CHM 1032C Tentative Grading Outline Fall 2015

Chapter 4: Molecular Compounds Homework Packet

B. _____(04*) Dot Structures of Molecules-Section 4.7 Answers

C. _____(02) Binary Molecular(Covalent) Compounds-Section 4.11 Answers

L. ____ (02) Bond Angles/Bond Lengths-Section 4.8 Answers

N. ____ (02) Geometry of Molecules-Section 4.8 Answers

O. ____ (02) Polarity of Molecules-Section 4.10 Answers

H. _____(03) Inorganic Compounds Sections 3.7, 3.8,3.9, 3.10, 3.11, 4.11 Answers

_______(15) Chapter 4 Total

Chapter 4. Molecular Compounds Table of Contents
4.1 Covalent Bonds
4.2 Covalent Bonds and the Periodic Table
4.3 Multiple Covalent Bonds
4.4 Coordinate Covalent Bonds
4.5 Characteristics of Molecular Compounds
4.6 Molecular Formulas and Lewis Structures
4.7 Drawing Lewis Structures
4.8 The Shapes of Molecules
4.9 Polar Covalent Bonds and Electronegativity
4.10 Polar Molecules
4.11 Naming Binary Molecular Compounds

See lab handout and McMurry GOB Section 4.7 (or Corwin Chapter 12 sections 12.4-12.5 for directions) or (Hein Chapter 11 Sections 11.7, 11.8, 11.9)

Submit the following dot structures as a separate Post Lab Report 10 points

Drag and Drop Interactive Web Site (Nothing required to turn in):

Check answers:

http://www.lsua.us/chem1001/dragdrop/menu.html

Module Four: Part B Dot Structures of Molecules 4 points

Using a periodic chart draw the electron dot structures of the following molecules:

(Choose One for each question or the one circled on the paper)

1. NH₃ CH₄ H₂O₂ H₂O 2. H₂SO₄ H₃PO₄ HClO₄ HClO₃

Submit these dot structures as a separate homework

3. HNO₃ H₂CO₃ HNO₂ 4. CO₂ HCN SO₃ SO₂

Submit these dot structures as a separate homework

5. HC₂H₃O₂ H₂C₂O₄ HCHO₂ 6. C₂H₄ C₂H₂ C₃H₈ C₂H₆

carbon to carbon by single covalent bond bond carbons to carbon

Submit these dot structures as a separate homework

7. CH₃CH₂OH CH₃COCH₃ CH₂O (HCHO)

(carbon to carbons by single covalent bonds-oxygen attach to carbon)

Submit these dot structures as a separate homework

8. CH₃OCH₃ CHONH₂ CH₃CH₂CH₂OH CH₃CHOHCH₃
oxygen separates the carbons O & N both bond to C (all three carbons single bonded and –OH attached to carbon)

Submit these dot structures as a separate homework

9. CH₂NH₂COOH CH₃CHNH₂COOH

carbon to carbons by single covalent bonds (-NH₂ amino on#2 carbon in both above)

Submit these dot structures as a separate homework

10. CH₃COOCH₂CH₃ HCOOCH₃

(-CH₂CH₃ also hooks to oxygen in#10, as well as - CH₃ )

Binary Molecular (Covalent) Compounds-Section 4.11

Binary Molecular compounds are explained after the ionic compounds in Chapter 4 GOB McMurry Section 4.11 (Corwin (7^th) Chapter 6 section 6.7, and inorganic acids are not covered until last in the chapter, sections 6.8 and 6.9) ( Hein includes Binary Molecular at the end of section 6.4 [Page 108 14^th] covering all Binary Compounds first ionic, then molecular).

The required Online Binary Covalent Molecular Homework 2 points each

The web site is:
C: Binary Molecular Names:
http://www.northcampus.net/Nomenclature/Molecules/32BinaryCovalent.html

C1: Binary Molecular Formulas:
http://www.northcampus.net/Nomenclature/MoleculeFormula/32BinaryMolecularFormula.html

Here is a brief tutorial for Part C:

PART C: BINARY COVALENT COMPOUNDS

 Both elements are nonmetals attached by covalent bonds.  These bonds may be single, double, or triple covalent.  Due to the covalent bonding there are many ratios of the same two elements making many different compounds_.  For this reason, the chemist states how many atoms of each element is present in the chemical formula in the formal name of the compound.

Prefixes are attached to each element to indicate how many.  Each student should learn the following prefixes:

MONO            =          ONE                                        HEXA              =          SIX

DI                    =          TWO                                       HEPTA            =          SEVEN

TRI                  =          THREE                                   OCTA             =          EIGHT

TETRA            =          FOUR                                      NONA            =          NINE

PENTA            =          FIVE                                        DECA              =          TEN

The element that is shown first in the chemical formula is written first using the proper prefix to indicate how may atoms of that element is contained in the compound.  If there is only one atom of that element it is often found without the prefix mono.  If you leave the prefix off then it is understood that you mean mono.

The element which is written second in the chemical formula is written second in the chemical name, but in addition to the prefix indicating how many, the suffix of the element’s name is changed to -ide. 

carbon becomes carbide                             chlorine becomes chloride

sulfur becomes sulfide                                oxygen becomes oxide

hydrogen becomes hydride                        nitrogen becomes nitride

Therefore, the following formulas of binary compounds would be spoken:

CCl₄                 carbon tetrachloride                                                     

SO₂                  sulfur dioxide

CO₂                 carbon dioxide

N₂O₃ dinitrogen trioxide

 BH₃                 boron trihydride

We use common names for NH₃, and H₂O. What would be their correct binary molecular names? Methane, CH₄, is the organic name for CH₄, what would its inorganic name be?

Chapter 4: Part C Binary Molecular Compounds 2 points

Using a periodic chart write the names or formulas of the following compounds depending on whether the formula or name is given:

Homework Packet Sample test: answer on grading outline

1. CO ____________________

2. SO₃ _____________________

3. N₂O₅ _____________________

4. N₂O₇ _____________________

5. N₂O _____________________

6. Phosphorus pentachloride _________

7. Boron trifluoride _________

8. Carbon dioxide _________

9. Sulfur Trioxide _________

10. Carbon Tetrachloride _________

Textbook Reference: McMurry Section 4.11 (Corwin Chapter 6 Section 6.7) (Hein Section 6.4 page 108)

Online Homework (2 Points Each Required):

C: Binary Molecular (Covalent) Homework: http://www.northcampus.net/Nomenclature/Molecules/25BinaryCovalent.html

C1. Binary Molecular (Covalent) Formulas: http://www.northcampus.net/Nomenclature/MoleculeFormula/25BinaryMolecularFormula.html

Submit grades on separate grading Sheet (goldenrod) when taking Chapter 4 Exam

Online Study Guide:
http://www.fccj.us/chm1025/AssignmentOutline/M4PartC.htm

Chapter 4 Part L: VSEPR and Bond Angles

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 CHM 1032C	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 CHM 2045C	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	r7r	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.

Chapter 4: Part L Bond Angles 02 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

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).

CHM 1032C 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.

Trigonal-bipyramidal Square Planer Seesaw T-shaped Octahedral

Module Four- Part N: Geometry of Molecules 2 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

.

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.

Chapter 4- Part O: Polarity of Molecules 2 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

Chapter 4: Part H Inorganic Compounds 3 points

The key to deciding which system to use in Part H is to look at the element written first.

 1. If a Metal is written first (or a polyatomic ion), then use the rules for ionic compounds (salts).

2. If a nonmetal is written first, then use the Covalent/Molecule System with prefixes. (If the compound is Organic Nomenclature of Organics is covered in Chapter 11, but for now use the prefix system of binary molecular nomenclature.

3. If hydrogen is written first (and it is in aqueous solution) then name it as an Acid

Chapter 4: Part H Inorganic Compounds 3 points

Using a periodic chart write the names or formulas of the following polyatomic ions depending on whether the formula or name is given:

1. H₂CO₃ _____________________

2. MgSO₄ _____________________

3. Ca₃(PO₃)₂ ____________________

4. HClO₃ _____________________

5. SO₃ _____________________

6. Fe₂O₃ _____________________

7. Aluminum Hydroxide ____________

8. Ammonium chloride _____________

9. Sodium Hypochlorite _____________

10. Nitrogen dioxide _____________

11. Calcium Phosphate ____________

12. Sulfuric acid ____________

Online Homework (2 Points Each Required):

H: Inorganic Compound Names Homework: http://www.northcampus.net/Nomenclature/Inorganic/32inorganic.html

H1. Inorganic Compound Formulas:

http://www.northcampus.net/Nomenclature/InorganicFormula/25inorganicFormula.html

Submit grades on separate grading Sheet when taking Chapter-4 Exam

Online Study Guide:

http://www.fccj.us/chm1025/AssignmentOutline/M4PartH.htm