Chapter 9 – Molecular Geometry and Bonding Theories Homework: 11, 13, 15, 19, 20, 21, 25, 26, 31, 34, 35, 36, 39, 41, 42, 43, 44, 47, 49, 51, 54, 56, 96, 100

9.2 – The VSEPR Model Two balloons Three balloons Four balloons Linear arrangement Three balloons Trigonal-planar arrangement Four balloons Tetrahedral arrangement

Electrons in molecules behave like balloons A single covalent bond forms between atoms when a pair of electrons is between the atoms A bonding pair of electrons defines a region in which the electrons are most likely to be found between two atoms This area we find electrons is called an electron domain A nonbinding pair (or lone pair) defines an electron domain located around one atom

Example Four electron domains here In general, each nonbinding pair, single bond or multiple bond produces an electron domain around the central atom

Because electron domains are negatively charged, they repel each other. The best arrangement of a given number of electron domains is the one that minimizes the repulsions between them. This is the basic idea behind the VSEPR model.

Similar to Balloons? You bet! Two domains makes linear arrangement Three domains makes trigonal-planar arrangement Four domains makes tetrahedral arrangement

pg. 349

The arrangement of electron domains about the central atom is called its electron–domain geometry. In contrast, the molecular geometry is the arrangement of only the atoms in a molecule or ion So any non-bonding pairs are not a part of the molecular geometry

The VSEPR model predicts electron-domain geometry From this and knowing how many domains are due to nonbinding pairs, we can predict the molecular geometry When all the electron domains in a molecule come from bonds, the molecular geometry is the same as the electron-domain geometry But if one or more domains comes from lone pairs, we must ignore those domains for molecular shape

pg. 351

Example NH3 Already done this. 4 electron domains around central atom So electron-domain geometry is tetrahedral We know 1 of those domains comes from lone pairs So the molecular geometry of NH3 is trigonal pyramidal Tetrahedral with one less end, see pg. 347

Steps using VSEPR model to predict shape of molecules Draw Lewis structure Count number of electron domains around central atom Determine electron-domain geometry Use table 9.1, 9.2 or 9.3 Use the arrangement of the bonded atoms to determine the molecular geometry Use table 9.2 or 9.3

Example CO2 Draw Lewis Structure How many electron domains around the central atom are there?

What is the electron-domain geometry for this? Linear What molecular geometry is possible?

Effect of Nonbonding Electrons and Multiple Bonds on Bond Angles We refine the VSEPR model to predict and explain slight variances from the ideal bond angles Methane (CH4), ammonia (NH3) and water (H2O) all have tetrahedral electron-domain geometries But their bond angles are a little different CH4 = 109.5º, NH3= 107º and H2O = 104.5º Differences based around which type of electron pairs make up the electron domains

Bond angles decrease as the # of nonbonding electron pairs increase. Bonding pair of electrons attracted by both nuclei of the bonded atoms Lone pair of electrons attracted primarily by one nucleus

Since H2O had the most lone pairs, it gets the shortest bond angles Because lone pair has less nuclear attraction, it’s domain becomes more spread out So electron domain for lone pairs exert more repulsive force on adjacent electron domains This compresses (lessens) the bond angles Since H2O had the most lone pairs, it gets the shortest bond angles

Multiple Bonds an Bond Angles Multiple bonds have a higher electron-charge density than single bonds Also creates larger electron domains So electron domains for multiple bonds exert a greater repulsive force on adjacent electron domains than single bonds do So multiple bonds (double or triple) will decrease the bond angles too

Phosgene (Cl2CO) Central atom has three electron domains 3 single bonds Trigonal planar geometry Double bond acts like a lone pair, reducing the Cl-C-Cl bond angle

How Do These all Compare? In terms of volume occupied by electron pairs In other words, who compresses the most? Lone pair > triple bonds > double bonds > single bonds

Molecules with Expanded Valence Shells So far we have assumed the molecules have no more than an octet of electrons But the most common exception to the octet rule is a central atom having greater than 8 valence electrons So we need to deal with molecules with 5 or 6 electron domains

pg. 354

Example Use the VSEPR model to predict the electron and molecular geometry of ClF3 Step 1: Lewis structure How many electron domains around central atom? 5

How many bonding domains? How many non-binding domains? 5 electron domains Gives us an electron geometry of trigonal bipyramidal How many bonding domains? 3 How many non-binding domains? 2 So its molecular geometry is T-shaped

Shapes of Larger Molecules The VSEPR model can be extended to more complex molecules than we’ve been dealing with. Consider acetic acid CH3COOH

Acetic acid has 3 interior atoms Carbon, and each oxygen We can use VSEPR to look at each central atom individually

9.3 – Molecular Shape and Molecular Polarity Remember that bond polarity measures how equally the electrons in a bond are shared between the two atoms Higher bond polarity = less equal sharing Higher electronegativity difference = higher bond polarity

For every bond in the molecule, we can look at the bond dipole The dipole moment depends on both the polarities of the bonds and the geometry of the molecule Last chapter we focused just on the polarity effect on the dipole moment For every bond in the molecule, we can look at the bond dipole The dipole moment that is due ONLY to the two atoms in the bond

Example CO2 O=C=O Each C=O bond is polar (O is more electronegative than C) Since we have two O=C bonds, the bonds are identical We end up with high electron density around the O, and low electron density in the middle

Bond dipoles and dipole moments are vectors The overall dipole moment is the sum of the bond dipoles that make it up But, must consider both the amount of the dipole, and the direction of the dipole We have two identical C=O bonds, so the amount of the dipoles are the same But the DIRECTION of the dipoles are opposite This causes the individual bond dipoles to cancel each other out So the geometry of CO2 indicates that it is a NONPOLAR molecule, even though it contains polar bonds.

Bond Dipole Activity Bond Dipole Activity

Steps to Determine Molecular Polarity Draw Lewis structure Determine molecular geometry Look at effects of electronegativity differences

9.4 – Covalent Bonding and Orbital Overlap The VSEPR gives as a method to predict the shape of molecules Does not explain WHY the bonds exist between atoms A mixture of Lewis’ notion of electron-pair bonds and atomic orbitals leads to a model of chemical bonding This mixture of views is called the valence-bond theory

In Lewis theory, covalent bonding occurs when atoms share electrons The sharing concentrates electron density between the two nuclei involved In valence-bond theory, the build-up of electron density between the nuclei is thought of as occurring when a valence atomic orbital of one atom merges with a valence atomic orbital of another atom

This merger of orbitals Means that they share a region of space Called overlap The overlap of orbitals allows two electrons of opposite spin to share the common space between the nuclei Forming an atomic bond See figure 9.14 on pg. 360

Distance There is always an optimum distance between the two bonded nuclei in a covalent bond Too close = too much repulsion between the nuclei Too far = not much overlap, not a strong bond

9.5 – Hybrid Orbitals