CHM 2045C

Module Four Part II: Chemical Bonding & Molecular Structure (Chapters 7

L. ____ (05) Bond Angles/Bond Lengths-Section 7.9 Answers

N. ____ (05) Geometry of Molecules-Section 7.9 Answers

O. ____ (05) Polarity of Molecules-Section 7.4 Answers

P._____ (05) Hybrid Orbitals Sections 7.11-7.12 Answers

______(20) Total Chapter 7

M-4ii Required Homework:

______(30) M-4ii Pretest Hardcopy Homework Packet

______(35) M-4ii Multiple Choice (MC) Practice (Blackboard Online)

Reference: VSEPR Video:

http://www.lsua.info/chem1001/VSEPR/VSEPRtheory.wmv

What is VSEPR?

VSEPR stands for Valence Shell Electron Pair Repulsion.  It's a complicated acronym, but it means something that's not difficult to understand.  Basically, the idea is that covalent bonds and lone pair electrons like to stay as far apart from each other as possible under all conditions.  This is because covalent bonds consist of electrons, and electrons don't like to hang around next to each other much because they have the same charge. 

This VSEPR thing explains why molecules have their shapes.  If carbon has four atoms stuck to it (as in methane), these four atoms want to get as far away from each other as they can.  This isn't because the atoms necessarily hate each other, it's because the electrons in the bonds hate each other.  That's the idea behind VSEPR.

What is a Bond Angle?

Let’s look at the water molecule:

VSEPR table

The bond angles in the table below are ideal angles from the simple VSEPR theory, followed by the actual angle for the example given in the following column where this differs. For many cases, such as trigonal pyramidal and bent, the actual angle for the example differs from the ideal angle, but all examples differ by different amounts. For example, the angle in H₂S (92°) differs from the tetrahedral angle by much more than the angle for H₂O (104.5°) does.

Bonding electron pairs	Lone pairs	Electron domains (Steric #)	Shape	Ideal bond angle (example's bond angle)	Example	Image
2	0	2	linear	180°	CO2
3	0	3	trigonal planar	120°	BF3
2	1	3	bent	120° (119°)	SO2
4	0	4	tetrahedral	109.5°	CH4
3	1	4	trigonal pyramidal	107°	NH3
2	2	4	bent	109.5° (104.5°)	H2O
5	0	5	trigonal bipyramidal	90°, 120°, 180°	PCl5
4	1	5	seesaw	180°, 120°, 90° (173.1°, 101.6°)	SF4
3	2	5	T-shaped	90°, 180° (87.5°, < 180°)	ClF3
2	3	5	linear	180°	XeF2
6	0	6	octahedral	90°, 180°	SF6
5	1	6	square pyramidal	90° (84.8°), 180°	BrF5
4	2	6	square planar	90°, 180°	XeF4
7	0	7	pentagonal bipyramidal	90°, 72°, 180°	IF7
6	1	7	pentagonal pyramidal	72°, 90°, 144°	XeOF5^−
5	2	7	planar pentagonal	72°, 144°	XeF5^−
8	0	8	square antiprismatic		XeF8^2−
9	0	9	tricapped trigonal prismatic		ReH9^2−

Example

Predict all bond angles in the following molecules.

a. CH₃Cl   b. CH₃CNl     c. CH₃COOH

Solution

a. The Lewis structure of methyl chloride is:

In the Lewis structure of CH₃Cl carbon is surrounded by four regions of high electron density, each of which forms a single bond. Based on the VSEPR model, we predict a tetrahedral distribution of electron clouds around carbon, H - C - H and H - C - Cl bond angles of 109.5°, and a tetrahedral shape for the molecule. Note the use of doted lines to represent a bond projecting behind the plane of the paper and a solid wedge to represent a bond projecting forward from the plane of the paper.

b. The Lewis structure of acetonitrile, CH₃CN is:

The methyl group, CH₃-, is tetrahedral. The carbon of the -CN group is in the middle of a straight line stretching from the carbon of the methyl group through the nitrogen.

c. The Lewis structure of acetic acid is:

Both the carbon bonded to three hydrogens and the oxygen bonded to carbon and hydrogen are centers of tetrahedral structures. The central carbon will have 120 7deg bond angles.

The geometry around the first carbon is tetrahedral, around the second carbon atom is trigonal planar, and around the oxygen is bent.

Module Four: Part L Bond Angles 05 points

What is the bond Angle in the following structures:

___1. ___4.

___2. ___5.

___3. ___6.

___7. ___8. ____9.

___10. ____11.

___12. ___13. ___16.

___14.

___15.

___17. ___20.

Bonus:

___21. ___23.

___22. ___24.

Steric Numbers do not predict bond angles within rings of carbons

Types of molecular structure

Main article: VSEPR theory #AXE method

Some common shapes of simple molecules include:

• Linear: In a linear model, atoms are connected in a straight line. The bond angles are set at 180°. A bond angle is very simply the geometric angle between two adjacent bonds. For example, carbon dioxide and nitric oxide have a linear molecular shape.
• Trigonal planar: Just from its name, it can easily be said that molecules with the trigonal planar shape are somewhat triangular and in one plane (flat). Consequently, the bond angles are set at 120°. An example of this is boron trifluoride.
• Bent: Bent or angular molecules have a non-linear shape. A good example is water, or H₂O, which has an angle of about 105°. A water molecule has two pairs of bonded electrons and two unshared lone pairs.
• Tetrahedral: Tetra- signifies four, and -hedral relates to a face of a solid, so "tetrahedral" literally means "having four faces". This shape is found when there are four bonds all on one central atom, with no extra unshared electron pairs. In accordance with the VSEPR (valence-shell electron pair repulsion theory), the bond angles between the electron bonds are arccos(−1/3) = 109.47°. An example of a tetrahedral molecule is methane (CH₄).
• Octahedral: Octa- signifies eight, and -hedral relates to a face of a solid, so "octahedral" literally means "having eight faces". The bond angle is 90 degrees. An example of an octahedral molecule is sulfur hexafluoride (SF₆).
• Pyramidal: Pyramidal-shaped molecules have pyramid-like shapes. Unlike the linear and trigonal planar shapes but similar to the tetrahedral orientation, pyramidal shapes require three dimensions in order to fully separate the electrons. Here, there are only three pairs of bonded electrons, leaving one unshared lone pair. Lone pair – bond pair repulsions change the angle from the tetrahedral angle to a slightly lower^{[citation needed]} value. An example is NH₃ (ammonia).

Common shapes you should know

There are a whole bunch of common shapes you need to know to accurately think of covalent molecules.  Here they are:

• Tetrahedral:  Tetrahedral molecules look like pyramids with four faces.  Each point on the pyramid corresponds to an atom that's attached to the central atom.  Bond angles are 109.5 degrees.
• Trigonal pyramidal:  It's like a tetrahedral molecule, except flatter.  It looks kind of like a squished pyramid because one of the atoms in the pyramid is replaced with a lone pair.  Bond angles are 107.5 degrees (it's less than tetrahedral molecules because the lone pair shoves the other atoms closer to each other).
• Trigonal planar:  It looks like the hood ornament of a Mercedes automobile, or like a peace sign with that bottom-most line gone.  The bond angles are 120 degrees.
• Bent:  They look, well, bent.  Bond angles can be either 118 degrees for molcules with one lone pair or 104.5 degrees for molecules with two lone pairs.
• Linear:  The atoms in the molecule are in a straight line.  This can be either because there are only two atoms in the molecule (in which case there is no bond angle, as there need to be three atoms to get a bond angle) or because the three atoms are lined up in a straight line (corresponding to a 180 degree bond angle).
• There are other types, but we won't worry about them.

Module Four- Part N: Geometry of Molecules 5 points

Use the dot/stick structures on the Part L page to state the geometry of the molecules:

Bent Linear Trigonal Planer Planer Trigonal Pyramidal Tetrahedral

Trigonal-bipyramidal Square Planer Seesaw T-shaped Octahedral

_____________1. H2O _____________2. CO2 _____________3. C2H4 _____________4. SO2 _____________5. SO3 _____________6. HCN _____________7. CH4 _____________8. NH3 _____________9. CH2O _____________10. C2H2 _____________Bonus. PF5 _____________Bonus SF6

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Module Four II- Part O: Polarity of Molecules 5 points

Use the dot/stick structures and sketch the molecule in three dimensions. Then draw the dipoles for each bond to state if the molecule is polar or nonpolar:

Electronegativities: F=4.0; O=3.5; N=3.0; Cl=3.0; Br=2.7; C=2.5; S=2.5; P=2.1; H=2.1

_____________1. H2O _____________2. CO2 _____________3. C2H4 _____________4. C2H2 _____________5. SO2 _____________6. SO3 _____________7. CH4 _____________8. NH3 _____________9. BH3 _____________10. HCN _____________Bonus. PCl5 _____________Bonus SCl6

Polar Covalent Bonds

• Covalent bonds result from the sharing of valence electrons.

• Often, the two atoms do not share the electrons equally. That is, one of the atoms holds onto the electrons more tightly than the other.

• When one of the atoms holds the shared electrons more tightly, the bond is polarized.

A polar covalent bond is one in which the electrons are not shared equally

Electronegativity

• Each element has an innate ability to attract valence electrons.

• Electronegativity is the ability of an atom to attract electrons in a chemical bond.

• Linus Pauling devised a method for measuring the electronegativity of each of the elements.

• Fluorine is the most electronegative element.

• Electronegativity increases as you go left to right across a period.

Electronegativity increases as you go from bottom to top in a family.

Electronegativity: The ability of an atom in a molecule to attract the shared electrons in a covalent bond.

Electronegativity Differences

• The electronegativity of H is 2.1; Cl is 3.0.

• Since there is a difference in electronegativity between the two elements (3.0 – 2.1 = 0.9), the bond in H – Cl is polar.

• Since Cl is more electronegative, the bonding electrons are attracted toward the Cl atom and away from the H atom. This will give the Cl atom a slightly negative charge and the H atom a slightly positive charge.

Nonpolar Covalent Bonds

• What if the two atoms in a covalent bond have the same or similar electronegativities?

• The bond is not polarized and it is a nonpolar covalent bond. If the electronegativity difference is less than 0.5, it is usually considered a nonpolar bond.

• The diatomic halogen molecules have nonpolar covalent bonds.

Hybrid Orbitals

Module 4II- Part P: Hybrid Orbitals of Molecules 5 points

Use the dot/stick structures in the table to predict the hybrid orbitals that overlap to form the covalent bond:

_____________1. H2O the sigma bond between either H-O _____________2. CO2 the sigma bond between either C=O _____________2a. CO2 the pi bond between either C=O _____________3. C2H4 the sigma bond between the two carbons _____________3a. C2H4 the pi bond between the two carbons _____________4. SO2 the sigma bond between S=O _____________4a. SO2 the pi bond between S=O ____________4b. SO2 the sigma bond between S-O ____________ 5. SO3 either sigma bond between S-O ____________ 5a. SO3 the sigma bond between S=O ____________ 5b. SO3 the pi bond between S=O ____________6. HCN the single (sigma) bond between H-C ____________6a. HCN the sigma bond between C=N ____________6b. HCN either pi bond between C=N ____________7. CH4 the sigma bond between any of the H-C ____________8. NH3 the sigma bond between any of the H-N ___________ 9. CH2O either sigma bond between the H-C ____________9a. CH2O the sigma bond between the C=O ____________9b. CH2O the pi bond between the C=O ___________ __10. C2H2 the sigma bond between the C=C _________ ____10a. C2H2 either pi bond between the C=C _____________ 10b. C2H2 the sigma bond between any of the H-C

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