section 8.13 molecular structure: the vsepr model ▪ covalent bonds are formed by the sharing of...

Section 8.13Molecular Structure: The VSEPR Model

▪ Covalent bonds are formed by the sharing of electrons; orbitals overlap to allow for this sharing.

▪ The mixing of two or more atomic orbitals of an atom forms hybrid orbitals. This process is called hybridization.

▪ Predicting hybridization is easy – just count the total number of steric numbers (domains).

▪ The AP test does not cover hybridization past 4 domains.

Hybridization


▪ The total number of steric numbers (also known as substituents - bonding plus non-bonding groups) is equal to the number of atomic orbitals that participate in the hybrid orbital.

# of substituents (steric numbers)

Hybridization Example Molecule

2 sp CO2

3 sp2 CH3

4 sp3CH4, NH3, H2O

5 sp3d PCl5, I3-

6 sp3d2 SF6


▪ In sp hybridization, only one p orbital is mixed with the s orbital

▪ Example: BeF2 - Steric Number: 2

▪ Electron configuration of Be: 1s22s2

▪ Electron configuration of F:1s22s22p5

The sp hybrid orbital


In the ground state, Be has no unpaired electrons – so how can the Be atom form a covalent bond with a fluorine?

Be obtains an unpaired electron by moving one electron from the 2s orbital to the 2p orbital resulting in two unpaired electrons, one in a 2s orbital and another in a 2p orbital

Be

F

Section 8.13Molecular Structure: The VSEPR Model▪ The Be atom can now form two covalent bonds

with fluorine atoms

▪ Although we would not expect these bonds to be identical (one is in a 2s electron orbital, the other is in a 2p electron orbital), the structure of BeF2 is linear and the bond lengths are identical

▪ The 2s and 2p electrons produced a "hybrid" orbital for both electrons

Be F


▪ In sp2 hybridization, two p orbitals are mixed with the s orbital to generate three new hybrids

▪ Example: BF3

The sp2 hybrid orbital


▪ In sp3 hybridization, all three p orbitals are mixed with the s orbital to generate four new hybrids

▪ Example: CH4

The sp3 hybrid orbital

Section 8.13Molecular Structure: The VSEPR ModelSigma (σ) and Pi Bonds (π)

Single bond 1 sigma bondDouble bond 1 sigma bond and 1 pi bondTriple bond 1 sigma bond and 2 pi bonds

Sigma bond: The first bond made with any other atom• Made from hybridized orbitals • s-s, s-p, or p-p head-on overlap between nucleus• Allows for free rotation Pi bond: Any 2nd or 3rd bond made with any other atom• Made from leftover p orbitals • parallel, sideways p-p overlap, nucleus above or below

overlap• Weaker bond than sigma• Fixed rotation

https://www.youtube.com/watch?v=cPDptc0wUYI





Section 8.13Molecular Structure: The VSEPR ModelSigma (σ) and Pi Bonds (π)

How many σ and π bonds are in the acetic acid (vinegar) molecule CH3COOH?

C

H

H

CH

O

O Hσ bonds = 6 + 1 =

7π bonds = 1


Given the structural formula for propyne:

1. Indicate the hybridization of each carbon atom in the structure above. 2. Indicate the total number of sigma (σ) and pi (π) bonds in the molecule.


Bond length and strength

▪The more electrons that are involved in bonding, the shorter the bond length and the stronger the bond (meaning higher bond energy).

Problem: Is the bond length between the two carbon atoms shorter in C2H6, C2H4, or C2H2. Why?

Problem: Use the bonding model to account for the fact that all the bond lengths in SO3 are identical and are shorter than a sulfur-oxygen single bond.


Problem

A. Draw the Lewis electron-dot structures for CO32-, CO2 and

CO, including resonance structures where appropriate.

Which of the three species has the shortest C-O bond length? Explain the reason for your answer.

Account for the fact that the carbon-oxygen bond length in CO3

2– is greater than the carbon-oxygen bond length in CO2.


Warm Up

▪Draw the following structures according to shape. Use arrows to indicate bond polarity. Arrows point to the more electronegative atom. Name shape and bond angle.

1. CH4 2.BF3 3.CO2 4. HCN 5. CH2Cl2 6. CH2O 7. NH3

Section 8.13Molecular Structure: The VSEPR ModelBond Polarity and Dipole Moments

▪Polar Molecules: Molecules with a somewhat negative end and a somewhat positive end (a dipole moment)

▪Use an arrow to represent a dipole moment.▪Point to the negative charge center with the tail of the arrow indicating the positive center of charge.

Copyright © Cengage Learning. All rights reserved 15


Dipole Moment

16

Bond Polarity and Dipole Moments


No Net Dipole Moment (Dipoles Cancel)


Bond Polarity and Dipole Moments


▪Polar molecules have a permanent dipole. ▪Polar molecules line up in the presence of an electric field; nonpolar molecules do not.

Polar Molecules


▪Molecules that have a polar bond may or may not show a dipole moment – the shape of the molecule must be considered.

Section 8.13Molecular Structure: The VSEPR ModelPolar vs. Nonpolar Molecules

Nonpolar molecules have a symmetrical charge distribution▪Diatomic molecules with the same atoms are nonpolar.

▪Ex: Cl2▪Linear, tetrahedral, trigonal planar shapes must have the same peripheral atoms to be nonpolar

▪Ex: CH4 BF3 CO2


Polar molecules have an asymmetrical charge distribution.

▪Always polar: trigonal pyramidal and bentEx: NH3, H2O

▪Diatomic molecules with different atoms are polar.

Ex: HCl▪Linear, tetrahedral, and trigonal planar shapes have different peripheral atoms to be polar

Ex: HCN, CH2Cl2, CH2O

Polar vs. Nonpolar Molecules


The melting points of polar substances are higher than the melting points of non-polar substances with similar sizes. Concept Check: Which of the following would have the higher boiling point?

Polar vs. Nonpolar Molecules


True or false: A molecule that has polar bonds will always be polar.

-If true, explain why. -If false, provide a counter-example.

Answer: False, a molecule may have polar bonds (like CO2) but the individual dipoles might cancel out so that the net dipole moment is zero.


CONCEPT CHECK!


Let’s Think About It

▪Draw the Lewis structure for CO2.

▪Does CO2 contain polar bonds?

▪Is the molecule polar or nonpolar overall? Why?



True or false: Lone pairs make a molecule polar.

-If true, explain why. -If false, provide a counter-example.

Answer: False, lone pairs do not always make a molecule polar. They might be arranged so that they are symmetrically distributed to minimize repulsions, such as XeF4.


CONCEPT CHECK!


Intramolecular vs. Intermolecular Forces

▪Intramolecular forces: The forces within individual molecules holding it together

(ex: covalent and ionic bonds)▪Intermolecular forces – weak interactions between molecules

Intermolecular forces can cause a condensed state of matter (liquids and solids).

▪Intermolecular forces are stronger in solids, weaker in liquids, and nearly absent in gases


Three types of Intermolecular forces (collectively called van der Waals forces)

▪London Dispersion▪Dipole-Dipole▪Hydrogen Bonding

Johannes van der Waals

http://en.wikipedia.org/wiki/Image:Johannes_Diderik_van_der_Waals.jpg

Section 8.13Molecular Structure: The VSEPR ModelLondon Dispersion

▪The weakest IM force▪Present in all molecules; the only type of IM force present in non-polar substances and Nobel gases.

▪Caused by instantaneous dipoles▪Random movement of electrons can create a momentary nonsymmetrical distribution of charge even in nonpolar molecules


▪ Instantaneous dipoles can induce a short-lived dipole in a neighboring molecule

London Dispersion

Section 8.13Molecular Structure: The VSEPR Model▪ Dispersion force increases as the number of electrons in the

molecule increases. (higher molecular weight, more electrons.) (Ex: CCl4 experiences greater London forces than CH4)

▪ Concept Check: Which Nobel gas would you predict to have the lowest boiling point? Why?

▪ Compare the boiling points of the noble gases:helium - 269°C experience fewer dispersion forcesneon - 246°Cargon - 186°Ckrypton - 152°Cxenon - 108°Cradon - 62°C experience more dispersion forces

London Dispersion


Dipole-Dipole Forces

▪Attraction between molecules with dipole moments (molecules that have a permanent dipole)

▪Occur in addition to dispersion forces▪Molecules orient themselves according to their poles

▪Maximizes (+,-) interactions▪Minimizes (+,+) and (-,-) interactions

Dipole-Dipole

http://www.google.com/imgres?safe=off&biw=983&bih=786&tbm=isch&tbnid=HthdQnn81mztCM:&imgrefurl=http://commons.wikimedia.org/wiki/File:Dipole-dipole-interaction-in-HCl-2D.png&docid=n4apq9bkeGuZoM&imgurl=http://upload.wikimedia.org/wikipedia/commons/5/59/Dipole-dipole-interaction-in-HCl-2D.png&w=1100&h=263&ei=bQGlUpfBOIP9oATniICgAw&zoom=1&ved=1t:3588,r:4,s:0,i:112&iact=rc&page=1&tbnh=76&tbnw=298&start=0&ndsp=11&tx=120&ty=32


▪Strength increases as polarity increases.

▪Stronger than dispersion, but yet only 1% the strength of ionic bonds.


Hydrogen Bonding

▪Special type of Dipole-Dipole

▪The strongest of the Intermolecular forces.


▪Hydrogen is bound to a highly electronegative atom (F, O, N) with a lone pair

▪Important in bonding of molecules such as water and DNA

▪Ex: NH3, H2O, HF can hydrogen bond.

▪Ex: HCN does not, why?

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