Chemistry Lab: Bonding Molecular Geometry

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Please complete these tables for my chemistry lab

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Data Table 1

Activity 1

Lewis Dot Structures

Group 1 Molecules

Group 2 Molecules

Group 3 Molecules

Chemical Formula

BeCl2

BF3

CH4

CO2

NH3

H2O

O2

HCN

H2CO

Dot Diagrams

Electrons

Central Atom

Lewis Structure

Activity 3

Molecular Geometry

Group 1 Molecules

Group 2 Molecules

Group 3 Molecules

Domains

Lone Pairs

Geometry

Bond Angle

Line Drawing

Activity 4

Bonding and Polarity

Group 1 Molecules

Group 2 Molecules

Group 3 Molecules

Bond Electronegativity

Bond Polarity / Type

Ionic, Polar Covalent, Non polar Covalent

Bond Dipole Moment

Molecule Polarity

Data Table 2

Bonding Pair

Electronegativity Difference

Type of Bond

Bond Polarity (position of dipole arrow)

C and H

N and H

H and O

Xe and F

H and Cl

O and O


Data Table 3

Straight-Chained Hydrocarbons

Insert drawings, not images from the internet.

Class

Type of Bonding

General Formula

Alkanes

Alkenes

Alkynes


Example

Structural Formula

Shape around the Carbon

Ethane

C2H6

Ethene

C2H4

Ethyne

C2H2

Cyclic and Acyclic Compounds

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Cyclohexane

Hexane



Skeletal Model

Unformatted Attachment Preview

CHEMISTRY Introduction to Molecules: A Molecular Bonding and Shapes Investigation Investigation Manual INTRODUCTION TO MOLECULES Table of Contents 2 Overview 2 Objectives 2 Time Requirements 3 Background 5 Materials 5 Safety 5 Preparation 6 Activity 1 8 Activity 2 9 Activity 3 11 Activity 4 14 Activity 5 21 Disposal and Cleanup 22 Data Tables Overview Learn about bonding and molecular geometry using molecular models. This investigation includes classifying bonds, drawing Lewis structures, predicting molecular geometry, constructing three-dimensional models, and determining polarity. Resonance structures and isomers will also be identified as needed. The activities will require approximately 2 hours to complete. Objectives • Use electronegativity difference to classify bonds as ionic, polar covalent, or nonpolar covalent. • Apply valence bond theory to draw appropriate Lewis structures. • Apply valence shell electron pair repulsion theory (VSEPR) to predict molecular geometry. • Construct three-dimensional molecular models. • Sketch three-dimensional line structures using dashed lines and wedges. • Determine whether the polarity of a molecule results in a dipole moment. • Identify resonance structures and isomers as needed. Time Requirements Preparation ..................................................................... 10 minutes Activity 1: Lewis Dot Structures ..................................... 30 minutes Activity 2: Electron Domain Geometries......................... 30 minutes Activity 3: Molecular Geometry ...................................... 30 minutes Activity 4: Bonding and Polarity .................................... 20 minutes Activity 5: Modeling Hydrocarbons ................................ 40 minutes Key Personal protective equipment (PPE) goggles gloves apron Made ADA compliant by NetCentric Technologies using the CommonLook® software 2 Carolina Distance Learning follow link to video photograph stopwatch results and required submit warning corrosion flammable toxic environment health hazard Background Few elements exist in nature as single atoms. Other than the noble gases, the atoms in most elements are bonded to other atoms. This is because atoms are more stable if their outer energy level is full. Electrons in the outer energy level are called valance electrons. The halogens (Group 17 in Figure 1) have seven valance electrons in their outer energy level. Figure 1. Each element in group 18 has a full outer energy level. For most elements, a full outer energy level has eight electrons, or an octet. Helium and hydrogen are exceptions in that they only require two valence electrons. continued on next page www.carolina.com/distancelearning 3 INTRODUCTION TO MOLECULES Background continued Atoms can fill their outer energy level by transferring or sharing electrons to form either ionic or covalent compounds (Figure 2). Figure 2. Ionic bond (electron is transferred to Y, giving it an octet) Covalent bond (two electrons are shared, giving both atoms a filled outer shell) The tendency for atoms to bond ionically or covalently is determined by the difference in their electronegativity. Electronegativity is a dimensionless number that is a measure of an atom’s ability to attract bonding valence electrons. Electronegativity values show periodic trends on the periodic table, as shown in Figure 3. Figure 3. 4 Carolina Distance Learning Excluding the noble gases, the most electronegative element is fluorine, which is assigned a value of 4.0. All the other elements have calculated values based on that of fluorine. Across each period the electronegativity tends to increase, with the nonmetal families of nitrogen, oxygen, and fluorine having the highest values. This is because these atoms have smaller radii within which the positive nuclei exert a greater attraction for the valence electrons. The alkali metals and alkaline earth metals (the groups on the left side of the periodic table) have the lowest electronegativity because their atoms have the largest radii, within which the positive nuclei exert a smaller attraction for valence electrons. Cesium and francium, with the largest radii, have the lowest electronegativity value of 0.7. Materials Included in the materials kit: 25 balloons Needed from the modeling kit: Molecular model set Needed from the equipment kit: Safety Read all instructions for this laboratory activity before beginning. Follow the instructions closely and observe established laboratory safety practices, including use of appropriate personal protective equipment (PPE) as described in the Safety and Procedure section. Model pieces can be a choking hazard. Keep these and all laboratory materials away from children. Latex balloons are included in this kit and can cause an allergic reaction in some individuals. Individuals sensitive to latex should not participate in any activities that may result in exposure. Preparation 1. Read the Student Guide. 2. Obtain all materials. Ruler Needed, but not supplied: • Calculator • Digital camera or smart phone • Pencil 3. Separate model pieces into tray compartments based on shape, color, and number of prongs. Activities 1, 3, and 4 are interrelated. The compounds used in Activity 1 are also used in Activities 3 and 4. Data in Data Table 1 will be used for these activities. Reorder Information: A replacement kit for Introduction to Molecules, item number 580306, can be ordered from Carolina Biological Supply Company. Call 800-334-5551 to order. www.carolina.com/distancelearning 5 ACTIVITY ACTIVITY 1 A Lewis Dot Structures A Lewis dot structure is one way to represent the arrangement of valence electrons in a molecule. In Lewis structures, each element symbol is surrounded by a specific number of dots representing valence electrons. Most atoms obey the octet rule*, which states that atoms tend to combine in such a way that they each have eight electrons in their valence shells, giving them the same electronic configuration as a noble gas. To fill their outer shells, elements form covalent bonds by sharing electrons. Lewis structures are often used to show those bonds. Covalent compounds may have single, double, or triple bonds between atoms. These bonds are represented in a Lewis structure by dashes between the chemical symbols of the bonded elements. The number of dashes corresponds to the number of bonds. Lewis structures also indicate the lone pairs of electrons on different atoms. Guidelines for drawing a Lewis structure for a given chemical formula are in the table below. Molecules to be used in Activities 1, 3, and 4 are found on page 22. *Exceptions: Hydrogen and helium need only two valence electrons to fill their outer shell. Boron and beryllium form compounds with less than eight electrons and elements in periods 3–6 may use more than eight. Use the Guidelines for Drawing Lewis Structures to complete the Data Table for Activity 1. Guidelines for Drawing Lewis Structures Example SiO2 1. Determine the number of valence electrons present for each atom in the compound (consult Figure 1 in the background section as needed). Silicon (Si): 4 valence electrons Oxygen (O): 6 valence electrons Oxygen (O): 6 valence electrons 2. Find the sum of valance electrons of all the atoms. If there is a charge on a compound, add an electron for each negative charge and subtract an electron for each positive charge. Silicon (Si): 4 valence electrons Oxygen (O): 6 valence electrons Oxygen (O): 6 valence electrons 16 total electrons 3. Determine which atom is to be the central atom. This is typically the least electronegative element (that is not hydrogen) and/or the first element in the chemical formula. Some compounds will have no central atom, but the first element in the formula is a good place to start. 4. Write the chemical symbol of the central atom, followed by the chemical symbols for other atoms around it. continued on next page 6 Carolina Distance Learning Guidelines for Drawing Lewis Structures Example SiO2 5. Use dots to show the appropriate number of valence electrons around each symbol in the compound, based on their position in the periodic table. 6. Determine the total number of electrons required for each element. Allow for exceptions to the octet rule as needed. Silicon (Si) needs 8 valence electrons Oxygen (O) needs 8 valence electrons Oxygen (O) needs 8 valence electrons 7. Create a skeleton structure by connecting unpaired electrons from the central atom to unpaired electrons on the atoms around it with a solid line. Each solid line now represents a (two-electron) covalent bond between the two atoms. 8. Redraw the structure as needed. If multiple bonds exist between two atoms, reorient the chemical symbols so that multiple bonds are with each other. 9. Sometimes there will be more than one way to add the bonds. To determine which Lewis structure is the best model, formal charge is often used. To determine the formal charge of atoms: a. Count the number of valence electrons that each atom typically has when unbound. b. In the Lewis structure being evaluated, count the number of assigned electrons for an atom by counting the number of electrons that exist around that atom as lone pairs and then counting the number of covalent bonds on that atom. Add the number of lone pair electrons and the number of bonds together. c. Subtract the number of assigned electrons from the number of valance electrons. The resulting value is the formal charge for the atom in that Lewis structure. 10. Lewis structures should show negative charges on the most electronegative elements and positive charges on the least electronegative elements. unbound electrons 6 4 6 assigned electrons -6 -4 -6 formal charge 0 0 0 unbound electrons 6 4 6 assigned electrons -7 -4 -5 formal charge -1 0 +1 is the preferred structure of Si02 www.carolina.com/distancelearning 7 ACTIVITY ACTIVITY 2 A Electron Domain Geometries The three-dimensional shape of a molecule is predicted using the valence-shell electron-pair repulsion (VSEPR) model. The VSEPR model is used to predict structures of molecules or ions that contain only nonmetals by minimizing the electrostatic repulsion between the regions of high electron density. In this activity, balloons will model electron domains. An electron domain is either a bond or a lone pair of electrons. Double and triple bonds count as only one domain when using the VSEPR model. When tied together, balloons naturally adopt the lowest energy arrangement predicted by the VSEPR model. Linear Geometry – Two Electron Domains 1. Inflate a balloon and tie the open end in a knot. 2. Repeat to prepare a second balloon of approximately the same size. 3. Tie the two balloons together in a knot. 4. Label the model and photograph it for future reference. Trigonal Planar Geometry – Three Electron Domains 1. Inflate and tie three balloons. 2. Tie the three balloons together. Start by tying two balloons together and then use the free ends to knot around the third. 3. 8 Label the model and photograph for future reference. Carolina Distance Learning Tetrahedral Geometry – Four Electron Domains 1. Inflate and tie four balloons. 2. Tie the balloons together to form two sets of two balloons. 3. Twist the two sets of two balloons together by crossing the knotted centers. 4. Label the model and photograph it for future reference. Trigonal Bipyramidal Geometry – Five Electron Domains 1. Inflate and tie five balloons. 2. Tie four balloons together to form two sets of two balloons. 3. Tie a third balloon to one of the sets. 4. Twist the set of three balloons and the set of two balloons together by crossing the knotted centers. 5. Label the model and photograph it for future reference. Octahedral Geometry – Six Electron Domains 1. Inflate and tie six balloons. 2. Tie the balloons together to form three sets of two balloons. 3. Twist two sets of two balloons together by crossing the knotted centers. Twist a third set of two balloons into the other two by crossing the knotted centers. 4. Label the model and photograph it for future reference. ACTIVITY 3 d. Black, four-pronged: Carbon A Molecular Geometry f. Red, four-pronged: 1. Construct a three-dimensional model of each molecule from Activity 1 using the following molecular model pieces to represent atoms. a. White one-pronged: b. Light green one-pronged: c. Dark brown three-pronged: e. Blue, four-pronged: Nitrogen Oxygen g. Purple five-pronged (and/or bonds): Atoms with five electron pairs Hydrogen h. Silver six-pronged: All other atoms Use white tubes to connect the pronged pieces and represent covalent bonds between atoms in each molecule. Atoms with only three electron pairs and/or bonds Atoms with six electron pairs and/or bonds As needed, connect the white, flat pieces for electron pairs on the central atom. continued on next page www.carolina.com/distancelearning 9 ACTIVITY ACTIVITY 3 continued 2. Use the white tubes to form bonds between the atoms. Each tube represents one bond (two electrons). For a double bond, use two tubes and two prongs on each atom. 3. 10 Label each model and photograph it for future reference. Carolina Distance Learning 4. Sketch a perspective drawing of each compound. If necessary, use wedges and dashes to show the three-dimensional appearance using the table below. 5. Build models for all the compounds in Activity 1. ACTIVITY 4 A Bonding and Polarity Chemical bonds are shared electrons between two atoms. If the atoms are different, the attraction for these shared electrons will also be different. Chemical bonds are classified as covalent or ionic. Electrons are equally shared in nonpolar covalent bonds, unequally shared in polar covalent bonds, and fully transferred to the more electronegative atom in ionic bonds. To determine the type of bond between two atoms, first calculate the difference between their electronegativity values. Only the absolute difference is important. Nonpolar Covalent: If the difference in electronegativity is < 0.3, the electrons are equally shared between the two atoms and the bond is a nonpolar covalent bond. Some examples of molecules with nonpolar covalent bonds are Cl2, H2, CH4, and CS2. Polar Covalent: When the difference in electronegativity is between 0.4 and 1.7, there is an unequal sharing of the electrons between the two atoms. Although all electrons are constantly moving within the bonding orbital, in polar covalent bonds the electrons are more attracted to the atom with higher electronegativity and will spend more time closer to that atom. This unequal sharing creates poles of charges, thus these bonds are termed polar covalent bonds. Some examples of polar covalent molecules are NH3 and H2O. Ionic: This type of bond occurs when there is complete transfer of the electrons between the two atoms. This occurs when the difference in electronegativity is >1.8. NaCl and MgCl2 are typical examples. Bond Polarity The term polar refers to the poles of partial charge. In a polar covalent bond, electrons are more attracted to the atom with greater Figure 4. Sample Bond Type Calculations Bonding Pair Electronegativity Difference Type of Bond Li and F | 1.0 – 4.0 | = 3.0 Ionic F and F | 4.0 – 4.0 | = 0 Nonpolar covalent C and F | 2.5 – 4.0 | = 1.5 Polar covalent continued on next page www.carolina.com/distancelearning 11 ACTIVITY ACTIVITY 4 continued electronegativity. This difference results in a partial negative charge on the more electronegative atom, and a partial positive charge on the less electronegative atom. Consider the compound hydrochloric acid, HCl. The electronegativity difference is 0.9 (|3.0 – 2.1| = 0.9). Chlorine is more electronegative than hydrogen and therefore pulls the bonding pair closer, giving the hydrogen a partial positive charge (delta plus, δ+) and the chlorine a partial negative charge (delta minus, δ–), as shown in Figure 5. Figure 5. Unequal sharing of electrons can also be represented by a vector indicating bond polarity. To show bond polarity, draw a vector from the positively charged atom and the negatively charged atom. Create a plus sign using a vertical line near the less electronegative atom. The vector arrow points toward the more electronegative atom, as in Figure 6. Figure 6. Molecular Polarity Molecular polarity is similar to bond polarity. A polar molecule has regions of partial positive and negative charge and will orient itself in an electric field. This orientation in an electric field, called a dipole moment, is caused when charges are distributed asymmetrically within the molecule. The molecular dipole moment is shown 12 Carolina Distance Learning as a vector to one side of the molecule. Once again, the vector arrow points toward the more electronegative atom, as in Figure 7. Figure 7. Molecules that demonstrate a dipole moment are characterized as polar. Molecules of hydrochloric acid are polar because the bond polarity is unbalanced. Not all molecules containing polar covalent bonds are polar. If bond polarity is balanced, the partial positive and partial negative charges are distributed symmetrically. If the positive and negative charges are distributed symmetrically, the molecule does not orient itself in an electric field and does not demonstrate a dipole moment. Molecules that do not demonstrate a dipole moment are characterized as nonpolar. Magnesium chloride (Cl – Mg – Cl) is an example of a nonpolar molecule that contains polar covalent bonds. Each Mg – Cl bond is polar covalent with an electronegativity difference of 0.8 (|1.2 – 3.0| = 0.8). However, the linear molecular geometry of magnesium chloride results in symmetrical bonds and a nonpolar molecule. Magnesium chloride does not have a dipole moment because the opposing bond polarities cancel each other out (see Figure 8). Figure 8. It is also possible to arrange polar covalent bonds symmetrically in three-dimensional continued on next page space, as in carbon tetrachloride. Once again, each C – Cl bond is a polar covalent bond; the electronegativity difference is 0.5 (|2.5 – 3.0| = 0.5). The molecular geometry of carbon tetrachloride (tetrahedral) results in three-dimensional symmetry and a nonpolar molecule. Carbon tetrachloride does not have a dipole moment because the opposing bond polarities cancel each other out (see Figure 9). Polar molecules are always asymmetric; this can assist in identifying a molecule as polar. Polar molecules are the result of unbalanced bond polarity and therefore all polar molecules are asymmetrical in some way. If one chlorine atom in carbon tetrachloride (CCl4) is replaced with a hydrogen atom, the molecule becomes chloroform (CHCl3). The tetrahedral molecule is no longer symmetric; it is now both asymmetric and polar. This molecule has a region of partial positive charge near the hydrogen atom and a region of partial negative charge near the chlorine atoms. This configuration results in a molecular dipole moment shown by the large vector on the far right (see Figure 10). Figure 9. Figure 10. The presence of lone pairs of electrons can also contribute to the polarity of a molecule. When lone pairs are present, they are regions of partial negative charge. However, not all molecules containing lone pairs are polar. The combination of molecular geometry and multiple lone pairs may result in a symmetric, nonpolar molecule. 1. Complete Data Table 2 using the Periodic Table of Electronegativities and the Bonding Scale to determine the type of bond that each set of atoms would exhibit if they formed a bond. For polar covalent bonds only, draw a vector between atoms showing bond polarity in the last column. 2. Determine the electronegativity difference and bond polarity for each bond in compounds in Data Table 1. Record yo ...
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