Pg. : 453 - 459 (Sect. 10.2, 10.3, 10.4): 2, 3, 4, 5, 30, 31, 32, 34, 42 (Sect. 10.5): 7, 38 (ignore last part of directions), 40 (Sect. 10.6, 10.7): 11, 65, 66 Ch. 10 Problem Assignment (Due test day)
I. Artificial Sweeteners: Fooled by Molecular Shape Artificial sweeteners such as aspartame (Nutrasweet) taste sweet like sugar but do not have the calories of sugar. This is because taste and nutritional value are independent properties. The caloric value of food depends on the amount of energy released when the food is metabolized in the body. Many artificial sweeteners are not even metabolized by the body- they just pass straight though. The taste of a food begins with specialized cells on the tongue that detect different molecules of food. These cells are so sensitive that the tongue can detect one molecule of sugar out of thousands of different molecules in a bite of food. We experience certain tastes when molecules of food fit into a special part of a taste receptor on our tongue, much in the same way a key fits into a lock. Artificial sweeteners “trick” our tongues into tasting sweetness because these molecules mimic the molecular shape of the sugar molecule and fit snuggly into our taste receptors.
- used to determine the 3- D shape of a molecule - based on the idea that electrons groups (bonds & lone pair e ─ ) repel one another to varying degrees - the combination of repulsions on the central atom of a molecule determines its 3-D shape VSEPR THEORY (VALENCE SHELL ELECTRON PAIR REPULSION): II. VSEPR Theory: The Five Basic Shapes *The basic idea behind VSEPR is that repulsions between electron groups determine molecular geometry. The preferred geometry is the one in which the electron groups have the maximum separation (and therefore the minimum energy) possible.
Preview We will first look at molecular geometries where there are two to six electron groups around a central atom and all of the electron groups are bonding groups (single or double bonds). We will then look at what happens to the shapes when some of the electron groups become lone pair electrons.
A. T WO E LECTRONS G ROUPS : L INEAR G EOMETRY Bond Angles: 180° Other examples: SiO 2 CS 2 Si OO C SS 16
B.T HREE E LECTRONS G ROUPS : T RIGONAL P LANAR G EOMETRY Bond Angles: 120° Other examples: BH 3 SO 3 B H HH S O OO In CH 2 O, the bond angles deviate slightly from 120° because there is more electron density in a double bond than in a single bond. Different electron groups repel each other in slightly different ways. 6 24
C. F OUR E LECTRONS G ROUPS : T ETRAHEDRAL G EOMETRY Bond Angles: 109.5° Other examples: SiCl 4 CF 4 Si Cl C F F F F 32
D. F IVE E LECTRONS G ROUPS : T RIGONAL B IPYRAMIDAL G EOMETRY Bond Angles: 90°, 120° Other examples: SOF 4 C l O 2 F 3 S O F F F F 40 Cl O F O F F
E. S IX E LECTRONS G ROUPS : O CTAHEDRAL G EOMETRY Bond Angles: 90° Other examples: SCl 6 S Cl 48
III. VSEPR Theory: The Effect of Lone PairsPreview We will now look at molecular geometries where there are four to six electron groups around a central atom and some of the electron groups are lone pairs. Keep in mind that these shapes are all variations of the geometries from the last section, but now one or more bonding groups have been replaced with lone pairs. The difference in shapes results from the fact that lone pair electrons repel other lone pair and bonding electrons to a greater extent than bonding electrons repel one another. THE ORDER OF ELECTRON PAIR REPULSION: lone pair - lone pair > lone pair- bonding pair > bonding pair- bonding pair Most repulsive Least repulsive * Want to be as far away from each other as possible
In this section, don’t get confused between electron and molecular geometries. In the last section, the electron and molecular geometries were the same (same name for each). In this section, they are different. There are only 5 electron geometries, but there are 11 molecular geometries. THE DIFFERENCE BETWEEN: Electron Geometry: Molecular Geometry: electron groups The geometrical arrangement of electron groups atoms The geometrical arrangement of the atoms
A.F OUR E LECTRON G ROUPS WITH L ONE P AIRS T ETRAHEDRAL E LECTRON G EOMETRY (T ETRAHEDRAL E LECTRON G EOMETRY ) Bond Angles: 107° T RIGONAL P YRAMIDAL (M OLEC. G EO.): O NE L ONE P AIR Other examples: PF 3 NCl 3 P F F F N Cl 26
Bond Angles: 104.5° B ENT (M OLEC. G EO.): T WO L ONE P AIRS Other examples: SCl 2 H 2 S A.F OUR E LECTRON G ROUPS WITH L ONE P AIRS T ETRAHEDRAL E LECTRON G EOMETRY (T ETRAHEDRAL E LECTRON G EOMETRY ) S Cl S H H 18 8
B. F IVE E LECTRON G ROUPS WITH L ONE P AIRS B IPYRAMIDAL E LECTRON G EOMETRY (B IPYRAMIDAL E LECTRON G EOMETRY ) Bond Angles: 90°, 120° S EESAW (M OLEC. G EO.): O NE L ONE P AIR Other examples: SeC l 4 IOF 3 Se Cl I O F F F 34 Lone pair must go equatorial 34
B. F IVE E LECTRON G ROUPS WITH L ONE P AIRS B IPYRAMIDAL E LECTRON G EOMETRY (B IPYRAMIDAL E LECTRON G EOMETRY ) Bond Angles: 90°, 120° Other examples: C l F 3 I 3 - Cl F F F I I I 28 Lone pairs must go equatorial
C. S IX E LECTRON G ROUPS WITH L ONE P AIRS (O CTAHEDRAL E LECTRON G EOMETRY ) Bond Angles: 90° S QUARE P YRAMIDAL (M OLEC. G EO.): O NE L ONE P AIR Other examples: XeOF 4 BrF 4 - Xe O F F FF Br F F FF S QUARE P LANAR (M OLEC. G EO.): T WO L ONE P AIRS Lone pairs must go axial 4236
Summary: There are 5 ELECTRON Geometries: LINEAR : 2 e - groups around central atom TRIGONAL PLANAR : 3 e - groups around central atom TETRAHEDRAL : 4 e - groups around central atom TRIGONAL BIPYRAMIDAL : 5 e - groups around central atom OCTAHEDRAL : 6 e - groups around central atom These electron geometries are also the MOLECULAR geometries For molecules where all e - groups are bonding groups.
Summary: There are 6 additional MOLECULAR Geometries. These occur when one or more of the e - groups are lone-pair e - Bent : 3 e - groups around central atom- 1 is a lone-pair 4 e - groups around central atom- 2 are a lone pair Trigonal Pyramidal : 4 e - groups around central atom- 1 is a lone-pair Seesaw: 5 e - groups around central atom- 1 is a lone-pair
T-Shaped: 5 e - groups around central atom- 2 are a lone-pair Square Pyramidal: 6 e - groups around central atom- 1 is a lone-pair Square Planar: 6 e - groups around central atom- 2 are a lone-pair
IV. VSEPR Theory: Predicting Molecular Geometries Procedure: 1. Draw a Lewis structure for the molecule & total the valence e ─ 2. Determine the total number of electron groups around the central atom: lone pairs, single, double, and triple bonds each count as one electron group 3. Determine the number of (1) bonding groups and (2) the number of lone pair groups around the central atom. These should total the sum from Step 2 4. Use the VSEPR Table to determine the electron geometry and the molecular geometry Use pgs. 414-415 in your text to fill out the table
V. Molecular Shape and Polarity To Determine if a Molecule is Polar: 1. Draw the molecule with the correct molecular geometry 2. Determine if each bond in the molecule is polar by considering electronegativity differences. If the bond is polar ( ≥ 0.4 ∆EN), draw an arrow over it pointing toward the more EN atom 3. Determine whether the polar bonds add together to form a net dipole moment, which makes the molecule polar (See Table 10.2)
Dipoles do NOT cancel; molecule is polar OCO Dipoles cancel; molecule is nonpolar H O H Dipoles do NOT cancel; molecule is polar Dipoles do NOT cancel; molecule is polar H N H H H C H H F F: 4.0 C: 2.5 H: 2.1 A B C D E Always polar Others become polar when bonds are NOT identical
In Lewis Theory, we use “dots” to represent valence electrons. Although this theory helps us to understand bonding, it is not an actual representation of how bonding between atoms truly occurs. In the next three sections we will learn about Valence Bond Theory and Molecular Orbital Theory. These theories attempt to explain how bonding actually occurs at the atomic level. VI. Valence Bond Theory e ─ reside in atomic orbitals and/or hybridized atomic orbitals bonds result when half-filled orbitals overlap bonds are localized between atoms VALENCE BOND THEORY:
Let’s apply the general concepts of valence bond theory to under- Stand the bonding in H 2 S:
VII. Valence Bond Theory: Hybridization of Orbitals Although the overlap of half-filled orbitals adequately explains the bonding in some molecules such as H 2 S, it cannot explain the bonding in many other molecules. Let’s see how this idea doesn’t predict the bonding of C and H in CH 4, a stable molecule that we know forms. Because C only has two half-filled orbitals, this model would suggest that carbon only bonds with 2 H’s. Again, we know it bonds with 4.
Some General Statements Regarding Hybridization: The number of standard atomic orbitals added together always equals the number of hybrid orbitals formed. In order words, the numbers of orbitals is preserved. The combination of orbitals determines the shapes and energies of the orbitals formed. The specific type of hybridization that forms for a molecule is the type that lowers the overall energy of the molecule. To better understand bonding in Valence Bond Theory, we must consider hybridization. a mathematical procedure in which atomic orbitals are combined to form new atomic orbitals called hybrid orbitals HYBRIDIZATION:
One s orbital and three p orbitals hybridize to form four sp 3 hybrid orbitals of equal energy. Head-on overlap of sp 3 orbitals only forms single bonds called σ (sigma) bonds. A. SP 3 HYBRIDIZATION
Now we get a better idea of how CH 4 bonds. Using sp 3 hybrid orbitals, C now has four half-filled orbitals to overlap with 4 H 1s orbitals, each of which is half-filled. Remember: according to valence bond theory, bonds form when half – filled orbitals overlap
One s orbital and two p orbitals hybridize to form three sp 2 hybrid orbitals of equal energy. Head-on overlap of sp 2 orbitals only forms single bonds called σ (sigma) bonds. B. SP 2 HYBRIDIZATION One p orbital remains unhybridized used to form a double bond Sideways overlap of two unhybridized p orbitals forms pi (π) bonds.
A DOUBLE BOND CONSISTS OF ONE SIGMA AND ONE PI BOND - Double bonds consist of 1 sigma σ and 1 pi π bond. - A σ bond forms from head-on overlap of two orbitals. - A π bond forms from sideways overlap of two p orbitals.
One s orbital and one p orbital hybridize to form two sp hybrid orbitals of equal energy. Head-on overlap of sp orbitals only forms single bonds called σ (sigma) bonds. C. SP HYBRIDIZATION Two p orbitals remains unhybridized used to form triple bonds Sideways overlap of four unhybridized p orbitals forms 2 π bonds.
DIFFERENCE IN BETWEEN DOUBLE AND TRIPLE BONDS: A double bond contains one sigma and one pi bond A triple bond contains one sigma and two pi bonds Double bond Triple bond
Summary of bond types: Bond Type# of σ# of π Single Double Triple 1 0 1 1 2
Note: Each of the following count as ONE electron group: A lone pair of e −, a single bond, a double bond, or a triple bond *An “interior atom” is an atom bonded to more than one other atom Number of electron groups around interior* atom Hybridization type Electron geometry around interior* atom HOW TO DETERMINE HYBRIDIZATION and MOLECULAR GEOMETRY for A MOLECULE: 4 sp 3 tetrahedral 3 sp 2 trigonal planar 2 sp linear
Ex. Determine the hybridization type of the interior atoms in each of the following molecules: C SS S O OO C F F F F
VIII. Molecular Orbital Theory Molecular orbital theory is different from valence bond theory but also attempts to explain how compounds actually bond at the atomic level. In valence bond theory, we treated orbitals, whether hybrid- ized or not, as being centered on individual atoms. In molecular orbital theory, we will treat orbitals as overlapping the entire molecule, not as “belonging” to individual atoms. - e - reside in orbitals that span the entire molecule - molecular orbitals are made from atomic orbitals MOLECULAR ORBITAL THEORY: To form our molecular orbitals, we will take a Linear Combination of Atomic Orbitals (LCAO). This is a mathematical procedure that is often performed via computers. As in valence bond theory, when atomic orbitals are combined to form molecular orbitals, the total number of orbitals does not change.
Molecular orbital theory is a complicated theory that you can understand at its basic level. It can help us to understand if bonds will form between atoms and which type (double, triple, etc) Because we can often come to the same conclusion via simpler theories, and because of time constraints, we will not study this theory in more detail at this point. The key is to realize that more complicated bonding theories do exist and they can more accurately explain and predict how bonding occurs
Let’s first examine how to take a LCAO for s orbitals: Whenever two atomic orbitals overlap, we get: a bonding molecular orbital which is lower in energy than the atomic orbitals an antibonding molecular orbital which is higher in energy than the atomic orbitals each orbital can hold up to two electrons
We can write molecular orbital energy diagrams to understand bonding in molecules. Here is a molecular orbital diagram for H 2 : Only valence e - go here
He atom Lets draw a molecular orbital diagram for He 2 :
We know that He 2 does NOT form. We can see why by looking at the bond order: Bond Order = A positive bond order means there are more electrons in bonding than in antibonding orbitals and the overall energy of the electrons is lower in the molecule than in the individual atoms A negative or zero bond order indicates that the energy of the electrons is overall higher than it would be if the atoms did not no bond forms bond, and therefore no bond forms. (# of e - in bonding MO’s) - (# of e - in antibonding MO’s) 2
Now let’s examine how to take a LCAO for p orbitals to understand how they form molecular orbitals. Remember that p orbitals exist in sets of three. When three p orbitals from each atom combine, we get six molecular orbitals. Three are bonding and three are anitbonding. Of the three bonding orbitals, one overlaps head-on and form a sigma bond and two overlap sideways to form two pi bonds. Sigma bonding and antibonding molecular orbitals From p atomic orbitals
Two pi bonding and antibonding molecular orbitals made from four p atomic orbitals
Here’s the part that’s slightly tricky: For B, C, and N the π MO is lower in energy than the σ MO. The π* MO is lower in energy than the σ* MO. For O, F, and Ne, it’s opposite. The σ MO is lower in energy than the π MO are lower in energy. BUT the π* MO is lower in energy than the σ* MO.
DISCLAIMER: MO theory is a very useful method that adequately explains bonding and other chemical and physical properties better than most other bonding theories combined. With the aid of computers, MO theory can be used to calculate molecular orbitals for entire molecules (as seen below for some simple examples). Keep in mind that we are just touching on the very basics of MO theory here just so that you have been introduced to the concept.
Neither Lewis Theory nor Valence Bond Theory adequately conveys the fact that all of the bonds in ozone (O 3 ) are equal in length and energy. However, a computer-generated molecular orbital is able to clearly show that all O-O bonds are equivalent.