Molecular Geometry. Objectives N H H. The objectives of this laboratory are to:

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1 Objectives The objectives of this laboratory are to: Molecular Geometry Write Lewis structure representations of the bonding and valence electrons in molecules. Use the VSEPR model to predict the molecular geometries (shapes) of molecules. Determine whether a molecule is polar (has a dipole moment) or not. Background The physical and chemical properties of a molecule are a direct consequence of the way in which the atoms are bonded together and the three-dimensional geometry, or shape, of the molecule. For example, water is a polar molecule, due in part to it s bent geometry: The polarity of a molecule (or lack thereof) influences physical properties such O solubility, boiling point, freezing point, vapor pressure, and chemical reactivity, H to name a few. An understanding of molecular geometry is central in H determining the chemical and physical properties of molecules. The first step in determining molecular geometry is to write the Lewis structure for a molecule, which gives the arrangement of valence electrons in terms of bonding pairs and lone pairs of electrons. The Valence Shell Electron Pair Repulsion (VSEPR) model is then applied to determine a molecule s threedimensional geometry (or shape). The basic premise of the VSEPR approach is that the electron pairs surrounding an atom in a molecule repel each other, and tend to minimize the repulsions by positioning themselves as far away from each other as possible. The application of VSEPR permits the prediction of molecular geometries and bond angles. This information can be used to subsequently determine other properties, such as molecular polarity. For example, consider ammonia, NH 3, which has the Lewis structure: H N H The nitrogen atom has four pairs of valence electrons, 3 bonding pairs H (shared with H atoms) and 1 lone pair (unshared). The VSEPR model predicts that the four pairs minimize repulsions by adopting a tetrahedral arrangement around the nitrogen. The arrangement of all electron pairs bonding pairs and lone pairs around the central atom is referred to as the electron pair arrangement. By contrast, the molecular geometry does not include the lone pairs, but only describes the geometric arrangement of the atoms in the molecule. Thus, in the example of NH 3, the electron pair arrangement is tetrahedral while the molecular geometry is trigonal pyramidal (shown at the right). VSEPR allows the prediction of ideal bond angles; for the tetrahedral electron pair arrangement of the NH 3 molecule, the ideal H-N-H angle would be H predicted to be 109.5, the tetrahedral angle. Table 1 lists the molecular geometries predicted via the application of the VSEPR model. N H H Before coming to lab please review these sections in the textbook (Zumdahl & ) with particular focus on Lewis structures, molecular geometry & VSEPR, and molecular polarity.

2 Procedure for Drawing Lewis Structures: 1. Add up the total # of valence electrons for all the atoms. Account for charge: If the species has a negative ( ) charge: add one valence electron for each negative charge; for a positively charged (+) species, subtract one electron for each positive charge. 2. Draw the molecular skeleton and connect the atoms with one bond. The central atom is generally the atom with the lowest electronegativity, but never H. 3. Satisfy the Octet rule: Distribute the remaining valence electrons by adding lone pairs to complete the octets of the outer atoms first (H only requires 2 electrons), then place any remaining e s on the central atom. 4. If there are too few valence electrons to give each atom an octet, multiple bond(s) may be required. In this case, convert outer atom lone pairs to bond pairs to form multiple bonds. 5. There are some exceptions to the octet rule : 3rd period or heavier elements may more than 8 electrons around them if needed (expanded valence shells). Atoms such as Be, B and Al may have less than an octet of electrons. In species with an odd number of electrons the least electronegative atom carries the odd electron. VSEPR Procedure: 1. Draw the Lewis structure 2. Determine the number of electron pair domains on the central atom. This is is equal to the number of atoms plus the number of lone pairs attached to the central atom. (Electron domains are lone pairs or bonding pairs of electrons, but multiple bonds count as one domain). 3. Determine the geometrical arrangement of electron domains (also see Table 1): 2 domains 3 domains 4 domains 5 domains 6 domains Linear Trigonal Planar Tetrahedral Trigonal Bipyramidal Octahedral 4. Re-draw the molecule in the correct electron domain geometry place the lone pairs in their proper positions in order to minimize repulsions (refer to Table 1). 5. Determine the molecular geometry (shape) of the molecule (the geometric arrangement of just the atoms ignore the lone pairs )

3 Table 1. Molecular Geometries (Shapes) Predicted by VSEPR Total Electron Domains No Lone Pairs One Lone Pair Two Lone Pairs Three Lone Pairs 2 Geometry: Linear: Angles: 180 : 3 Geometry: Trigonal Planar Bent Angles: 120 : ~120 : 4 Geometry: Angles: Tetrahedral Trigonal-pyramid ~109.5 bent ~ Geometry: Angles: Trigonal Bipyramidal 90, 120 See-saw 90, 120 T-shaped 90 Linear Geometry: Angles: Octahedral 90 Square-pyramid 90 Square-planar 90

4 Prediction of Polarities of Molecules from VSEPR and Electronegativities A polar covalent bond is formed when a pair of electrons is shared between two atoms with different electronegativities. The atom with the greater electronegativity exerts a greater attraction for the electrons in the bond than does the atom with the smaller electronegativity. The electron density is displaced toward the more electronegative atom, with the net effect that the more electronegative atom has a partial negative character and the less electronegative atom has a partial positive character. The HCl molecule is an example of a molecule possessing a polar covalent bond: δ+ δ H Cl The Cl atom has a greater share of the bonding electrons than H atom due to the larger electronegativity of the Cl atom. The HCl molecule is said to possess a dipole moment due to the separation of the partial positive and negative charges. The dipole moment (or polarity of the bond) is indicated by an arrow pointing towards the negative end of the bond ( and molecule in this case since there is only one bond). In molecules containing two or more bonds, the overall polarity of the molecule (or lack thereof) depends on the net sum of the individual bond polarities. In some cases the individual bond polarities may cancel, giving rise to no net dipole moment and an overall non-polar molecule. In the case of H 2 O, the individual O-H bond polarities do not cancel due to the bent shape of H 2 O, and the result is a polar molecule. O O H H H H bond dipoles molecular diople In the carbon dioxide (CO 2 ) molecule, the linear geometry of the molecule results in the individual bond polarities exactly canceling one another and the result is that CO 2 is a non-polar molecule: O C O In general, a molecule will be non-polar when there is a symmetrical arrangement of the same type of bonds, such that the individual bond polarities cancel one another. The five basic electron domain geometries from VSEPR (linear, trigonal planar, tetrahedral, trigonal bipyramidal and octahedral) plus the square planar geometry are symmetrical arrangements. Polar molecules result when there is an asymmetric arrangement of bonds, such that individual bond polarities do not exactly cancel one another.

5 Experimental Procedures During this lab period you will determine the Lewis Structures molecular geometries, and predicted polarities for a number of molecular species beginning only with the molecular formula. As well, you will need to create a 3-dimensional model of the molecule with ~ correct bond angles using a set of toothpicks and Styrofoam balls. You may not be as familiar with the geometries that include 5 or 6 electron domains, so be sure to refer to the table of geometries at the end of the lab. You will be assigned a set of molecular species (see list below). You are to do the following tasks for each species, and write the answers in the space provided on the worksheets: 1. Determine the number of valence electrons in the molecule (or polyatomic ion). 2. Determine and draw the Lewis structure for the molecule (or polyatomic ion). 3. Determine the molecular geometry (shape) using VSEPR 4. Construct a model of the molecule using the materials provided 5. Draw a correct 3 D representation of the molecule from your results and the model 6. Determine whether the molecule (or polyatomic ion) is polar or non-polar. AXE Method The "AXE method" of electron counting is commonly used when applying the VSEPR theory. The A represents the central atom and always has an implied subscript one. The X represents the number of sigma bonds between the central atoms and outside atoms. Multiple covalent bonds (double, triple, etc.) count as one X. The E represents the number of lone electron pairs surrounding the central atom. The sum of X and E is known as the steric number. The number of lone pair electrons can often be determined by multiplying the number of sigma bonds (X) times 8 and subtracting this from the total number of electrons. The number of lone pairs is this difference divided by 2. In other words: E = (# electrons 8X)/2 Molecular Species List 1. SiCl 4 2. PCl 3 3. NO 3-4. SF 6 5. PBr 5 6. IF 3 7. XeF CO 2 9. BrF SO H 2 O 12. SF XeF 4

6 Molecular Geometry Worksheet Name Section Lab Partner Molecule List # 1. Formula: # of valence electrons = 2. Formula: # of valence electrons = 3 D Drawing (include bond angles) Polar or Non-Polar?

7 3. Formula: # of valence electrons = 3 D Drawing (include bond angles) Polar or Non-Polar? 4. Formula: # of valence electrons = 3 D Drawing (include bond angles) Polar or Non-Polar?

8 5. Formula: # of valence electrons = 6. Formula: # of valence electrons =

9 7. Formula: # of valence electrons = 8. Formula: # of valence electrons =

10 9. Formula: # of valence electrons = 10. Formula: # of valence electrons =

11 11. Formula: # of valence electrons = 12. Formula: # of valence electrons =

12 13. Formula: # of valence electrons =