VSEPR Theory As shown at right, diatomic molecules, like those of (a) hydrogen, H 2, and (b) hydrogen chloride, HCl, can only be linear because they consist

VSEPR Theory As shown at right, diatomic molecules, like those of (a) hydrogen, H 2, and (b) hydrogen chloride, HCl, can only be linear because they consist of only two atoms. To predict the geometries of more-complicated molecules, one must consider the locations of all electron pairs surrounding the bonding atoms. This is the basis of VSEPR theory. Section 5 Molecular Geometry Chapter 6

Section 5 Molecular Geometry Chapter 6 Lewis Structures show bonding and lone pairs, but do not denote shape. However, we use Lewis Structures to help us determine shapes. Here we see some common shapes for molecules with two or three atoms connected to a central atom.

Section 5 Molecular Geometry Chapter 6 What Determines the Shape of a Molecule? Simply put, electron pairs, whether they be bonding or nonbonding, repel each other. By assuming the electron pairs are placed as far as possible from each other, we can predict the shape of the molecule. This is the Valence-Shell Electron-Pair Repulsion (VSEPR) model.

Electron Domains We can refer to the directions to which electrons point as electron domains. This is true whether there is one or more electron pairs pointing in that direction. The central atom in this molecule, A, has four electron domains. Chapter 6

Valence-Shell Electron-Pair Repulsion (VSEPR) Model The best arrangement of a given number of electron domains is the one that minimizes the repulsions among them. (The balloon analogy in the figure to the left demonstrates the maximum distances, which minimize repulsions.) Chapter 6

Electron-Domain Geometries The Table shows the electron-domain geometries for two through six electron domains around a central atom. To determine the electron- domain geometry, count the total number of lone pairs, single, double, and triple bonds on the central atom. Chapter 6

Molecular Geometries Once you have determined the electron-domain geometry, use the arrangement of the bonded atoms to determine the molecular geometry. Tables 9.2 and 9.3 show the potential molecular geometries. We will look at each electron domain to see what molecular geometries are possible. Chapter 6

Linear Electron Domain In the linear domain, there is only one molecular geometry: linear. NOTE: If there are only two atoms in the molecule, the molecule will be linear no matter what the electron domain is. Chapter 6

Trigonal Planar Electron Domain There are two molecular geometries: trigonal planar, if all electron domains are bonding, and bent, if one of the domains is a nonbonding pair. Chapter 6

Tetrahedral Electron Domain There are three molecular geometries: tetrahedral, if all are bonding pairs, trigonal pyramidal, if one is a nonbonding pair, and bent, if there are two nonbonding pairs. Chapter 6

Nonbonding Pairs and Bond Angle Nonbonding pairs are physically larger than bonding pairs. Therefore, their repulsions are greater; this tends to compress bond angles. Chapter 6

Multiple Bonds and Bond Angles Double and triple bonds have larger electron domains than single bonds. They exert a greater repulsive force than single bonds, making their bond angles greater. Chapter 6

Expanding beyond the Octet Rule Remember that some elements can break the octet rule and make more than four bonds (or have more than four electron domains). The result is two more possible electron domains: five = trigonal bipyramidal; six = octahedral (as was seen in the slide on electron-domain geometries). Chapter 6

Trigonal Bipyramidal Electron Domain There are two distinct positions in this geometry: Axial Equatorial Lone pairs occupy equatorial positions. Chapter 6

Trigonal Bipyramidal Electron Domain There are four distinct molecular geometries in this domain: Trigonal bipyramidal Seesaw T-shaped Linear Chapter 6

Octahedral Electron Domain All positions are equivalent in the octahedral domain. There are three molecular geometries: Octahedral Square pyramidal Square planar Chapter 6

VSEPR Theory, continued Sample Problem E Use VSEPR theory to predict the molecular geometry of boron trichloride, BCl 3. Section 5 Molecular Geometry Chapter 6

VSEPR Theory, continued Sample Problem E Solution Section 5 Molecular Geometry Chapter 6 Boron trichloride is an AB 3 type of molecule. Its geometry should therefore be trigonal-planar.

Click below to watch the Visual Concept. Visual Concept Chapter 6 Section 5 Molecular Geometry VSEPR and Lone Electron Pairs

Sample Problem F a.Use VSEPR theory to predict the shape of a molecule of carbon dioxide, CO 2. b.Use VSEPR theory to predict the shape of a chlorate ion,. Section 5 Molecular Geometry Chapter 6 VSEPR Theory, continued

Sample Problem F Solution a.Draw the Lewis structure of carbon dioxide. There are two carbon-oxygen double bonds and no unshared electron pairs on the carbon atom. This is an AB 2 molecule, which is linear. Section 5 Molecular Geometry Chapter 6 VSEPR Theory, continued

Sample Problem F Solution, continued b.Draw the Lewis structure of the chlorate ion. There are three oxygen atoms bonded to the central chlorine atom, which has an unshared electron pair. This is an AB 3 E molecule, which is trigonal-pyramidal. Section 5 Molecular Geometry Chapter 6 VSEPR Theory, continued

VSEPR and Hybrid Orbitals VSEPR predicts shapes of molecules very well. How does that fit with orbitals? Lets use H 2 O as an example: If we draw the best Lewis structure to assign VSEPR, it becomes bent. If we look at oxygen, its electron configuration is 1s 2 2s 2 2p 4. If it shares two electrons to fill its valence shell, they should be in 2p. Wouldnt that make the angle 90? Why is it 104.5? Section 5 Molecular Geometry Chapter 6

Hybrid Orbitals Hybrid orbitals form by mixing of atomic orbitals to create new orbitals of equal energy, called degenerate orbitals. When two orbitals mix they create two orbitals; when three orbitals mix, they create three orbitals; etc. Section 5 Molecular Geometry Chapter 6

Besp hybridization When we look at the orbital diagram for beryllium (Be), we see that there are only paired electrons in full sub-levels. Be makes electron deficient compounds with two bonds for Be. Why? sp hybridization (mixing of one s orbital and one p orbital) Section 5 Molecular Geometry Chapter 6

sp Orbitals Mixing the s and p orbitals yields two degenerate orbitals that are hybrids of the two orbitals. These sp hybrid orbitals have two lobes like a p orbital. One of the lobes is larger and more rounded, as is the s orbital. Section 5 Molecular Geometry Chapter 6

Position of sp Orbitals These two degenerate orbitals would align themselves 180 from each other. This is consistent with the observed geometry of Be compounds (like BeF 2 ) and VSEPR: linear. Section 5 Molecular Geometry Chapter 6

BoronThree Electron Domains Gives sp 2 Hybridization Using a similar model for boron leads to three degenerate sp 2 orbitals. Section 5 Molecular Geometry Chapter 6

Carbon: sp 3 Hybridization With carbon, we get four degenerate sp 3 orbitals. Section 5 Molecular Geometry Chapter 6

What Happens with Water? We started this discussion with H 2 O and the angle question: Why is it 104.5 instead of 90? Oxygen has two bonds and two lone pairsfour electron domains. The result is sp 3 hybridization! Section 5 Molecular Geometry Chapter 6

Click below to watch the Visual Concept. Visual Concept Chapter 6 Section 5 Molecular Geometry Hybrid Orbitals

Section 5 Molecular Geometry Chapter 6

Copyright Cengage Learning. All rights reserved 34 Types of Intermolecular Force Strongest to weakest forces: dipoledipole forces Dipole-induced dipole iondipole forces hydrogen bonding (a special dipoledipole force) dispersion forces (or London dispersion forces) Section 5 Molecular Geometry Chapter 6

Intermolecular Forces The forces of attraction between molecules are known as intermolecular forces. The boiling point of a liquid is a good measure of the intermolecular forces between its molecules: the higher the boiling point, the stronger the forces between the molecules. Intermolecular forces vary in strength but are generally weaker than bonds between atoms within molecules, ions in ionic compounds, or metal atoms in solid metals. Boiling points for ionic compounds and metals tend to be much higher than those for molecular substances: forces between molecules are weaker than those between metal atoms or ions. Section 5 Molecular Geometry Chapter 6

Comparing Ionic and Molecular Substances Section 5 Molecular Geometry Chapter 6

Intermolecular Forces, continued The strongest intermolecular forces exist between polar molecules. Because of their uneven charge distribution, polar molecules have dipoles. A dipole is created by equal but opposite charges that are separated by a short distance. The direction of a dipole is from the dipoles positive pole to its negative pole. Section 5 Molecular Geometry Chapter 6

Intermolecular Forces, continued A dipole is represented by an arrow with its head pointing toward the negative pole and a crossed tail at the positive pole. The dipole created by a hydrogen chloride molecule is indicated as follows: Section 5 Molecular Geometry Chapter 6

Section 5 Molecular Geometry Chapter 6

Intermolecular Forces, continued The negative region in one polar molecule attracts the positive region in adjacent molecules. So the molecules all attract each other from opposite sides. Such forces of attraction between polar molecules are known as dipole-dipole forces. Section 5 Molecular Geometry Chapter 6

Intermolecular Forces, continued Dipole-dipole forces act at short range, only between nearby molecules. Rapidly becomes weaker as distance between molecules increases. Dipole-dipole forces explain, for example the difference between the boiling points of iodine chloride, ICl (97C), and bromine, BrBr (59C). Section 5 Molecular Geometry Chapter 6

Comparing Dipole-Dipole Forces Section 5 Molecular Geometry Chapter 6

Click below to watch the Visual Concept. Visual Concept Chapter 6 Section 5 Molecular Geometry Dipole-Dipole Forces

Intermolecular Forces, continued A polar molecule can induce a dipole in a nonpolar molecule by temporarily attracting its electrons. The result is a short-range intermolecular force that is somewhat weaker than the dipole-dipole force. Induced dipoles account for the fact that a nonpolar molecule, oxygen, O 2, is able to dissolve in water, a polar molecule. Section 5 Molecular Geometry Chapter 6

Dipole-induced Dipole Section 5 Molecular Geometry Chapter 6

Click below to watch the Visual Concept. Visual Concept Chapter 6 Section 5 Molecular Geometry Dipole-Induced Dipole Interaction

IonDipole Interactions Iondipole interactions are found in solutions of ions. The strength of these forces is what makes it possible for ionic substances to dissolve in polar solvents. Section 5 Molecular Geometry Chapter 6

Intermolecular Forces, continued Some hydrogen-containing compounds have unusually high boiling points. This is explained by a particularly strong type of dipole-dipole force. In compounds containing HF, HO, or HN bonds, the large electronegativity differences between hydrogen atoms and the atoms they are bonded to make their bonds highly polar. This gives the hydrogen atom a strong positive charge. Section 5 Molecular Geometry Chapter 6

Intermolecular Forces, continued The small size of the hydrogen atom allows the atom to come very close to an unshared pair of electrons in an adjacent molecule. The intermolecular force in which a hydrogen atom that is bonded to a highly electronegative atom is attracted to an unshared pair of electrons of an electronegative atom in a nearby molecule is known as hydrogen bonding. Section 5 Molecular Geometry Chapter 6

Intermolecular Forces Hydrogen bonds are usually represented by dotted lines connecting the hydrogen-bonded hydrogen to the unshared electron pair of the electronegative atom to which it is attracted. An excellent example of hydrogen bonding is that which occurs between water molecules. The strong hydrogen bonding between water molecules accounts for many of waters characteristic properties. Section 5 Molecular Geometry Chapter 6

Intermolecular Forces Section 5 Molecular Geometry Chapter 6

Visual Concepts Hydrogen Bonding Chapter 6

Intermolecular Forces, continued London Dispersion Forces Even noble gas atoms and nonpolar molecules can experience weak intermolecular attraction. In any atom or moleculepolar or nonpolarthe electrons are in continuous motion. As a result, at any instant the electron distribution may be uneven. A momentary uneven charge can create a positive pole at one end of an atom of molecule and a negative pole at the other. Section 5 Molecular Geometry Chapter 6

Intermolecular Forces, continued London Dispersion Forces, continued This temporary dipole can then induce a dipole in an adjacent atom or molecule. The two are held together for an instant by the weak attraction between temporary dipoles. The intermolecular attractions resulting from the constant motion of electrons and the creation of instantaneous dipoles are called London dispersion forces. Section 5 Molecular Geometry Chapter 6

Section 5 Molecular Geometry Chapter 6 Occurs in all molecules, including nonpolar ones. The figure below shows how a nonpolar particle (in this case a helium atom) can be temporarily polarized to allow dispersion force to form. The tendency of an electron cloud to distort is called its polarizability.

Fritz London first proposed their existence in 1930. Fritz London- 1900-1954 Born in Germany, later was a professor at Duke Section 5 Molecular Geometry Chapter 6

Click below to watch the Visual Concept. Visual Concept Chapter 6 Section 5 Molecular Geometry London Dispersion Force

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VSEPR Theory As shown at right, diatomic molecules, like those of (a) hydrogen, H 2, and (b) hydrogen chloride, HCl, can only be linear because they consist