 hybrid orbitals in ethene STEP TWO To account for the ~120º bond angles (based on three electron domains), the next step involves hybridization of the three atomic orbitals 2s, 2px, and 2py. These combine to form a set of three new sp 2 hybrid orbitals. The 2pz orbital remains unhybridized. The three new sp 2 hybrid orbitals on the carbon atom are entirely equivalent with 33% s character and 66.7% p character. They point to the corners of a trigonal planar system.
The formation of sp 2 hybrid orbitals in ethene STEP THREE The next step involves formation of the sigma bond along the internuclear axis by the overlap of two sp 2 hybrid orbitals, one on each carbon atom. A pi bond is formed from the sideways overlap of the two pz unhybridized atomic orbitals, with overlap regions above the internuclear axis.
The formation of sp 2 hybrid orbitals in ethene STEP FOUR The final step involves overlap of each remaining sp 2 orbital on carbon with a hydrogen 1s atomic orbital Any molecule with a trigonal-planar electron domain geometry will be predicted to have sp 2 hybridization!!!
The formation of sp hybrid orbitals in ethyne Ethyne, C2H2, is an alkyne and has one triple bond. Around each carbon atom there are two electron domains so the electron domain geometry is linear, with a 180º bond angle 180º To deduce the hybridization scheme of the carbon atom in ethyne, we will start with the orbital diagram as we did with methane [He] 2s 2 2px 1 2py 1 2pz 0
The formation of sp hybrid orbitals in ethyne STEP ONE One of the electrons in the 2s orbital of the ground state configuration is promoted to the vacant 2pz orbital to form an excited-state [He] 2s 2 2px 1 2py 1 2pz 0
The formation of sp hybrid orbitals in ethyne STEP TWO To account for the 180º bond angles (based on two electron domains), the next step involves hybridization of the two atomic orbitals 2s and 2px. These combine to form a set of two new sp hybrid orbitals. The 2py and 2pz orbital remains unhybridized. The two new sp hybrid orbitals on the carbon atom are entirely equivalent with 50% s character and 50% p character. They point in opposite directions and each contain one electron. The remaining orbitals are the 2py and 2pz unhybridized atomic orbitals, also containing one electron.
The formation of sp hybrid orbitals in ethyne STEP THREE The next step involves formation of the sigma bond along the internuclear axis by the overlap of two sp hybrid orbitals, one on each carbon atom. Two pi bonds are formed from the sideways overlap of the two py and pz unhybridized atomic orbitals, with overlap regions above and below the internuclear axis.
The formation of sp hybrid orbitals in ethyne STEP FOUR The final step involves overlap of each remaining sp orbital on carbon with a hydrogen 1s atomic orbital Any molecule with a linear electron domain geometry on its interior central atom would be predicted to have sp hybridization!!!
SUMMARY Number of electron domains Electron domain geometry Hybridization 4 tetrahedral sp 3 3 trigonal planar sp 2 2 linear sp
Practice Problem... I Do... Deduce the hybridization of the central nitrogen interior atom in ammonia a) NH3
Covalent structures The basic molecular geometries can therefore be summarized on the basis of two, three, or four pairs of electrons. Each pair of electrons is described as occupying an electronic domain - a field of electronic density. Number of electron domains Molecular geometry Bond angle Examples of molecules or ions having this shape two linear 180º AB2 BeCl2, CO2 three trigonal planar 120º AB3 BF3, [NO3] - four tetrahedral 109.5º AB4 CH4, [NH4] +, [ClO4] -
Covalent structures Number of electron domains Electron domain geometry Molecular geometry Bond angle Examples of molecules or ions having this shape three trigonal planar AB2E v-shaped (bent) <120º [NO2] -, SO2 four tetrahedral AB3E trigonal pyramidal <109.5º NH3, [SO3] 2-, H3O + four tetrahedral AB2E2 v-shaped (bent) <109.5º H2O, [NH2] -
Practice Problem... We Do... For each interior atom A, B, and C in a molecule of methyl propanoate deduce the electron domain geometry, the molecular geometry, the bond angles, and the hybridization state. A C B
Methyl propanoate Interior atom A B C Number of electron domains Electron domain geometry 4 3 4 tetrahedral trigonal planar tetrahedral Molecular geometry tetrahedral trigonal planar V-shaped (bent) Bond angle(s) H C H 109.5º C C O 120º C O C <109.5º Hybridization sp 3 sp 2 sp 3
Practice Problem... You Do... The condensed structural formula of phenylamine (traditional name aniline) is C6H5NH2. a. Using VESPR theory, deduce the electron domain and molecular geometries of the carbon and nitrogen atoms in phenylamine. b. A model of the molecule is shown (figure 18). In the model the two hydrogen atoms attache to nitrogen appear to be above the horizontal plane of the molecule.
Practice Problem... You Do... Figure 18: Molecular model of phenylamine Figure 19: Space-filling model of pheylamine
Practice Problem... You Do... Deduce the hybridization of the nitrogen from figure 18 and figure 19 and comment on the two models. c.a theoretical study of the electronic structure of phenylamine found the H N H bond angle in phenylamine to be 112.79º, which is very close to the experimental value from gas-phase microwave studies. Discuss what you may conclude about the molecular geometry around the nitrogen in the NH2 group in the structure of phenylamine, and deduce its hybridization state on this basis.
Covalent structures The basic molecular geometries can therefore be summarized on the basis of two, three, or four pairs of electrons. Each pair of electrons is described as occupying an electronic domain - a field of electronic density. Number of electron domains Molecular geometry Bond angle Examples of molecules or ions having this shape two linear 180º AB2 BeCl2, CO2 three trigonal planar 120º AB3 BF3, [NO3] - four tetrahedral 109.5º AB4 CH4, [NH4] +, [ClO4] -
Covalent structures Number of electron domains Electron domain geometry Molecular geometry Bond angle Examples of molecules or ions having this shape three trigonal planar AB2E v-shaped (bent) <120º [NO2] -, SO2 four tetrahedral AB3E trigonal pyramidal <109.5º NH3, [SO3] 2-, H3O + four tetrahedral AB2E2 v-shaped (bent) <109.5º H2O, [NH2] -
Topic 14.2 Hybridization A hybrid orbital results from the mixing of different types of atomic orbitals on the same atom.