CHEM 106 Lab 6

pdf

School

CUNY Hunter College *

*We aren’t endorsed by this school

Course

106.LB

Subject

Chemistry

Date

Apr 3, 2024

Type

pdf

Pages

Uploaded by PresidentInternet12837

3/29/22 Experiment 6: VSEPR and Molecular Shape “How Does It Look” Observations and Experimental Data for Part 1 Molecule Name Shape Bond Angle Predicted Bond Angle PF6- Octahedral 89.8 90 [BrF6]- Octahedral 91.8 90 I3- Linear 180 180 ln(CH3)3 Trigonal Planar 116.2 120 [BeF4]2- Tetrahedral 109.7 109.5 NH4+ Tetrahedral 110.3 109.5 [SbF6]- Octahedral 89.9 90 Data for Part 2 Molecule Name Shape XeF5- Square Planar H2O Bent [ClF4]- Square Planar SbBr52- Square Pyramidal Data for Part 3 Molecules Predicted Geometry Actual Crystal Structure Predicted Angles Actual Angles Di-bromodimethylsel enium Seesaw Seesaw 180, 90 177.7, 98º SO2 Bent Bent 120 113.0 NH3 Trigonal pyramidal Trigonal pyramidal 109.5 102.9 Dichloro-diphenyl-se Tetrahedral Seesaw 180, 90 175, 100

lenium Boric acid Trigonal planar Trigonal planar 120 120.7 Focus Question 1. What is VSEPR theory, and how can it be used to predict molecular shapes? The VSEPR theory is used to predict the 3D structure of a molecule by analyzing the number of electron groups around the central atom. There are two components of this theory: electron geometry and molecular geometry. Electron geometry only considers the electron groups around the center, including lone pairs as a group, and gives an overall shape for the molecule. Molecular geometry also looks at the electron group around the center but it considers the impact lone pairs can have on a molecular structure. For structures with a lone pair, they have different names from their electron geometry and a lower expected bond angle. 2. a) Can the structure of simple molecular substances be illustrated by drawing or building models? Why are why not? This is possible because when given a simple molecular substance and its molecular formula, its Lewis structure can be drawn, which can help us to determine the number of electron groups around the central atom. This information can then be used to find the molecular geometry, which can be easily drawn. These structures can also be built as shown in this experiment. b) What about more complex molecules? Explain. Complex molecules are harder to draw since their molecular formula tends to have more atoms to each element. For instance, in the experiment, di-bromodimethylselenium and dichloro-diphenyl-selenium were given. From the compound name, it can be hard to know the chemical formula. In drawing the Lewis structure, it gets more complicated to draw an accurate structure. However, these molecules can be built on a program performed in this experiment. Hence, it isn’t impossible to draw and predict the molecular structure, but it can be difficult without the help of a program. 3. How are models and theories useful in helping explain the structure and behavior of matter? Models and theories like the VSEPR theory can help predict the 3D structure of a molecule – its shape, bond angle, and polarity. A molecule with a greater amount of lone pairs will decrease in bond angle and length, impacting the way the molecule looks. It can also show how atoms are positioned, helping us see the molecule’s shape and polarity – whether there is symmetry or if the dipole moments cancel out. Post-Lab Questions 1. What are molecular geometries and how do they differ from electron domain geometry Molecular geometry predicts a molecule’s structure by looking at the number of electron groups around the central atom and the impact lone pairs from the center can have on the overall structure. This is seen by the many names and shapes of molecules with different amounts of lone pairs. For instance, a molecule with 5 electron groups and 1 lone pair is a seesaw structure

while a molecule with the same amount of electron groups but with 2 lone pairs is a t-shape structure. Electron geometry does not consider the impact of lone pairs and solely looks at the number of electron groups around the central atom. 2. What is the electronic geometry about a central atom which has the following number of regions of electron density? Give one example of each. a. Three regions of electron density Trigonal Planar. An example is sulfur dioxide. b. Four regions of electron density Tetrahedral. An example is methane. c. Five regions of electron density Trigonal bipyramidal. An example is sulfur tetrafluoride. 3. Is PH3 molecule symmetrical or unsymmetrical? What about BI3? Explain your reasoning for both. PH3 is not symmetrical because its molecular shape is trigonal pyramidal – a tetrahedral electron structure with 1 lone pair. Although lone pairs aren’t shown in molecular structures, it is still a part of the molecule and affects the overall symmetry of the molecule. It can decrease the bond angle and length of atoms in the structure. BI3 is symmetrical because it is a trigonal planar with no lone pairs. Boron is an exception to the octet rule and only needs three bonds. Since it does not need a lone pair, the structure remains symmetrical. 4. Keeping the VSEPR model in mind, draw the Lewis structure for ethanol, (C2H5OH) and dimethyl ether (CH3OCH3). Determine the electron pair geometry and molecular geometry around the oxygen atom in each. Ethanol Electron Groups: 4 Electron geometry: Tetrahedral Molecular geometry: Bent Dimethyl Ether Electron Groups: 4 Electron geometry: Tetrahedral Molecular geometry: Bent 5. Cubane C8H8 is a cubic-shaped hydrocarbon with a carbon atom at each corner of the cube. a. According to the VSEPR theory, what should be the shape around each carbon atom? What bond angle is associated with this shape? The VSEPR theory would predict a tetrahedral shape around each carbon with a bond angle of 109.5º.

Your preview ends here

Eager to read complete document? Join bartleby learn and gain access to the full version

Access to all documents
Unlimited textbook solutions
24/7 expert homework help

b. If you assume an ideal cubic shape, what are the actual bond angles around each carbon? The bond angles would be 90º around each carbon. c. Explain how your answer to questions 5A and 5B suggest why this molecule is so unstable. Since cubane has a tetrahedral shape around each carbon, it is supposed to have a bond angle of 109.5º. However, as shown in the given picture and mentioned in 5B, its actual bond angle is closer to 90º. Since the actual bond angle is far from the expected value for structures with four electron groups, it could explain why the molecule is unstable. 6. What are the steps to determine the shape of a molecule given the formula? Explain and show a few examples. In order to find the molecular or electron geometry, the Lewis structure for the molecule must be drawn to determine the number of electron groups around the central atom – this includes the lone pairs.To draw a Lewis structure, you need three pieces of information: total number of electrons (with octet and duet rule), total electrons used in the structure, and number of lone pairs. The total electrons of a molecule (unshared) can be found by adding the total amount of electrons for each element. Most elements follow the octet rule except for hydrogen (2 electrons) and boron (6 electrons). Next, you find the total number of electrons in the structure by looking at the group an element is in on the periodic table, and multiplying the amount appropriately to the number of atoms it contains. After doing this to each element, you add them all together to get the electrons in the structure – this includes bonded and lone electrons. To find the lone pairs of the center, you subtract the first piece of information we found (electrons with the octet/duet rule) by the electrons used. Using these pieces of information, you can construct your Lewis structure. After drawing the structure, you count the total amount of electron groups around the center and determine the electron geometry. Linear– 2 groups, trigonal planar– 3 groups, tetrahedral– 4 groups, trigonal bipyramidal– 5 groups, octahedral– 6. To find the molecular structure, you find the electron geometry and count the number of lone pairs around the atom. Each electron geometrical shape has a different structure depending on the number of lone pairs. * Example: H2O Total e-: 8+2(2)= 12 e- Electron groups: 4 E- in structure: 6+2(1)= 8 e- Electron geometry: tetrahedral Lone pairs: 12-8= 4 e- (2 lp) Molecular geometry: bent *You would determine the name from this list: Lone pair= lp Trigonal planar 1 lp– bent Tetrahedral 1 lp – trigonal pyramidal, 2lp– bent Trigonal bipyramidal 1 lp– seesaw, 2 lp– T-shape, 3 lp– linear Octahedral 1 lp– square pyramidal, 2 lp– square planar, 3 lp– T-shape, 4 lp– linear