Chapter 6: How Atoms Bond

Part A: Dot Structures Atoms Answers
Part B: Ionic Character Answers
Part C: Bond Recognition Answers
Part D: Dot Structures of Molecules Answers
Part E: Bond Angles (Octet Rule-Steric Numbers 2-3-4 Answers
Part F: Molecular Geometry (Octet Rule) Steric Numbers 2-3-4 Answers
Part G: Polarity(Octet Rules) Steric Numbers 2-3-4
Part H: Polar/Nonpolar Molecules (Octet Rules) Steric Numbers 2-3-4 Answers
Part V: Chapter 6 Vocabulary p187 Answers

Part M: Chapter 6 Multiple Choice (Blackboard - Course Content
-Required Path 2 MC Quizzes – Chapter 6)
Part Z: Conceptual Chemistry Spotlight: Toxic Wastes and Superfund Act 193-194

Part A: Dot Structures Atoms

See:
Chapter 43 Study Pack Part P1: Electron Dot Structures using the Periodic Table Answers

Part B: Ionic Character
See: Chapter 3 Study Pack Part P2: Periodic Ionic Character using the Periodic Table Answers

Part C: Bond Recognition

There are three types of chemical bonds between two elements:
Ionic, Covalent, and Metallic. 

In General:

Metal-Metal = Metallic Bond (example: Ag(5)-Au(14)-Cu(5) = 14 Karat Gold)

Metal-Nonmetal = Ionic Bond (example: Na-Cl)

Nonmetal-nonmetal = Covalent Bond (example: H₂O)

There is a more exact way to predict if two atoms will transfer their electrons or share their electron in pairs making a compound. In Suchocki’s 5^th edition read carefully section 6.7 pages 180-182. See Figure 6.27 page 181 Periodic table of Electronegativity showing the electronegativity of each element on the periodic chart.

 If the difference in electronegativity between two atoms is greater than 1.7 (Suchocki), 1.8 (Corwin), or(2.0) Hei,n the electrons will transfer from one atom to the other to make ions and Ionic Compounds.

Ionic (sometimes called Electrovalent) Compounds are also called salts and in nature they are called minerals and in Sports medicine Body Electrolytes. We will over simplify this concept to say if a metal meets a nonmetal, ionic bonds are formed (Just a Rule of Thumb) if we do not have access to the Table of electronegativity. Hein (14^th) states on page 227 if the difference in electronegativity is greater than 2.0 the bonding is strongly ionic, while less than 1.5 strongly covalent. Then he states between 1.7-1.9 the bonding will be more ionic than covalent.

 For this course, if the difference between the electronegativity of two atoms is less than 1.7 then the two atoms will share electrons in pairs. Two types of sharing bonds are formed.

Metallic and Covalent.

Metallic Bonds are formed when two metals share electrons such as alloys of metals. 24 karat gold is pure gold and is very soft. But Jewelry is usually 10-18 Karat Gold, meaning that another metal is mixed with gold to make the solid harder. We will not study Metallic Bonds in this course, but you should know that two metals share electrons in pairs to make Metallic Bonds.

“Metallic bonding occurs as a result of electromagnetism and describes the electrostatic attractive force that occurs between conduction electrons (in the form of an electron cloud of delocalized electrons) and positively charged metal ions. It may be described as the sharing of free electrons among a lattice of positively charged ions (cations). In a more quantum-mechanical view, the conduction electrons divide their density equally over all atoms that function as neutral (non-charged) entities.^{[citation needed]} Metallic bonding accounts for many physical properties of metals, such as strength, ductility, thermal and electrical resistivity and conductivity, opacity, and luster.^[1][2][3][4]

Metallic bonding is not the only type of chemical bonding a metal can exhibit, even as a pure substance. For example, elemental gallium consists of covalently-bound pairs of atoms in both liquid and solid state—these pairs form a crystal lattice with metallic bonding between them. Another example of a metal–metal covalent bond is mercurous ion (Hg₂⁺²).“

Covalent Bonds are formed when two nonmetals bond together. The elements carbon, oxygen, hydrogen, sulfur, nitrogen, phosphorus, chlorine, and bromine will be the main nonmetals studied in drawing dot structures of molecules. Bonds between these nonmetals are always Covalent.

Part D of Chapter 6 should now be easy. Predict what type of bond will be made if two atoms combine:

Chapter 6: Part C Sample Bond Recognition Exam

Using a periodic chart (Rule of Thumb), predict the bond that would form between the two elements:

1. Fe-Al ________________

2. P-S ________________

3. C-O ________________

4. B-Cl ________________

5. K-I ________________

For the following element pairs use the electronegativity table below to determine if the bond is ionic or covalent.

6. Na-P ________________

7. Ca-Br ________________

8. Ge-O ________________

9. P-H ________________

10. Be-Cl ________________

Chapter 6: Part D Dot Structures of Molecules

Section 6.5 pages 173-177 describe how two nonmetal atoms bond by sharing electron pairs. Every text has a mathematical set of rules to construct dot structures of molecules. John Suchocki has this mathematical approach by counting total valence electrons.

The Bishop text has a nice summary of the bonding variations of the number of covalent bonds possible between two nonmetallic atoms bondin:.

Dot Structure Table from Bishop Text:

John Taylor’s Method for Drawing Dot Structures

On John Taylor’s web site, he has a lengthy study guide for Polyatomic ions:
http://www.fccj.us/PolyatomicIons/PolyatomicIonsIntro.htm

From that study guide he modifies the mathematical approach given in each of the books on the previous pages. The following are his seven steps:

You can use paper atoms to construct dot structures of molecules:

Dot/Stick Structures of Atoms: Dot Structures of Atoms:
Oxygen-Carbon Dot/Stick Atoms O, H, S, e-1 atoms
Hydrogen-Chlorine-Nitrogen Dot/Stick Atoms
Hydrogen-Phosphorus-Sulfur Dot/Stick Atoms P, N, Cl, e-1 atoms
Oxygen-Hydrogen-Carbon-Chlorine Dot Stick Atoms


Place the nonmetal which is not oxygen or hydrogen in the middle of your desk.

1. First using all single bonds, hook all oxygens in the formula to the central nonmetal (Simple covalent or coordinate covalent bonds).

Never hook oxygen to oxygen except in peroxides, O₂, O₃.

2. If oxygen is present, hook the hydrogen written first in the chemical formula to an oxygen. Hydrogen requires only two electrons to fill it's orbital. Notice the change in the polyion's charge when a hydrogen is placed on the oxygen. If hydrogen is written second in the formula after another nonmetal, then hook that hydrogen to that nonmetal, not to oxygen such as (CH₃-COOH written organically) acetic acid HC₂H₃O₂ or oxalic acid H₂C₂O₄ both hydrogens are attached to the oxygens.
3. If using all single bonds connecting the oxygen to the central nonmetal, the count of electrons around each element should total eight electrons (octet rule). This includes the element's original outer surface (valence) electrons plus the electrons being shared from the bonded element.

4. If the count is 7-7, then add a second bond, a double covalent bond (four electrons being shared between two atoms.

5. If the count is 6-6, then make a triple covalent bond between the two elements (six electrons shared between the two atoms).
   

6. If the count is 8-6, (or 9-7) then make a coordinate covalent bond. Hook the six (vacant) orbital onto an unshared pair of the eight. A coordinate covalent bond is still a single bond. In making double, triple bonds you may also use one or two coordinate covalent bonds to predict a structure using the octet rule, if necessary as in carbon monoxide.
   

Extras: Never have an unshared pair or lone pair of electrons on a carbon atom (except carbon monoxide, CO). These are two dimensional structures, so there are many variations of the answer shown on the web site. Never have more than two bonds to any oxygen (except CO). If you place hydrogen to an oxygen, then oxygen HAS to hook to another element by a single bond, never a double bond.

Table 9.5 from Kotz displays common oxoacids and their anions:

Chapter 6: Part D Dot Structures of Molecules

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 )

Part E: Bond Angles (Octet Rule-Steric Numbers 2-3-4

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.

Chapter 6: Part E Bond Angles

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.

___21. ___23.

___22. ___24.

Steric Numbers do not predict bond angles within rings of carbons

Part F: Molecular Geometry (Octet Rule) Steric Numbers 2-3-4

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).

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.

Chapter 6- Part F: Geometry of Molecules
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

.

Part G: Polarity(Octet Rules) Steric Numbers 2-3-4
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.

Part H: Polar/Nonpolar Molecules (Octet Rules) Steric Numbers 2-3-4

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