Hybridization Intermolecular Forces Chapter 6 Section 5 Molecular Geometry Objectives Explain VSEPR theory. Predict the shapes of molecules or polyatomic ions using VSEPR theory. Explain how the shapes of molecules are accounted for by hybridization theory.
Chapter 6 Section 5 Molecular Geometry Objectives, continued Describe dipole-dipole forces, hydrogen bonding, induced dipoles, and London dispersion forces and their effects on properties such as boiling and melting points. Explain the shapes of molecules or polyatomic ions using VSEPR theory. Chapter 6
Section 5 Molecular Geometry Molecular Geometry The properties of molecules depend not only on the bonding of atoms but also on molecular geometry: the threedimensional arrangement of a molecules atoms. The polarity of each bond, along with the geometry of the molecule, determines molecular polarity, or the uneven distribution of molecular shape. Molecular polarity strongly influences the forces that act between molecules in liquids and solids. A chemical formula, by itself, reveals little information about a molecules geometry. Chapter 6
Section 5 Molecular Geometry VSEPR Theory As shown at right, diatomic molecules, like those of (a) hydrogen, H2, 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. Chapter 6 Section 5 Molecular Geometry VSEPR Theory The abbreviation VSEPR (say it VES-pur) stands for valence-shell electron-pair repulsion. VSEPR theory states that repulsion between the sets of valence-level electrons surrounding an atom causes these sets to be oriented as far apart as possible. example: BeF2 The central beryllium atom is surrounded by only the two
electron pairs it shares with the fluorine atoms. According to VSEPR, the shared pairs will be as far away from each other as possible, so the bonds to fluorine will be 180 apart from each other. The molecule will therefore be linear: F Be F Chapter 6 Section 5 Molecular Geometry VSEPR Theory Representing the central atom in a molecule by A and
the atoms bonded to the central atom by B, then according to VSEPR theory, BeF2 is an example of an AB2 molecule, which is linear. In an AB3 molecule, the three AB bonds stay farthest apart by pointing to the corners of an equilateral triangle, giving 120 angles between the bonds. In an AB4 molecule, the distance between electron pairs is maximized if each AB bond points to one of four corners of a tetrahedron. Chapter 6 Section 5 Molecular Geometry
VSEPR Theory, continued Sample Problem E Use VSEPR theory to predict the molecular geometry of boron trichloride, BCl3. Chapter 6 Section 5 Molecular Geometry VSEPR Theory, continued Sample Problem E Solution First write the Lewis structure for BCl3. Boron is in Group 13 and has three valence electrons.
B Chlorine is in Group 17, so each chlorine atom has seven valence electrons. Cl Chapter 6 Section 5 Molecular Geometry VSEPR Theory, continued Sample Problem E Solution, continued The total number of electrons is calculated as shown below.
B 3Cl 1 3e 3 7e = 3e = 21e 24e The following Lewis structure uses all 24 electrons. Cl
Cl B Cl Chapter 6 Section 5 Molecular Geometry VSEPR Theory, continued Sample Problem E Solution, continued Cl Cl B
Cl Boron trichloride is an AB3 type of molecule. Its geometry should therefore be trigonal-planar. Chapter 6 Section 5 Molecular Geometry VSEPR Theory, continued VSEPR theory can also account for the geometries of molecules with unshared electron pairs. examples: ammonia, NH3, and water, H2O.
The Lewis structure of ammonia shows that the central nitrogen atom has an unshared electron pair: HNH H VSEPR theory postulates that the lone pair occupies space around the nitrogen atom just as the bonding pairs do. Chapter 6 Section 5 Molecular Geometry VSEPR Theory, continued
Taking into account its unshared electron pair, NH3 takes a tetrahedral shape, as in an AB4 molecule. The shape of a molecule refers to the positions of atoms only. The geometry of an ammonia molecule is that of a pyramid with a triangular base. H2O has two unshared pairs, and its molecular geometry takes the shape of a bent, or angular, molecule. Chapter 6 Section 5 Molecular Geometry
VSEPR Theory, continued Unshared electron pairs repel other electron pairs more strongly than bonding pairs do. This is why the bond angles in ammonia and water are somewhat less than the 109.5 bond angles of a perfectly tetrahedral molecule. Chapter 6 Section 5 Molecular Geometry VSEPR Theory, continued The same basic principles of VSEPR theory that have
been described can be used to determine the geometry of several additional types of molecules, such as AB2E, AB2E2, AB5, and AB6. Treat double and triple bonds the same way as single bonds. Treat polyatomic ions similarly to molecules. The next slide shows several more examples of molecular geometries determined by VSEPR theory. Chapter 6 Section 5 Molecular Geometry
VSEPR and Molecular Geometry Chapter 6 Section 5 Molecular Geometry VSEPR and Molecular Geometry Chapter 6 Section 5 Molecular Geometry VSEPR Theory, continued Sample Problem F
a. Use VSEPR theory to predict the shape of a molecule of carbon dioxide, CO2. b. Use VSEPR theory to predict the shape of a ClO chlorate3 ion, . Chapter 6 Section 5 Molecular Geometry VSEPR Theory, continued Sample Problem F Solution
a. Draw the Lewis structure of carbon dioxide. O C O There are two carbon-oxygen double bonds and no unshared electron pairs on the carbon atom. This is an AB2 molecule, which is linear. Chapter 6
Section 5 Molecular Geometry VSEPR Theory, continued Sample Problem F Solution, continued a. Draw the Lewis structure of the chlorate ion. Cl O O O There are three oxygen atoms bonded to the central chlorine atom, which has an unshared electron pair. This is an AB3E molecule, which is trigonal-pyramidal.
Chapter 6 Section 5 Molecular Geometry Hybridization VSEPR theory is useful for predicting and explaining the shapes of molecules. A step further must be taken to explain how the orbitals of an atom are rearranged when the atom forms covalent bonds. For this purpose, we use the model of hybridization, which is the mixing of two or more atomic orbitals of similar energies on the same atom to produce new hybrid atomic orbitals of equal energies.
Chapter 6 Section 5 Molecular Geometry Hybridization Take the simple example of methane, CH4. The carbon atom has four valence electrons, two in the 2s orbital and two in 2p orbitals. Experiments have determined that a methane molecule is tetrahedral. How does carbon form four equivalent, tetrahedrally arranged, covalent bonds? Recall that s and p orbitals have different shapes. To achieve four equivalent bonds, carbons 2s and three 2p
orbitals hybridize to form four new, identical orbitals called sp3 orbitals. The superscript 3 on the p indicates that there are three p orbitals included in the hybridization. The superscript 1 on the s is left out, like in a chemical formula. Chapter 6 Section 5 Molecular Geometry Hybridization, continued The four (s + p + p + p) hybrid orbitals in the sp3-hybridized methane molecule are equivalent: they all have the same energy, which is greater than that
of the 2s orbital but less than that of the 2p orbitals. Hybrid orbitals are orbitals of equal energy produced by the combination of two or more orbitals on the same atom. Hybridization explains the bonding and geometry of many molecules. Chapter 6 Geometry of Hybrid Orbitals Section 5 Molecular Geometry
Chapter 6 Section 5 Molecular Geometry 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. Chapter 6 Section 5 Molecular Geometry Comparing Ionic and Molecular Substances Chapter 6 Section 5 Molecular Geometry
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. Chapter 6 Section 5 Molecular Geometry
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: H Cl Chapter 6 Section 5 Molecular Geometry 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. Dipole-dipole forces act at short range, only between nearby molecules. Dipole-dipole forces explain, for example the difference between the boiling points of iodine chloride, ICl (97C), and bromine, BrBr (59C). Chapter 6 Section 5 Molecular Geometry
Comparing Dipole-Dipole Forces Chapter 6 Section 5 Molecular Geometry 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, O2, is able to dissolve
in water, a polar molecule. Chapter 6 Section 5 Molecular Geometry 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 positive charge that is almost half as large as that of a bare proton. Chapter 6 Section 5 Molecular Geometry 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. Chapter 6 Section 5 Molecular Geometry 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. Chapter 6 Visual Concepts Hydrogen Bonding Chapter 6 Section 5 Molecular Geometry 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. Chapter 6 Section 5 Molecular Geometry
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. Fritz London first proposed their existence in 1930.
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