Molecular Geometry and Electron group Geometry - Chemistry

Chapter 10: Chemical Bonding II Molecular shapes and Bonding Theories Principles of Chemistry: A Molecular Approach,1st Ed. Nivaldo J. Tro

Dr. Azra Ghumman Memorial University of Newfoundland

Chemical Bonding II

10.2

The Valence-Shell Electron-Pair Repulsion (VSEPR) Theory: Five basic shapes VSEPR Theory: The effect of lone pairs VSEPR Theory: Predicting Molecular Geometries Dipole Moment and Molecular Geometry Valence Bond Theory: Orbital overlap as a Chemical bond. Valence Bond Theory: Hybridization of Atomic Orbitals Suggested reading

10.3 10.4 10.5 10.6 10.7 10.1

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Chapter 10: Chemical bonding II A. Ghumman

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Molecular Geometry and Electron group Geometry Shape of the molecules (three dimensional structure) effects their chemical and physical properties. Molecular Geometry (shape): The general shape of a molecule determined by the relative positions of the atomic nuclei. Electron group geometry (EGG): The geometry of a molecule determined by the arrangement of electron groups around the central atom Electron Groups(e- domains)- Lone pairs, single bond, multiple bonds and even single electrons

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The Valence Shell Electron Pair Repulsion (VSEPR) Theory Valence-Shell Electron Pair Repulsion model (VSEPR)- It predicts the shape of the molecules and ions by assuming the valence-shell electron-pairs are arranged as far away from one another as possible to minimize e- pairs repulsions. Electron groups repulsion determines the shape of the molecule.

It does not explain chemical bonding can be explained by Valence Bond Theory (based on Quantum mechanics)

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VSEPR Theory: The Five basic Shapes Five basic arrangements of e- groups around central atom leads to five different basic molecular shapes Linear, trigonal planar, tetrahedral, trigonal pyramidal and octahedral Points to remember For molecules that exhibit resonance, you can choose any resonance structure Electron group geometry is almost same as molecular geometry when all e- pairs are bonding

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Linear Geometry To find the geometry of the molecule AXn , you first determine the number of electron groups around the central atom A (X = electron group) When all e- pairs are bonding e- group geometry is almost same as molecular geometry Two electron groups (AX2) linear: bond angle 180°, e.g. CO2, HCN, BeCl2. Three electron groups (AX3):triangular planar three electron groups are farthest apart when they lie in the same plane and point towards the corners of the equilateral triangle, Ideal bond angle120°,e.g. BF3and Chem 1011


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Tetrahedral Geometry Four electron groups(AX4)- Tetrahedral geometry Tetrahedron- A three dimensional geometrical shape with four identical faces, each an equilateral triangle Four electron groups around the central atom will occupy positions in the shape of a tetrahedron around the central atom. The ideal bond angle is 109.5°e.g CH4

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Trigonal Bipyramidal Geometry Five electron groups (AX5)-Trigonal bipyramidal geometry Two types of positions: Equatorial (trigonal plane)Three electron pairs lie in a plane and point toward the corners of an equilateral triangle (bond angle 120°), Axial -the 4th electron pair points directly above and the fifth pair points below the central atom at 90° angle

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Octahedral Six electron groups(AX6)- Octahedral geometry Octahedron- Eight sided geometric shape the shape of two square base pyramids joined base to base, with the central atom in the center of the shared bases All positions are equivalent The bond angle is 90°

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Effect of Lone pair of electrons on bond angles Molecular geometry is different than e- group geometry Observed Bond angles are somewhat small than in ideal geometry a.The lone pair of e-s occupy larger volume than bonding electron pairs, pushing them closer to each other, reducing the bond angles than usual e.g. NH3 and H2O. b. Relative sizes of repulsive interactions are: Lone pair – Lone pair > Lone pair – Bonding pair > Bonding pair – Bonding pair c. Triple bond bond occupy more space than a double or single bond, in the following order Triple bond > double bond > single bond


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Three electron groups(AX3) Electron geometry Trigonal planar

Molecular geometry

Trigonal Planar

AX2E

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Bent


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Four electron Groups(AX4) (tetrahedral arrangement). Electron geometry Tetrahedral

Molecular geometry Tetrahedral

AX3E Trigonal pyramidal

AX2E2 Bent

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Geometries of Five electron Group(AX5) structures

Trigonal bipyramidal


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Subclasses : See-saw and T-shapes and linear Electron geometry

Electron geometry



AX3E2

AX4E

Molecular geometry See-saw Bond angle < 120, 90 and 180° Chem 1011

Molecular geometry T-shape Bond angle 90,180°


Position of lone pair

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Geometries of Six electron Groups with lone pairs Subclass AX5E: Five bonding e- groups and one lone pair a square pyramidal shape. AX4E2: Five bonding e- groups, and two lone pairs, a square planar shape.


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Geometries of six electron groups with lone pairs Electron geometry

Molecular geometry

Octahedral

Octahedral, Square pyramidal, square planar

Square pyramidal Bond angle < 90° Chem 1011


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Square planar Bond angle 90° 17

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Class Example Molecular Shape Bond angles AX2 BeF2 linear 180 ° AX3 BF3 Trigonal planar 120° AX2E SO2 bent < 120° AX4 CH4 Tetrahedral 109.5 ° AX3E NH3 trigonal pyramid, < 109° AX2 E2 H2O bent < 109° Molecules with expanded octet AX5 PF5 trigonal bipyramidal 90°, 120°, 180° AX4E SF4 Seesaw 90°, 120°, AX3E2 ClF3 T shaped 90° 180 AX2E3 XeF2 AX6 SF6 AX5E BrF5 AX4E2 XeF4 Chem 1011

linear octahedral square pyramidal square planar

180° 90° 90° 90°


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Rules for Predicting Molecular Geometry by VSEPR 1. Draw a Lewis structure for the molecule. 2. Determine how many electrons groups are around the central atom, including bonding and nonbonding pairs. Count a multiple bond (double or triple bond)as one pair or one group of electrons. In case of two or more resonace structures, you can use any one of them. 3. Determine the electron group geometry by noting the arrangement of the electron groups around that central atom. The electron pairs will orient themselves so that they stay as far away as possible from each other 4. Obtain the molecular geometry from the directions of bonding pairs for this arrangement (Table 10.1 )

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Sample Problem Predict the electron pair geometry and molecular geometry of the following molecules according to VSEPR model. NO2-, H3O+, ICl3, SF4, ClF5, , I3-

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To represent the 3-D shape of a molecule on a flat surface (paper) A straight solid line represents a bond in the plane of page. A dashed line( hashed wedge) represents a bond extending behind the page. A solid wedge represents a bond extending above (in front of) the page

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Shapes of larger molecules Molecules with more than one central atom we describe the shape around each central atom in sequence Ethylene (H2C CH2) 6 atoms in a plane; 120 ° Each C has trigonal planar geometry, as a result the entire molecule is planar.

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Dipole Moment and Molecular Geometry Dipole Moment : is a quantitative measure of the degrees of the charge separation in a molecule. Any molecule that has a net separation of charge (partial charges + and -) has a dipole moment e.g. H-Cl Units of Dipole moment: Debyes ‘D’ -30 1D = 3.34 x 10 C.m. (coulomb-meters) We can view the polarity of individual bonds within a molecule as vector quantities.

Thus, molecules that are perfectly symmetric have a zero dipole moment. considered nonpolar.

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Dipole Moment and Molecular Geometry You can relate the presence or absence of a dipole moment in a molecule to its molecular geometry. Bond dipole (the polarity of individual bonds) can be viewed within a molecule as vector quantity that has magnitude and direction (can be added and subtracted) Perfectly symmetric molecules: have a zero dipole moment and are considered as nonpolar (see Table 10.2)

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Dipole Moment and Molecular Geometry However, molecules that exhibit any asymmetry in the arrangement of electron pairs would have a nonzero dipole moment. These molecules are considered polar.

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Dipole Moment and Molecular Geometry

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Determining the polarity of the molecule To determine whether a molecule is polar or non polar 1. Draw the Lewis structure of the molecule and determine molecular geometry 2. Draw the bond dipoles by vector arrows pointing toward more electronegative atoms 3. Determine whether the bond dipoles add together to for net dipole moment

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Molecular Polarity Affects Solubility in Water Polar molecules are attracted to other polar molecules. Since water is a polar molecule, other polar molecules dissolve well in water. and ionic compounds as well

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Effect of polarity on Molecular properties Cis -1,2-dichloroethene

trans-1,2-dichloroethene

boiling point: 60 °C

48°C

Why the boiling point of cis-1,2-dichloroethene is higher than trans-1,2-dichloroethene? Chem 1011


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Valence Bond Theory Valence bond theory:Valence electrons in a molecule reside in atomic orbitals ( s,p,d,f or hybrid orbitals) A bond forms between two atoms when Two half-filled atomic orbitals “overlap” (or less commonly the overlap of completely filled orbital with an empty orbital) Greater the orbital overlap, the stronger the bond. The total number of electrons in both orbitals is no more than two. The shape of the molecule is determined by geometry of the overlapping orbitals For orbitals other than s orbital, gives a direction to the bond.

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Valence bond model for H2 and F2 Generally number of covalent bonds formed by an atom is determined by the number of unpaired valence electrons in the orbitals e.g. F2,O2

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Valence Bond Theory: Hybridization of Atomic Orbitals Unhybridized C orbitals predict the wrong bonding and geometry:


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Hybridization in carbon

Energy

The bonding in carbon might be explained as follows: Four unpaired electrons are formed as an electron from the 2s orbital is promoted (excited) to the vacant 2p orbital.

2p

2p

2s

2s 1s

1s

Chapter 10: Chemical bonding II A.

C atom (ground state)Ghumman

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C atom (promoted) 35

Hbyridization This does not explain the fact that the four bonds in CH4 appear to be identical Hybridization is a mathematical procedure of mixing different types of orbitals to make a new set of degenerate orbitals (orbitals of equal energy) sp, sp2, sp3, sp3d, sp3d2 Some atoms hybridize their orbitals to maximize bonding. more bonds = more full orbitals = more stability Same type of atom can have different hybridizations depending on the compound. C = sp, sp2, sp3

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Hybrid Orbitals Hybrid orbitals- used to describe bonding obtained by the combinations of atomic orbitals of the isolated atoms. # of atomic orbitals added = # of hybrid orbitals formed The particular kind of hybridization that occurs is the one that yields the lowest overall energy for the molecule. The number and type of standard atomic orbitals combined determine the shape of the hybrid orbitals. In CH4, 1 s + 3 p orbitals = 4 sp3 hybrid orbitals. Chem 1011


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sp3 Hybridization of C • Valence bond theory assumes that the four available atomic orbitals in carbon combine to make four equivalent “hybrid” orbitals.

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sp3 Hybrid Orbitals The four sp3 hybrid orbitals (4sp3) take the shape of a tetrahedron Each sp3 hybrid orbital has two lobes, one larger than the other Four large lobes point toward the corners of a tetrahedron Bonds formed with sp3 orbitals are very strong. The bond angles between hybrid orbitals are 109.5° A tetrahedral arrangement of the electron groups always implies sp3 hybridization. Chem 1011


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Other kinds of hybridization 3sp2 Hybrid orbital = one s + two p trigonal planar 2sp Hybrid orbital = one s + one p linear Only atoms of the elements in the third or lower rows in the periodic table can form five hybrid atomic orbitals (e.g. P and S) 5 sp3d Hybrid orbital = one s + three p + one d trigonal bipyramidal 6 sp3d2 Hybrid orbital = one s + three p + two d octahedral shape

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Multiple Bonds More than one orbital from each bonding atom might overlap to form multiple bonds. A sigma ( ) bond: results when the interacting atomic orbitals point along the axis connecting the two bonding nuclei (head-on overlap). In a

bond electron density is along the bond axis.

A (pi) bond: results from “side-to-side” overlap of parallel “p” orbitals, creating an electron distribution above and below the bond axis. A (pi) bond occurs when two parallel orbitals are still available after strong bonds have formed. Chapter 10: Chemical bonding II A.

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Pi bond

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sp2 Hybridization and double bonds

sp2 Hybridiztion and Double Bond Hydbridization of C in Ethylene (C2H4): One unhybridized p orbital on each carbon is perpendicular to the plane of the sp2 hybrid orbitals This remaining “unhybridized” 2p orbitals on each of the carbon atoms overlap side-to-side forming a bond. C=C double bond: one bond and one bond (In class example)

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Bond Rotation Because orbitals that form the bond point along the internuclear axis, rotation around that bond does not require breaking the interaction between the orbitals. But the orbitals that form the bond interact above and below the internuclear axis, so rotation around the axis requires the breaking of the interaction between the orbitals.

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Isomerism

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Sp Hybridization and Triple bond Bonding in Acetylene (C2H2)- C atom with two electron groups linear molecule with180° bond angle A triple bond between two carbon atoms composed of one bond and two bonds

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sp3d Hybridization Atom with five electron groups around it sp3d :Trigonal bipyramidal shape with bond angles120° and 90° seesaw, T-shape, linear use empty d orbitals from valence shell d orbitals can be used to make bonds

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sp3d2 An atom with six electron groups around it sp3d2 octahedral shape square pyramidal, square planar 90° bond angles empty d orbitals from valence shell are used to form hybrid d orbitals can be used to make bonds

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Hybrid Orbitals and Geometry There is a relationship between the type of hybrid orbitals and the geometric arrangement of those orbitals Thus, if you know the geometric arrangement, you know which hybrid orbitals to use in the bonding description Two e- group around central atom will need two hybrid orbitals One for each bond pair and for each lone pair Table10.3 summarizes the types of hybridization and their spatial arrangements

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Writing hybridization and Bonding Scheme To obtain the bonding description of any atom in a molecule, you proceed as follows: 1. Write the Lewis electron-dot formula for the molecule. 2. Use the VSEPR theory to predict the electron pairs geometry around the atom. 3. From the geometric arrangement of the electron pairs, obtain the hybridization for the central atom (see Table 10.3). 4. Sketch the molecule, beginning with the central atom and its hybrid orbitals .Show overlap with the appropriate orbitals on the terminal atoms 5. Label all bonds using the and p bond notation followed by the type of overlapping Chem 1011


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Practice problem What hybridization would be expected for the central atom in each of the following? a. XeF2, b. ClF4+, c. PCl5, d. IF4-, e. BrF5

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Operational Skills Predicting molecular geometries. Relating dipole moment and molecular geometry. Applying valence bond theory to explain bonding in molecules

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