Module in Chemistry

7/29/2019 Module in Chemistry

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Project in Chemistry

Molecular GeometryAnd Bonding Theories

Gilean Carla Dalida

Lodie Jane C. dela Cruz

III Ptolemy


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Molecular Geometry

I. Molecular Shapes

II. The VSEPR model

III. Polarity of Polyatomic Molecules

IV. Covalent bonding and Orbital overlap

V. Hybrid orbital

We have seen that molecules are extremely small. Nevertheless, the size and shape of amolecule of a particular substance together with the strength and polarity of its bonds, largely

determine the properties of that substance. Some of the most dramatic examples of theimportance of size and shape are seen in biochemical reactions.

We will discuss three models that relate to molecular geometry and bonding. First, we considerhow the shapes of molecules can be describes and predicted using a simple model basedlargely on Lewis structures and the idea o electron-electron repulsions (VSEPR model).

We then examine the model of a molecular bonding (the valence-bond theory) that helps usunderstand why molecules form bonds and why they have shapes as they do. Finally, wediscuss a model of chemical bonding (the molecular orbital theory) that provides additionalinsight into the electronic structures of molecules.


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Common Student Misconceptions

Students find it difficult to think in three dimensions. Often, they believe that a square planararrangement is the best arrangement for the least repulsion of four electron domains.

Students often confuse the electron domain geometry and the molecular geometry (shape).

Students need to realize that in order to determine whether a molecule is polar, they need toestablish the correct molecular geometry.

Students need to realize that large molecules with several central atoms do not have easilydescribable molecular shapes; geometry about each central atom has to be determinedindividually.

Students often find it difficult to understand how a molecule with polar bonds can be nonpolar;a review of basic vector algebra to illustrate how the net dipole is derived may be needed.

Students often attempt to determine polarity of ions.

Students do not realize that hybridization is related to the electron domain geometry, not themolecular geometry.

Students need to realize that in wave mechanics bonding orbitals result from constructiveinterference and antibonding orbitals form destructive interference.


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I. Molecular Shapes

Lewis structures give atomic connectivity: they tell us which atoms are physicallyconnected to which, but do not indicate the shapes of molecules. They simply show the numberand types of bonds between atoms.

E.g. the Lewis structure of CCl4 tells us only that four Cl atoms are bonded to a central C atom.

The shape of a molecule is determined by its bond angles.

The angles made by the lines joining the nuclei of the atoms in a molecule are the bondangles.

Consider CCl4:

Experimentally we find all ClCCl bond angles are 109.5o.

Therefore, the molecule cannot be planar.

All Cl atoms are located at the vertices of a tetrahedron with the C at its center.

In order to predict molecular shape, we assume that the valence electrons repel each other.

Therefore, the molecule adopts the three-dimensional geometry that minimizes this repulsion.

We call this model the Valence Shell Electron PairRepulsion (VSEPR) model.


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Derivatives from the ABngeometries.

Additional molecular shapes can be obtained by removing corner atoms from the basicmolecular geometries. Here we begin with a tetrahedron and successively remove corners,producing first a trigonal-pyramidal geometry and then a bent geometry, each having ideal bondangles of 109.5. Molecular shape is meaningful only when there are at least three atoms. Ifthere are only two, they must be arranged next to each other, and no special name is given todescribe the molecule.


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II. The VSEPR Model

A covalent bond forms between two atoms when a pair of electrons occupies the spacebetween the atoms.

This is a bonding pair of electrons.

Such a region is an electron domain.

A nonbonding pair orlone pairof electrons defines an electron domain located principally onone atom.

Example: NH3 has three bonding pairs and one lone pair.

Electron domains.

A bonding pair of electrons defines a region in which the electrons will most likely be found.We will refer to such a region as an electron domain. Likewise, a nonbonding pair (orlonepair) of electrons defines an electron domain that is located principally on one atom. Forexample, the Lewis structure of NH3 has four electron domains around the central nitrogenatom (three bonding pairs and one nonbonding pair). VSEPR predicts that the best arrangementof electron domains is the one that minimizes the repulsions among them.

The arrangement of electron domains about the central atom of an ABn molecule is itselectron-domain geometry.

There are five different electron-domain geometries:

Linear (two electron domains), trigonal planar (three domains), tetrahedral (four domains),trigonal

bipyramidal (five domains) and octahedral (six domains).


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The molecular geometry is the arrangement of the atoms in space.

To determine the shape of a molecule we distinguish between lone pairs and bonding pairs.We use the electron domain geometry to help us predict the molecular geometry.

Draw the Lewis structure.

Count the total number of electron pairs around the central atom.

Arrange the electron pairs in one of the above geometries to minimize electron-electronrepulsion.

Next, determine the three-dimensional structure of the molecule:

We ignore lone pairs in the molecular geometry.

Describe the molecular geometry in terms of the angular arrangement of the bonded atoms.

Multiple bonds are counted as one electron domain.

Sample ExerciseUse the VSEPR model to predict the molecular geometries of

a) O3

b) SnCl3-

Practice ExercisePredict the electron-domain geometry and the molecular geometry for

a) SeCl2

b) CO3


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The Effect of Nonbonding Electrons and Multiple Bonds on Bond Angles

We determine the electron pair geometry only looking at electrons

We name the molecular geometry by the positions of the atoms

We ignore lone pairs in the molecular geometry.

All the atoms that obey the octet rule have tetrahedral electron pair geometries.

We refine VSEPR to predict and explain slight distortions from ideal geometries.

Consider three molecules with tetrahedral electron domain geometries:

CH4, NH3, and H2O.

By experiment, the HXH bond angle decreases from C (109.5o in CH4) to N (107 o in NH3)to O

(104.5 o in H2O).

Since electrons in a bond are attracted by two nuclei, they do not repel as much as lonepairsTherefore,

the bond angle decreases as the number of lone pairs increases.


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A bonding pair of electrons is attracted by two nuclei. They do not repel as much as lone pairswhich are primarily attracted by only one nucleus.

Electron domains for nonbonding electron pairs thus exert greater repulsive forces on adjacentelectron domains.

They tend to compress the bond angles. The bond angle decreases as the number of nonbonding pairs increases.

Similarly, electrons in multiple bonds repel more than electrons in single bonds. (e.g. in Cl2COthe OCCl angle is 124.3 o, and the ClCCl bond angle is 111.4 o).

We will encounter 11 basic molecular shapes:

Three atoms (AB2)

Linear

Bent

Four atoms (AB3):

Trigonal planar

Trigonal pyramidal

T-shaped


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Five atoms (AB4):

Tetrahedral

Square planar

See-saw

Six atoms (AB5):

Trigonal bipyramidal

Square pyramidal

Seven atoms (AB6):

Octahedral

Molecules with Expanded Valence Shells

Atoms that have expanded octets have AB5 (trigonal bipyramidal) or AB6 (octahedral) electrondomain

geometries.

Trigonal bipyramidal structures have a plane containing three electron pairs.

The fourth and fifth electron pairs are located above and below this plane.

In this structure two trigonal pyramids share a base.

For octahedral structures, there is a plane containing four electron pairs.

Similarly, the fifth and sixth electron pairs are located above and below this plane.

Two square pyramids share a base.

Consider a trigonal bipyramid:

The three electron pairs in the plane are called equatorial;

The two electron pairs above and below this plane are called axial;

The axial electron pairs are 180 o apart and 90 o to the equatorial electrons;

The equatorial electron pairs are 120 o apart;

To minimize electronelectron repulsion, nonbonding pairs are always placed in equatorialpositions

and bonding pairs in either axial or equatorial positions.


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III. Molecular Shape and Molecular Polarity

Polar molecules interact with electric fields.

We previously saw that binary compounds are polar if their centers of negative and positivecharge do not

coincide.

If two charges, equal in magnitude and opposite in sign, are separated by a distance d, then adipoleis

established.

The dipole moment, , is given by

= Qr

where Qis the magnitude of the charge

We can extend this to polyatomic molecules.

For each bond in a polyatomic molecule, we can consider the bond dipole.

The dipole moment due only to the two atoms in the bond is the bond dipo le.

Because bond dipoles and dipole moments are vector quantities, the orientation of theseindividual

dipole moments determines whether the molecule has an overall dipole moment.

Examples:

In CO2 each +CO- dipole is canceled because the molecule is linear.

In H2O, the +HO- dipoles do not cancel because the molecule is bent.


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It is possible for a molecule with polar bonds to be either polar or nonpolar.

Example:

For diatomic molecules:

polar bonds always result in an overall dipole moment.

For triatomic molecules:

if the molecular geometry is bent, there is an overall dipole moment.

if the molecular geometry is linear, and the B atoms are the same, there is no overall dipole

moment.

if the molecular geometry is linear and the B atoms are different, there is an overall dipole

moment.

For molecules with four atoms:

if the molecular geometry is trigonal pyramidal, there is an overall dipole moment;

if the molecular geometry is trigonal planar, and the B atoms are identical, there is no

overall dipole moment;

if the molecular geometry is trigonal planar and the B atoms are different, there is an

overall dipole moment.

The overall polarity of a molecule depends on its molecular geometry.


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Molecules containing polar bonds.

Two of these molecules have a zero dipole moment because their bond dipoles cancel oneanother, while the other

molecules are polar.

IV. Covalent Bonding and Orbital Overlap

Lewis structures and VSEPR theory give us the shape and location of electrons in a molecule.

They do not explain why a chemical bond forms.

How can quantum mechanics be used to account for molecular shape? What are the orbitalsthat are

involved in bonding?

We use valence-bond theory:

A covalent bond forms when the orbitals on two a toms overlap.

The shared region of space between the orbitals is called the orbital overlap.


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There are two electrons (usually one from each atom) of opposite spin in the orbital overlap.

As two nuclei approach each other their atomic orbitals overlap.

As the amount of overlap increases, the energy of the interaction decreases.

At some distance the minimum energy is reached.

The minimum energy corresponds to the bonding distance (or bond length).

As the two atoms get closer, their nuclei begin to repel and the energy increases.

At the bonding distance, the attractive forces between nuclei and electrons just balance therepulsive forces (nucleus-nucleus, electron-electron).


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V. Hybrid Orbitals

We can apply the idea of orbital overlap and valence-bond theory to polyatomic molecules.

spHybrid Orbitals

Consider the BeF2 molecule:

Be has a 1s22s2 electron configuration.

There is no unpaired electron available for bonding.

We conclude that the atomic orbitals are not adequate to describe orbitals in molecules.

We know that the FBeF bond angle is 180 o (VSEPR theory).

We also know that one electron from Be is shared with each one of the unpaired electrons

from F.

We assume that the Be orbitals in the BeF bond are 180 o apart.

We could promote an electron from the 2sorbital on Be to the 2porbital to get two unpairedelectrons for bonding.

BUT the geometry is still not explained.


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We can solve the problem by allowing the 2sand one 2porbital on Be to mix or form two newhybrid orbitals (a process called hybridization).

The two equivalent hybrid orbitals that result from mixing an sand a porbital and are called sphybrid orbitals.

The two lobes of an sphybrid orbital are 180 o apart.

According to the valence-bond model, a linear arrangement of electron domains implies sphybridization.

Since only one of 2porbitals of Be has been used in hybridization, there are two unhybridizedporbitals remaining on Be.

The electrons in the sphybrid orbitals form shared electron bonds with the two fluorine atoms.


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Formation of two equivalent BeF bonds in BeF2.

Each sphybrid orbital on Be overlaps with a 2porbital on F to form a bond. The two bonds areequivalent to each other and

form an angle of 180.

sp2 and sp3 Hybrid Orbitals

Important: When we mix natomic orbitals we must get nhybrid orbitals.

Three sp2 hybrid orbitals are formed from hybridization of one sand two porbitals.

Thus, there is one unhybridized porbital remaining.

The large lobes of the sp2 hybrids lie in a trigonal plane.

Molecules with trigonal planar electron-pair geometries have sp2 orbitals on the central atom.


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Foursp3 hybrid orbitals are formed from hybridization of one sand three porbitals.

Therefore, there are four large lobes.

Each lobe points towards the vertex of a tetrahedron.

The angle between the large lobes is 109.5 o

Molecules with tetrahedral electron pair geometries are sp3 hybridized.


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Valence-bond description of H2O.

The bonding in a water molecule can be envisioned as sp3 hybridization of the orbitals on O.Two of the four hybrid orbitals

overlap with 1sorbitals of H to form covalent bonds. The other two hybrid orbitals are occupiedby nonbonding pairs of

electrons.

Hybridization Involving dOrbitals

Since there are only three porbitals, trigonal bipyramidal and octahedral electron-pairgeometries must

involve dorbitals.

Trigonal bipyramidal electron pair geometries require sp3dhybridization.

Octahedral electron pair geometries require sp3d2 hybridization.

Note that the electron pair VSEPR geometry corresponds well with the hybridization.

Use ofdorbitals in making hybrid orbitals corresponds well with the idea of an expanded octet.

Hybrid Orbital Summary

We need to know the electron-domain geometry before we can assign hybridization.

To assign hybridization:

Draw a Lewis structure.


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Assign the electron-domain geometry using VSEPR theory.

Specify the hybridization required to accommodate the electron pairs based on their geometric

arrangement.

Name the geometry by the positions of the atoms.

Bonding in NH3.

The hybrid orbitals used by N in the NH3 molecule are predicted by first drawing the Lewisstructure, then using the VSEPR model to determine the electron-domain geometry, and thenspecifying the hybrid orbitals that correspond to that geometry. This is essentially the sameprocedure as that used to determine molecular structure (Figure 9.6), except we focus on theorbitals used to make bonds and to hold nonbonding pairs.


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Module in Chemistry