The geometrical structure of molecules lab answers

Molecular Modeling and Lewis Structures Hands-On Labs, Inc. Version 42-0080-00-02

Review the safety materials and wear goggles when working with chemicals. Read the entire exercise before you begin. Take time to organize the materials you will need and set aside a safe work space in which to complete the exercise.

Experiment Summary:

In this experiment, you will draw Lewis structures for a series of molecules and then create the VSEPR model for the molecule using the modeling kit.

EXPERIMENT

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Learning Objectives Upon completion of this laboratory, you will be able to:

● Define allotropes, valence electrons, and lone pairs.

● Describe the duet rule and octet rule.

● Define and create Lewis structures.

● Describe the valence shell electron pair repulsion (VSEPR) model.

● Draw Lewis structures for molecules.

● Create VSEPR models of molecules with molecular modeling kits.

● Identify the number of valence electrons of elements using the periodic table.

● Diagram resonance structures.

● Classify the VSEPR model of a molecule as: linear, trigonal planar, tetrahedral, trigonal bipyramidal, or octahedral.

● Assemble molecules with molecular modeling kits, preparing single bonds, double bonds, and resonance structures.

Time Allocation: 3.5 hours

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Experiment Molecular Modeling and Lewis Structures

Materials Student Supplied Materials

Quantity Item Description 1 Camera, digital or smartphone 1 Pen 1 Sheet of white paper

HOL Supplied Materials

Quantity Item Description 1 Modeling kit (Molecular Modeling and Lewis Structures):

6 - Single bonds 4 - Double bonds 18 - Lone pairs 5 - White (1 hole) 1 - Pink (2 holes) 1 - Gray (3 holes) 2 - Red (4 holes) 2 - Black (4 holes) 6 - Green (4 holes) 1 - Blue (4 holes) 4 - Yellow (4 holes) 1 - White (5 holes) 1 - Yellow (6 holes)

Note: To fully and accurately complete all lab exercises, you will need access to:

1. A computer to upload digital camera images.

2. Basic photo editing software, such as Microsoft Word® or PowerPoint®, to add labels, leader lines, or text to digital photos.

3. Subject-specific textbook or appropriate reference resources from lecture content or other suggested resources.

Note: The packaging and/or materials in this LabPaq kit may differ slightly from that which is listed above. For an exact listing of materials, refer to the Contents List included in your LabPaq kit.

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Experiment Molecular Modeling and Lewis Structures

Background Structure and bonding

The structure and molecular bonding of molecules is an important factor in the course of chemical reactions. Both diamond and graphite are variants of the carbon atom, and while both substances are exclusively composed of carbon, it is the difference in the structure and molecular bonding of the carbon atoms that result in the two extraordinarily different minerals. Different structural modifications of an element are referred to as allotropes. See Figure 1.

Figure 1. Allotropes of the element carbon. A. Crystalline structure of diamond. B. Crystalline structure of graphite. © Michael Ströck

Many properties and characteristics are involved in chemical bonding and molecular structure and function including bond strength, polarity, and atomic orbitals. The focus of this lab is the localized electron model. This model assumes that a molecule is bonded through the sharing of valence electron pairs. The localized electron model also highlights valence electron arrangement, Lewis structures, and molecular shape.

Valence electrons and Lewis structures

Valence electrons are the electrons of an atom located in the outermost shell of an atom. The number of valence electrons that an atom (element) contains can be found in the periodic table. Elements in the same group (vertical column) of the periodic table contain the same number of valence electrons. See Figure 2. For example, all elements in group 7 (VIIA), including fluorine, chlorine, bromine, and iodide, have 7 valence electrons. All of the elements in group 6 (VIA), such as oxygen and sulfur, have 6 valence electrons.

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Experiment Molecular Modeling and Lewis Structures

Figure 2. Periodic Table of Elements. The number of valence electrons of an atom is determined by the group number, shown directly above each group in roman numerals. Click to Download

Printable Version.

For example, all elements in group 6 (VIA); oxygen, sulfur, chlorine, bromine, iodide, and astatine, contain 6 valence electrons. Likewise, all elements in group 4 (IVA); carbon, silicon, germanium, tin, and lead, contain 4 valence electrons. See Figure 3.

Figure 3. Valence electrons examples.

Molecules are surrounded by an electron cloud: the electrons belong to the entire molecule, rather than the individual atoms. It is useful, however, to model atoms and electrons in an organized manner to better understand the structure of a molecule.

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Experiment Molecular Modeling and Lewis Structures

http://holscience.com/sites/default/files/Periodic_Table.pdf
http://holscience.com/sites/default/files/Periodic_Table.pdf
A Lewis structure shows how valence electrons are arranged among atoms in a molecule. In arranging valence electrons, the duet and octet rules are very important. The duet rule applies to molecules containing hydrogen, as hydrogen is most stable when sharing two valence electrons. The octet rule is based upon the observation that atoms (other than hydrogen) are most stable when surrounded by eight valence electrons. These 8 valence electrons can either be shared (bonds) or not shared (lone pairs).

Consider the bonding of hydrogen and fluorine. Hydrogen (Group IA) has 1 valence electron and fluorine (Group VIIA) has 7 valence electrons. See Figure 4. The atoms form the molecule hydrogen fluoride (HF), which has a total of 8 valence electrons in its electron cloud. The hydrogen follows the duet rule, and the fluorine follows the octet rule. Two electrons are shared between them. Six electrons surrounding fluorine are not shared, and are considered lone pairs. There are a total of three lone pairs around fluorine, which are shown as three pairs of dots.

Figure 4. Bonding of hydrogen fluoride (HF). Hydrogen has two valence electrons and obeys the duet rule; fluorine has eight valence electrons and obeys the octet rule.

The Lewis structure of the molecule can be drawn so that a single dash represents the shared electrons, as shown in Figure 5.

Figure 5. A. Hydrogen and fluorine share one pair of electrons (shown in the circle); there are also three lone pairs. B. The shared pair may be represented as a single dash that signifies a

bond, while the lone pairs are still drawn as dots.

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Experiment Molecular Modeling and Lewis Structures

Creating Lewis Structures

Lewis structures are created in just a few simple steps. Follow along with the procedures listed below and practice creating the Lewis structure for carbon tetrachloride.

Example 1: Carbon tetrachloride (CCl4) will be used:

Step 1) Calculate the total number of valence electrons in the molecule.

Carbon is located in group 4 of the periodic table; thus, carbon has 4 valence electrons. Chlorine, located in group 7, has 7 valence electrons and there are 4 chlorine atoms total. The total number of valence electrons in the molecule is 32.

Step 2) Arrange atoms and create single bonds.

The first atom listed in the molecular formula is often the central atom in the Lewis structure. (A more precise description is that the least electronegative atom is usually the central atom.) A wrong choice in the central atom will usually result in the inability to create the Lewis structure.

A single bond is composed of two valence electrons and is noted as either (● ●) or (—). In this example, 8 electrons are used to create one single bond between each of the C-Cl atoms.

The C atom is placed in the center and surrounded by the Cl atoms. Single lines are drawn between the atoms representing shared pairs of electrons, which may be thought of as single bonds.

When actors lose an academy award to another actor they often say, “It was an honor just to be nominated.”

While not an actor, Gilbert Newton Lewis, an American chemist for whom the Lewis structure

model was named, can certainly relate to the honor of a nomination. While Dr. Lewis was nominated for a Nobel Prize 35 times, he never won. Although Lewis was denied winning the Nobel Prize multiple times, this did not reduce the impact he made on science;

rather he has the distinct honor of having been nominated for the award more than any

other scientist (thus far!).

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Experiment Molecular Modeling and Lewis Structures

Step 3) Calculate the number of remaining valence electrons; then distribute the electrons with the goal of satisfying the duet rule or octet rule.

Since each single bond (represented by a line) contains 2 shared electrons, 8 electrons have already been added to the Lewis structure. As determined in step 1, the molecule contains 32 valence electrons total. Thus, there are 24 more electrons (32-8=24) to add to the Lewis structure.

C already satisfies the octet rule.

Placing the remaining 24 valence electrons around the chlorine atoms satisfies the octet rule for each Cl. The lone pairs are represented by pairs of dots.

Step 4) Check your work: ensure the duet and octet rules are satisfied and count the total number of valence electrons. Review the molecule to ensure that all atoms are surrounded by 8 valence electrons, satisfying the octet rule. Ensure that hydrogen atoms (when present) satisfy the duet rule. Count the number of electrons represented in your diagram. Ensure that the number of elec- trons you count matches the number calculated in step 1.

Example 2: Oxygen gas (O2)

Step 1) Calculate the total number of valence electrons in the molecule.

Oxygen atoms have 6 valence electrons, and there are a total of 2 atoms. There are 12 valence electrons total.

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Experiment Molecular Modeling and Lewis Structures

Step 2) Arrange atoms and create single bonds.

The two oxygen atoms should be placed in a line and linked with a single bond that represents shared electrons.

Step 3) Calculate the number of remaining valence electrons; then distribute the electrons with the goal of satisfying the duet rule or octet rule.

Two electrons (held in 1 single bond) have already been added to the Lewis structure. This molecule has 12 valence electrons total, as determined in step 1. Ten more electrons must be added, and this can be accomplished through a trial-and-error process. Creating a double bond between the atoms and then adding lone pairs satisfies the octet rule for all atoms and also generates the correct number of valence electrons. (Double bonds represent 4 shared electrons.)

Consider the outcome if the Lewis structure was drawn with a single bond, and the octet rule was fulfilled. The total number of electrons would be 14. In step 1, we determined that the molecule has 12 valence electrons, so we know that the Lewis structure below is incorrect, even though the octet rule is fulfilled.

Consider the outcome if the Lewis structure was drawn with the correct number of total valence electrons and a single bond. The Lewis structure below is incorrect because, although 12 electrons are present, the octet rule is not fulfilled.

Note: Multiple bonds are used ONLY when there are not enough lone pairs present for each atom to fulfill the octet rule. In the incorrect Lewis structure above, we know a multiple bond is needed because the correct total valence electrons are present but the octet rule is not fulfilled.

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Experiment Molecular Modeling and Lewis Structures

Step 4) Check your work: ensure the duet and octet rules are satisfied and count the total number of valence electrons.

A review of the correct molecule shows that all atoms are surrounded by 8 valence electrons, satisfying the octet rule. The number of electrons drawn in the Lewis structure (in the lone pairs and the double bond) totals 12, which matches the calculation in step 1.

Example 3: Cyanide (CN-)

Step 1) Calculate the total number of valence electrons in the molecule.

As indicated in the molecular formula, this molecule has a negative charge. Therefore, an additional electron must be accounted for. Carbon contributes 4 valence electrons, nitrogen contributes 5 valence electrons, and the negative charge in the chemical formula indicates 1 additional valence electron.

There are 10 valence electrons total in CN-.

Note: In this example, a “–“charge was shown. If a “+” charge were present, then valence electron(s) would be subtracted rather than added.

Step 2) Arrange the atoms and create single bonds.

The two atoms are placed in a line and linked with a single bond.

Step 3) Calculate the number of remaining valence electrons; then distribute the electrons with the goal of satisfying the duet rule or octet rule.

Two electrons were added to the Lewis structure through the addition of the single bond; 8 more electrons were required. Through trial-and-error, it can be determined that the two atoms share 6 electrons in a triple bond. The C and N atoms each have one set of lone pairs, satisfying the octet rule.

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Experiment Molecular Modeling and Lewis Structures

Consider the outcome if the Lewis structure was drawn with a single bond, and the octet rule was fulfilled. The total number of electrons would be 14. In step 1, we determined that the molecule has 10 valence electrons, so we know that the Lewis structure below is incorrect, even though the octet rule is fulfilled.

Consider the outcome if the Lewis structure was drawn with the correct number of total valence electrons and a single bond. The Lewis structures below are both incorrect because, although 10 electrons are present in both structures, the octet rule is not fulfilled.

Consider the outcome if the Lewis structure was drawn with a double bond and the octet rule was fulfilled. The Lewis structure below is incorrect because, yet again, the number of valence electrons is incorrect. The total number of electrons in the structure below is 12, and it was determined that the molecule has 10 valence electrons.

The correct Lewis structure has a triple bond; carbon has one lone pair, and nitrogen has one lone pair. The octet rule is fulfilled, and the valence electrons total 10, as calculated in step 1.

However, the Lewis structure is not complete. The charge of the molecule is denoted by placing the structure in brackets and writing the charge in the upper right-hand corner.

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Experiment Molecular Modeling and Lewis Structures

Resonance Structures

Resonance occurs when more than one valid Lewis structure exists. In this scenario, the various molecular structures may be referred to resonance structures.

Consider the Lewis structure for the polyatomic ion CO3 -2. Following the steps for drawing Lewis

structures, we find that CO3 -2 has a total of 24 electrons (4+(3)6+2=24). Through trial-and-error, it

is determined that one double bond, two single bonds, and 8 lone pairs exist.

A Lewis structure for CO3 -2 can be drawn as follows:

In the diagram above, the oxygen atoms are evenly distributed around the central carbon atom. However, notice that there is more than one option for the placement of the double bond. The double bond could be placed between the C atom and any one of the O atoms. Thus, the Lewis structure may be represented in three drawings.

The CO3 -2 molecule has resonance, which occurs when more than one valid Lewis structure

exists. In reality, the electron structure of CO3 -2 is a combination of all three resonance structures.

Resonance is represented by double-headed arrows as follows. Notice that the -2 charge is denoted for each resonance structure.

Exceptions to the Octet Rule

Oxygen, fluorine, nitrogen, and carbon always obey the octet rule. However, it is often said in jest that the only rule in science is, “There is an exception to every rule.” Indeed, the octet rule applies to most atoms in molecules, but there are exceptions. Some elements, such as Boron, tend to have fewer than 8 valence electrons. Likewise, other elements can have more than 8 valence electrons. Only the elements located in or below period 3 (row 3) of the periodic table tend to exceed an octet.

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Experiment Molecular Modeling and Lewis Structures

Consider the Lewis structure for iodine tetrachloride (ICl4 -). The sum of the valence electrons is

36 (7+(4)7+1=36). Drawing the molecule and placing lone pairs around the Cl atoms results in a representation of only 32 electrons.

Where do the remaining 4 electrons go? Since iodine is located below period 3 of the periodic table, the remaining lone pairs may be placed around the central iodine atom. (You may learn later in your course that this occurs because elements like iodine can exceed the octet by using their empty valence d orbitals.) In the case of ICl4

-, iodide breaks the octet rule, while the Cl atoms obey the octet rule.

As shown above, the total number of valence electrons (36) are represented in the Lewis structure. The placement of the Cl atoms and the lone pairs are adjusted so that the central iodide atom is evenly surrounded.

Valence shell electron pair repulsion (VSEPR)

While Lewis structure models describe atoms in the two-dimensional sense, the valence shell electron pair repulsion (VSEPR) model describes the three-dimensional arrangement of the molecule.

The VSEPR model arranges atoms in a manner that minimizes electron pair repulsion, maintaining the most stable form of the molecule. Consider the two-dimensional Lewis structure for methane (CH4) shown in Figure 6. The three-dimensional VSEPR model has bond angles of 109.5 degrees, placing the hydrogen atoms as far from one another as possible

Figure 6. Lewis structure and VSEPR model of methane (CH4)

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Experiment Molecular Modeling and Lewis Structures

The molecular geometry (or molecular shape) of CH4 is classified as a tetrahedral arrangement. In fact, anytime one atom is surrounded by four other atoms (and no lone pairs), the geometry will be tetrahedral and the molecule will exist in its most stable form.

A selection of molecular geometries is shown in Figure 7. Central atoms are represented by grey spheres, and surrounding atoms are represented by red spheres. The black bars represent bonds, which could be single or double bonds. Lone pairs are represented by green clouds.

Figure 7. Selection of molecular geometries showing central atoms (grey Spheres), surrounding atoms (red spheres), single/double bonds (black bars), and electron pairs (green clouds).

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Experiment Molecular Modeling and Lewis Structures

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Experiment Molecular Modeling and Lewis Structures

Lone pairs are not considered part of the molecular geometry, but the presence of lone pairs around the central atom affects the arrangement of the other atoms and dictates the shape of the entire molecule. For example, compare the trigonal planar and trigonal pyramidal structures in Figure 7. Both structures have one central atom (grey) and three surrounding atoms (red). Notice that the trigonal planar arrangement has no lone pairs and it is flat. The trigonal pyramidal arrangement does have a lone pair and the structure is not flat- it is “lifted” like a pyramid. In your course, you may also learn about electron-pair geometry, which does take into account the shapes of the electron configuration; however, the focus of this lab experience will be molecular geometries around the central atom.

Review all of the geometries in Figure 7. Try to predict which geometries will allow the central atom to have more than eight valence electrons.

Note: Please note that there are additional shapes that are not shown in Figure 7. For example, a central atom surrounded by one lone pair and four atoms is an “irregular tetrahedron” (not shown). Variations on the geometries listed also exist. A bent molecule can include a central atom surrounded by two atoms and only one lone pair as opposed to the two lone pairs shown in Figure 7. The figure includes all of the geometries you need to be successful in the Experimentation. For more geometries, consult a textbook or a reliable internet source.

Molecular Modeling Kits

Molecular modeling kits are used to visualize the three-dimensional structure of a molecule. Kits typically contain balls that represent atoms, sticks that represent bonds, and paddles that represent lone pairs. See Figure 8.

Figure 8. Molecular modeling kit.

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Experiment Molecular Modeling and Lewis Structures

The Lewis structure and molecular model for chloromethane (CH3Cl) are shown in Figure 9. The Lewis structure includes a carbon atom surrounded by three hydrogen atoms and one chlorine atom. The chlorine atom has 3 lone pairs. Creating the molecule with a modeling kit shows the three-dimensional placement of the atoms and lone pairs. The model includes one carbon atom (black) surrounded by three hydrogen atoms (small, white), and one chlorine atom (dark green) surrounded by three lone pairs (pink paddles).

Figure 9. Methyl chloride (CH3Cl): Lewis structure and molecular model.

To identify the molecular geometry, the central atom must be identified. Then the geometry can be determined using Figure 7 as a guide. For example, the central atom in CH3Cl is carbon. Using Figure 7 as a guide, it may be determined that the geometry around the carbon atom is tetrahedral. Remember, the lone pairs around the chlorine atom are not included when determining the molecular geometry.

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Experiment Molecular Modeling and Lewis Structures

Geometry around Multiple Central Atoms

When a molecule has more than one interior atom, more than one geometry may be used to describe the molecule. Consider the Lewis structure and molecular model of ethylene (C2H4) in Figure 10. Note that the two carbons are connected by a double bond. Each carbon is considered a central atom.

Figure 10. Ethylene C2H4: Lewis structure and molecular model.

Focus on only one interior atom at a time. Using Figure 7 as a guide, the geometry around only the leftmost carbon atom of C2H4 is trigonal planar. Taking only the rightmost carbon into account, the geometry is again trigonal planar. For C2H4 it may be concluded that the geometry is “trigonal planar around each carbon atom.”

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Experiment Molecular Modeling and Lewis Structures

Exercise 1: Lewis Structures and Molecular Modeling In this exercise, the student will draw Lewis structures for a series of molecules and then create the VSEPR model for the molecule using the modeling kit.

Procedure

1. Gather the modeling kit, a white sheet of paper, a pen, and the digital camera.

2. Review the key in Table 1. Check that the contents of your modeling kit match the key.

Table 1. Element, color, and holes.

Element/bond/etc. Color Holes # Included Hydrogen White 1 5 Magnesium Pink 2 1 Aluminum Gray 3 1 Carbon Black 4 2 Oxygen Red 4 2 Chlorine and Florine Green 4 6 Lead and Nitrogen Blue 4 1 Sulfur and Iodine Yellow 4 4 Tin White 5 1 Sulfur and Iodine Yellow 6 1 Single Bond Gray Short and rigid 6 Double Bond Gray Long and flexible 4 Lone Pair (2 valence electrons) Pink Paddle shaped 18

Part 1: Practice Describing Molecular Structures

3. Practice describing the molecular structure of CHO2 - (See Lab Report Assistant for Answers).

Number of Valence Electrons:

● How many valence electrons does CHO2 - have? Use the periodic table to calculate the

total.

Lewis Structures:

● What is the Lewis structure for CHO2 - ? Practice drawing the Lewis structure on a separate

sheet of paper.

Hint: CHO2 - has resonance structures, and there are two forms of the drawn molecule. Draw both

structures.

VSEPR Models:

● Use your molecular modeling kit to create a CHO2 - molecule. Although the molecule has

two Lewis structures, you only need to build one molecule.

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Experiment Molecular Modeling and Lewis Structures

Note: Consult Table 1 to determine which pieces represent the C, H, and O atoms. To create a double bond, use TWO of the long, flexible gray connectors. To create a single bond, use one of the short, inflexible connectors. Pink paddles represent lone pairs. The completed molecule should have no “open” or unfilled holes.

Atoms:

● What is the central atom? If there is more than one interior atom, list each.

● How many bonds and electron pairs surround the central atom(s)?

Geometry:

● Identify the molecular geometry of the molecule. Refer to Figure 11 as needed.

Figure 11. Molecular geometries.

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Experiment Molecular Modeling and Lewis Structures

4. Practice the questions above until you feel comfortable describing the CHO2 - molecule.

Note: If you would like additional practice, describe the molecules in the background section.

Part 2: Describing Molecular Structures

5. Using Table 1 as a guide, label and organize the atoms of the molecular modeling kit. See Figure 12.

Figure 12. Organized molecular modeling kit.

Note: You might not use all of the pieces included in the modeling kit during this exercise.

6. For each molecule listed below, you will: 1- calculate the number of valence electrons, 2- draw the Lewis structure(s), 3- upload an image of the VSEPR model, 4- list the number of bonds and lone pairs surrounding the central atom(s), and 5- identify the structure geometry; or identify multiple geometries if there is more than one central atom.

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Experiment Molecular Modeling and Lewis Structures

7. Use the periodic table in Figure 2 of the Background to calculate the number of valence electrons in CCl4. Click here to download a printable copy of the periodic table. Record in Data Table 1 of your Lab Report Assistant.

8. Draw the Lewis structure for CCl4 and insert into Data Table 1. Ensure that the features of the molecule are large enough and clear enough for your instructor to grade. Refer to the background for step-be-step instructions on drawing Lewis structures.

Note: To draw the Lewis structure, either use a computer drawing program, such as Power Point, or other computerized drawing tool and cut and paste into Data Table 1, or draw by hand, scan, resize, and then insert into Data Table 1, or draw by hand and take a photograph, resize and then insert into Data Table 1. Refer to the appendix entitled, “How to Label an Image” for guidance Refer to the appendix entitled, “Resizing an Image” for guidance. If resonance structures exist for any of the molecules, make sure to draw all of the structures.

9. Use the molecular modeling kit to build CCl4. Ensure that the appropriate atoms are used, as described in Table 1. The completed molecule should have no “open” or unfilled holes.

10. Place the molecular model of CCl4 on a white sheet of paper with your name and the date written on it, take a photograph of the model, as shown in Figure 13. Resize and insert the image into Data Table 1.

Figure 13. NCH3CH3CH3: molecular model with student name and date.

11. Determine the number of bonds and lone pairs that surround the central atom. Record the data in Data Table 2 of your Lab Report Assistant.

Note: If there is more than one central atom, identify each atom and list the number of electron clouds that surround each.

http://holscience.com/sites/default/files/Periodic_Table.pdf
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Experiment Molecular Modeling and Lewis Structures

12. Use Figure 11 to determine the name of the geometric structure for ClC4. Structure names include: linear, bent, trigonal planar, trigonal pyramidal, tetrahedral, square pyramidal, and octahedral. Record the name Data Table 2.

13. Repeat steps 5-12 for the remaining molecules.

Hints:

● Some molecules have more than one central atom: multiple geometries will need to be described.

● Some molecules have a charge: draw brackets and indicate the charge in the Lewis structure.

● Some molecules have resonance: draw all of the forms for the structure.

● Most, but not all, of the molecules fulfill the duet and octet rules.

14. When you are finished uploading photos and data into your Lab Report Assistant, save and zip your file to send to your instructor. Refer to the appendix entitled “Saving Correctly,” and the appendix entitled “Zipping Files,” for guidance with saving the Lab Report Assistant in the correct format.

Cleanup:

15. Return all items to the kit for future use.

Questions A. Which, if any, of the molecules in Exercise 1 had resonance structures? How many resonance

structures did each molecule have?

B. Both chlorine and fluorine are represented by a green modeling piece that has 4 holes. Is using the same piece for two different atoms acceptable? Why or why not?

C. List all of the possible geometric structures of a molecule that contains two atoms.

D. Explain why only the lone pairs on the central atom are taken into consideration when predicting molecular shape. What molecules from the lab activity are exceptions to this?