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0 | CHAPTER 1 <br> Structure and Bonding -
FIGURE 1.1 The enzyme HMG-CoA reductase, shown here as a so-called ribbon model, catalyzes a crucial step in the body's synthesis of cholesterol. Understanding how this enzyme functions has led to the development of drugs credited with saving millions of lives. (credit: image fro... |
1 | CHAPTER CONTENTS -
1.1 Atomic Structure: The Nucleus
1.2 Atomic Structure: Orbitals
1.3 Atomic Structure: Electron Configurations
1.4 Development of Chemical Bonding Theory
1.5 Describing Chemical Bonds: Valence Bond Theory
$1.6 s p^{3}$ Hybrid Orbitals and the Structure of Methane
$1.7 s p^{3}$ Hybrid Orbitals and th... |
2 | CHAPTER CONTENTS -
Oxycodone (OxyContin)
Cholesterol
Benzylpenicillin
Historically, the term organic chemistry dates to the mid-1700s, when it was used to mean the chemistry of substances found in living organisms. Little was known about chemistry at that time, and the behavior of the "organic" substances isolate... |
3 | CHAPTER CONTENTS -
| Grou 1A | | | | | | | | | | | | | | | | | 8A |
| :---: | :---: | :---: | :---: | :---: | :---: | :---: | :---: | :---: | :---: | :---: | :---: | :---: | :---: | :---: | :---: | :---: | :---: |
| H | 2A | | | | | | | | | | | 3A | 4A | 5A | 6A | 7A | He |
| Li | Be | | | ... |
4 | CHAPTER CONTENTS - 1.1 Atomic Structure: The Nucleus
As you might remember from your general chemistry course, an atom consists of a dense, positively charged nucleus surrounded at a relatively large distance by negatively charged electrons (FIGURE 1.3). The nucleus consists of subatomic particles called neutrons, whic... |
5 | CHAPTER CONTENTS - 1.1 Atomic Structure: The Nucleus
A specific atom is described by its atomic number ( $Z$ ), which gives the number of protons (or electrons) it contains, and its mass number ( $\boldsymbol{A}$, which gives the total number of protons and neutrons in its nucleus. All the atoms of a given element have... |
6 | CHAPTER CONTENTS - 1.2 Atomic Structure: Orbitals
How are the electrons distributed in an atom? You might recall from your general chemistry course that, according to the quantum mechanical model, the behavior of a specific electron in an atom can be described by a mathematical expression called a wave equation-the sam... |
7 | CHAPTER CONTENTS - 1.2 Atomic Structure: Orbitals
The orbitals in an atom are organized into different layers around the nucleus called electron shells, which are centered around the nucleus and have successively larger size and energy. Different shells contain different numbers and kinds of orbitals, and each orbital ... |
8 | CHAPTER CONTENTS - 1.2 Atomic Structure: Orbitals
A $2 p_{\mathrm{x}}$ orbital
A $2 p_{\mathrm{y}}$ orbital
A $\mathbf{2} p_{z}$ orbital
FIGURE 1.6 Shapes of the $2 p$ orbitals. Each of the three mutually perpendicular, dumbbell-shaped orbitals has two lobes separated by a node. The two lobes have different algebr... |
9 | CHAPTER CONTENTS - 1.3 Atomic Structure: Electron Configurations
The lowest-energy arrangement, or ground-state electron configuration, of an atom is a list of the orbitals occupied by its electrons. We can predict this arrangement by following three rules. |
10 | RULE 1 -
The lowest-energy orbitals fill up first, $1 s \rightarrow 2 s \rightarrow 2 p \rightarrow 3 s \rightarrow 3 p \rightarrow 4 s \rightarrow 3 d$, according to the following graphic, a statement called the Aufbau principle. Note that the $4 s$ orbital lies between the $3 p$ and $3 d$ orbitals in energy.
RULE ... |
11 | RULE 1 - 1.4 Development of Chemical Bonding Theory
By the mid-1800s, the new science of chemistry was developing rapidly, especially in Europe, and chemists had begun to probe the forces holding compounds together. In 1858, the German chemist August Kekulé and the Scottish chemist Archibald Couper independently propos... |
12 | RULE 1 - 1.4 Development of Chemical Bonding Theory
Why, though, do atoms bond together, and how can chemical bonds be described electronically? The why question is relatively easy to answer: atoms bond together because the compound that results is more stable and lower in energy than the separate atoms. Energy-usually... |
13 | RULE 1 - 1.4 Development of Chemical Bonding Theory
But how do elements closer to the middle of the periodic table form bonds? Look at methane, $\mathrm{CH}_{4}$, the main constituent of natural gas, for example. The bonding in methane is not ionic because it would take too much energy for carbon ( $1 s^{2} 2 s^{2} 2 p... |
14 | RULE 1 - 1.4 Development of Chemical Bonding Theory
| Electron-dot structures (Lewis structures) | <smiles></smiles> | $\begin{gathered} H: \ddot{\mathrm{N}}: \mathrm{H} \\ \ddot{\mathrm{H}} \end{gathered}$ | $H: \ddot{O}: H$ | <smiles></smiles> |
| :---: | :---: | :---: | :---: | :---: |
| Line-bond structures (Kekulé... |
15 | RULE 1 - 1.4 Development of Chemical Bonding Theory
Valence electrons that are not used for bonding remain as dots in structures and are called lone-pair electrons, or nonbonding electrons. The nitrogen atom in ammonia, $\mathrm{NH}_{3}$, for instance, shares six valence electrons in three covalent bonds and has its re... |
16 | Predicting the Number of Bonds Formed by Atoms in Molecules -
How many hydrogen atoms does phosphorus bond to in forming phosphine, $\mathrm{PH}_{\text {? }}$ ? |
17 | Strategy -
Identify the periodic group of phosphorus, and find from that how many electrons (bonds) are needed to make an octet. |
18 | Solution -
Phosphorus is in group 5A of the periodic table and has five valence electrons. It thus needs to share three more electrons to make an octet and therefore bonds to three hydrogen atoms, giving $\mathrm{PH}_{3}$. |
19 | Drawing Electron-Dot and Line-Bond Structures -
Draw both electron-dot and line-bond structures for chloromethane, $\mathrm{CH}_{3} \mathrm{Cl}$. |
20 | Strategy -
Remember that a covalent bond-that is, a pair of shared electrons-is represented as a line between atoms. |
21 | Solution -
Hydrogen has one valence electron, carbon has four valence electrons, and chlorine has seven valence electrons. Thus, chloromethane is represented as
Chloromethane
PROBLEM Draw a molecule of chloroform, $\mathrm{CHCl}_{3}$, using solid, wedged, and dashed lines to show its tetrahedral 1-3 geometry.
PROB... |
22 | Ethane -
PROBLEM What are likely formulas for the following substances?
1-5 (a) $\mathrm{CCl}_{\text {? }}$
(b) AlH ?
(c) $\mathrm{CH}_{?} \mathrm{Cl}_{2}$
(d) SiF
(e) $\mathrm{CH}_{3} \mathrm{NH}_{?}$
PROBLEM Write line-bond structures for the following substances, showing all nonbonding electrons:
1-6 (a) $\mathrm{... |
23 | 1-7 - 1.5 Describing Chemical Bonds: Valence Bond Theory
How does electron sharing lead to bonding between atoms? Two models have been developed to describe covalent bonding: valence bond theory and molecular orbital theory. Each model has its strengths and weaknesses, and chemists tend to use them interchangeably depe... |
24 | 1-7 - 1.5 Describing Chemical Bonds: Valence Bond Theory
During the bond-forming reaction $2 \mathrm{H} \cdot \longrightarrow \mathrm{H} 2,436 \mathrm{~kJ} / \mathrm{mol}(104 \mathrm{kcal} / \mathrm{mol})$ of energy is released. Because the product $\mathrm{H}_{2}$ molecule has $436 \mathrm{~kJ} / \mathrm{mol}$ less en... |
25 | 1-7 - 1.5 Describing Chemical Bonds: Valence Bond Theory
How close are the two nuclei in the $\mathrm{H}_{2}$ molecule? If they are too close, they will repel each other because both are positively charged. Yet if they're too far apart, they won't be able to share the bonding electrons. Thus, there is an optimum distan... |
26 | $1.6 s p^{3}$ Hybrid Orbitals and the Structure of Methane -
The bonding in the hydrogen molecule is fairly straightforward, but the situation is more complicated in organic molecules with tetravalent carbon atoms. Take methane, $\mathrm{CH}_{4}$, for instance. As we've seen, carbon has four valence electrons $\left(2... |
27 | $1.6 s p^{3}$ Hybrid Orbitals and the Structure of Methane -
The concept of hybridization explains how carbon forms four equivalent tetrahedral bonds but not why it does so. The shape of the hybrid orbital suggests the answer to why. When an $s$ orbital hybridizes with three $p$ orbitals, the resultant $s p^{3}$ hybri... |
28 | $1.7 s p^{3}$ Hybrid Orbitals and the Structure of Ethane -
The same kind of orbital hybridization that accounts for the methane structure also accounts for the bonding together of carbon atoms into chains and rings to make possible many millions of organic compounds. Ethane, $\mathrm{C}_{2} \mathrm{H}_{6}$, is the si... |
29 | $1.8 s p^{2}$ Hybrid Orbitals and the Structure of Ethylene -
The bonds we've seen in methane and ethane are called single bonds because they result from the sharing of one electron pair between bonded atoms. It was recognized nearly 150 years ago, however, that carbon atoms can also form double bonds by sharing two e... |
30 | $1.8 s p^{2}$ Hybrid Orbitals and the Structure of Ethylene -
When two carbons with $s p^{2}$ hybridization approach each other, they form a strong $\sigma$ bond by $s p^{2}-s p^{2}$ head-on overlap. At the same time, the unhybridized $p$ orbitals interact by sideways overlap to form what is called a pi ( $\boldsymbol... |
31 | $1.8 s p^{2}$ Hybrid Orbitals and the Structure of Ethylene -
As you might expect, the carbon-carbon double bond in ethylene is both shorter and stronger than the single bond in ethane because it has four electrons bonding the nuclei together rather than two. Ethylene has a $\mathrm{C}=\mathrm{C}$ bond length of 134 p... |
32 | Drawing Electron-Dot and Line-Bond Structures -
Commonly used in biology as a tissue preservative, formaldehyde, $\mathrm{CH}_{2} \mathrm{O}$, contains a carbon-oxygen double bond. Draw electron-dot and line-bond structures of formaldehyde, and indicate the hybridization of the carbon orbitals. |
33 | Strategy -
We know that hydrogen forms one covalent bond, carbon forms four, and oxygen forms two. Trial and error, combined with intuition, is needed to fit the atoms together. |
34 | Solution -
There is only one way that two hydrogens, one carbon, and one oxygen can combine:
Like the carbon atoms in ethylene, the carbon atom in formaldehyde is in a double bond and its orbitals are therefore $s p^{2}$-hybridized.
PROBLEM Draw a line-bond structure for propene, $\mathrm{CH}_{3} \mathrm{CH}=\mathr... |
35 | 1.9 sp Hybrid Orbitals and the Structure of Acetylene -
In addition to forming single and double bonds by sharing two and four electrons, respectively, carbon can also form a triple bond by sharing six electrons. To account for the triple bond in a molecule such as acetylene, $\mathrm{H}-\mathrm{C} \equiv \mathrm{C}-\... |
36 | 1.9 sp Hybrid Orbitals and the Structure of Acetylene -
FIGURE 1.17 The structure of acetylene. The two carbon atoms are joined by one $s p-s p \sigma$ bond and two $p-p \pi$ bonds.
As suggested by $s p$ hybridization, acetylene is a linear molecule with $\mathrm{H}-\mathrm{C}-\mathrm{C}$ bond angles of $180^{\circ}$.... |
37 | 1.9 sp Hybrid Orbitals and the Structure of Acetylene -
| Molecule | Bond | Bond strength | | Bond length (pm) |
| :---: | :---: | :---: | :---: | :---: |
| | | (kJ/mol) | ( $\mathrm{kcal} / \mathrm{mol}$ ) | |
| Methane, $\mathrm{CH}_{4}$ | $\left(s p^{3}\right) \mathrm{C}-\mathrm{H}$ | 439 | 105 | 109 |
| Ethane... |
38 | 1.9 sp Hybrid Orbitals and the Structure of Acetylene - 1.10 Hybridization of Nitrogen, Oxygen, Phosphorus, and Sulfur
The valence-bond concept of orbital hybridization described in the previous four sections is not limited to carbon. Covalent bonds formed by other elements can also be described using hybrid orbitals. ... |
39 | 1.9 sp Hybrid Orbitals and the Structure of Acetylene - 1.10 Hybridization of Nitrogen, Oxygen, Phosphorus, and Sulfur
In the periodic table, phosphorus and sulfur are the third-row analogs of nitrogen and oxygen, and the bonding in both can be described using hybrid orbitals. Because of their positions in the third ro... |
40 | 1.9 sp Hybrid Orbitals and the Structure of Acetylene - 1.10 Hybridization of Nitrogen, Oxygen, Phosphorus, and Sulfur
Methanethiol
Dimethyl sulfide
PROBLEM Identify all nonbonding lone pairs of electrons in the following molecules, and tell what geometry
1-14 you expect for each of the indicated atoms.
(a) The oxyg... |
41 | 1.9 sp Hybrid Orbitals and the Structure of Acetylene - 1.11 Describing Chemical Bonds: Molecular Orbital Theory
We said in Section 1.5 that chemists use two models for describing covalent bonds: valence bond theory and molecular orbital theory. Having now seen the valence bond approach, which uses hybrid atomic orbita... |
42 | 1.9 sp Hybrid Orbitals and the Structure of Acetylene - 1.11 Describing Chemical Bonds: Molecular Orbital Theory
The additive combination is lower in energy than the two hydrogen $1 s$ atomic orbitals and is called a bonding MO because electrons in this MO spend most of their time in the region between the two nuclei, ... |
43 | 1.9 sp Hybrid Orbitals and the Structure of Acetylene - 1.12 Drawing Chemical Structures
Let's cover just one more point before ending this introductory chapter. In the structures we've been drawing until now, a line between atoms has represented the two electrons in a covalent bond. Drawing every bond and every atom i... |
44 | RULE 1 -
Carbon atoms aren't usually shown. Instead, a carbon atom is assumed to be at each intersection of two lines (bonds) and at the end of each line. Occasionally, a carbon atom might be indicated for emphasis or clarity. |
45 | RULE 2 -
Hydrogen atoms bonded to carbon aren't shown. Because carbon always has a valence of 4, we mentally supply the correct number of hydrogen atoms for each carbon. |
46 | RULE 3 -
Atoms other than carbon and hydrogen are shown.
One further comment: Although such groupings as $-\mathrm{CH}_{3},-\mathrm{OH}$, and $-\mathrm{NH}_{2}$ are usually written with the C , O , or N atom first and the H atom second, the order of writing is sometimes inverted to $\mathrm{H}_{3} \mathrm{C}-, \mathrm... |
47 | Interpreting a Line-Bond Structure -
Carvone, a substance responsible for the odor of spearmint, has the following structure. Tell how many hydrogens are bonded to each carbon, and give the molecular formula of carvone.
Carvone |
48 | Strategy -
The end of a line represents a carbon atom with 3 hydrogens, $\mathrm{CH}_{3}$; a two-way intersection is a carbon atom with 2 hydrogens, $\mathrm{CH}_{2}$; a three-way intersection is a carbon atom with 1 hydrogen, CH ; and a four-way intersection is a carbon atom with no attached hydrogens.
Solution
Ca... |
49 | Organic Foods: Risk versus Benefit -
Contrary to what you may hear in supermarkets or on television, all foods are organic-that is, complex mixtures of organic molecules. Even so, when applied to food, the word organic has come to mean an absence of synthetic chemicals, typically pesticides, antibiotics, and preservat... |
50 | Organic Foods: Risk versus Benefit -
Atrazine
How can the potential hazards from a chemical like atrazine be determined? Risk evaluation of chemicals is carried out by exposing test animals, usually mice or rats, to the chemical and then monitoring the animals for signs of harm. To limit the expense and time needed, t... |
51 | Organic Foods: Risk versus Benefit -
So, should we still use atrazine? All decisions involve tradeoffs, and the answer is rarely obvious. Does the benefit of increased food production outweigh possible health risks of a pesticide? Do the beneficial effects of a new drug outweigh a potentially dangerous side effect in ... |
52 | Key Terms -
- antibonding MO
- atomic number ( $Z$ )
- Aufbau principle
- bond angle
- bond length
- bond strength
- bonding MO
- condensed structure
- covalent bond
- electron shell
- electron-dot structure
- ground-state electron configuration
- Hund's rule
- ionic bond
- isotope
- Kekulé structure
- Lewis structure... |
53 | Summary -
The purpose of this chapter has been to get you up to speed-to review some ideas about atoms, bonds, and molecular geometry. As we've seen, organic chemistry is the study of carbon compounds. Although a division into organic and inorganic chemistry occurred historically, there is no scientific reason for the... |
54 | Summary -
Organic molecules are usually drawn using either condensed structures or skeletal structures. In condensed structures, carbon-carbon and carbon-hydrogen bonds aren't shown. In skeletal structures, only the bonds and not the atoms are shown. A carbon atom is assumed to be at the ends and at the junctions of l... |
55 | WHY YOU SHOULD WORK PROBLEMS -
There's no surer way to learn organic chemistry than by working problems. Although careful reading and rereading of this text are important, reading alone isn't enough. You must also be able to use the information you've read and be able to apply your knowledge in new situations. Working... |
56 | Visualizing Chemistry -
PROBLEM Convert each of the following molecular models into a skeletal structure, and give the formula of
1-18 each. Only the connections between atoms are shown; multiple bonds are not indicated (black = C, red $=0$, blue $=\mathrm{N}$, gray $=\mathrm{H}$ ) .
(a)
Coniine (the toxic substance... |
57 | Electron Configurations -
PROBLEM How many valence electrons does each of the following dietary trace elements have?
1-22
(a) Zinc (b) Iodine (c) Silicon (d) Iron
PROBLEM Give the ground-state electron configuration for each of the following elements:
1-23
(a) Potassium (b) Arsenic
(c) Aluminum
(d) Germanium |
58 | Electron-Dot and Line-Bond Structures -
PROBLEM What are likely formulas for the following molecules?
1-24 (a) $\mathrm{NH}_{?} \mathrm{OH}$
(b) $\mathrm{AlCl}_{\text {? }}$
(c) $\mathrm{CF}_{2} \mathrm{Cl}_{?}$
(d) $\mathrm{CH}_{?} \mathrm{O}$
PROBLEM Why can't molecules with the following formulas exist?
$\mathbf{1... |
59 | Electron-Dot and Line-Bond Structures -
Acetate ion
PROBLEM Convert the following line-bond structures into molecular formulas:
1-29
(a)
(b)
(c)
Vitamin C (ascorbic acid)
Nicotine
(d)
Glucose
PROBLEM Convert the following molecular formulas into line-bond structures that are consistent with valence
1-30 rules:
(... |
60 | Electron-Dot and Line-Bond Structures -
PROBLEM Oxaloacetic acid, an important intermediate in food metabolism, has the formula $\mathrm{C}_{4} \mathrm{H}_{4} \mathrm{O}_{5}$ and
1-32 contains three $\mathrm{C}=\mathrm{O}$ bonds and two $\mathrm{O}-\mathrm{H}$ bonds. Propose two possible structures.
PROBLEM Draw struc... |
61 | Electron-Dot and Line-Bond Structures -
(b)
Pyridine
(c)
Lactic acid
(in sour milk)
PROBLEM Propose structures for molecules that meet the following descriptions:
1-38 (a) Contains two $s p^{2}$-hybridized carbons and two $s p^{3}$-hybridized carbons
(b) Contains only four carbons, all of which are $s p^{2}$-hybrid... |
62 | General Problems -
PROBLEM Why do you suppose no one has ever been able to make cyclopentyne as a stable molecule?
1-45
Cyclopentyne
PROBLEM Allene, $\mathrm{H}_{2} \mathrm{C}=\mathrm{C}=\mathrm{CH}_{2}$, has two adjacent double bonds. Draw a picture showing the orbitals involved
1-46 in the $\sigma$ and $\pi$ bond... |
63 | A carbocation -
(a) How many valence electrons does the positively charged carbon atom have?
(b) What hybridization do you expect this carbon atom to have?
(c) What geometry is the carbocation likely to have?
PROBLEM A carbanion is a species that contains a negatively charged, trivalent carbon.
1-50
A carbanion
(a)... |
64 | A carbocation -
Ibuprofen
Naproxen
Acetaminophen
(a) How many $s p^{3}$-hybridized carbons does each molecule have?
(b) How many $s p^{2}$-hybridized carbons does each molecule have?
(c) What similarities can you see in their structures? |
65 | CHAPTER 2 -
Polar Covalent Bonds; Acids and Bases
FIGURE 2.1 The opium poppy is the source of morphine, one of the first "vegetable alkali," or alkaloids, to be isolated. (credit: modification of work "Papaver somniferum" by Liz West/Flickr, CC BY 2.0) |
66 | CHAPTER CONTENTS -
2.1 Polar Covalent Bonds and Electronegativity
2.2 Polar Covalent Bonds and Dipole Moments
2.3 Formal Charges
2.4 Resonance
2.5 Rules for Resonance Forms
2.6 Drawing Resonance Forms
2.7 Acids and Bases: The Brønsted-Lowry Definition
2.8 Acid and Base Strength
2.9 Predicting Acid-Base Reactions from ... |
67 | CHAPTER CONTENTS - 2.1 Polar Covalent Bonds and Electronegativity
Up to this point, we've treated chemical bonds as either ionic or covalent. The bond in sodium chloride, for instance, is ionic. Sodium transfers an electron to chlorine to produce $\mathrm{Na}^{+}$and $\mathrm{Cl}^{-}$ions, which are held together in th... |
68 | CHAPTER CONTENTS - 2.1 Polar Covalent Bonds and Electronegativity
| $\begin{gathered} \mathrm{H} \\ 2.1 \end{gathered}$ | | | | | | | | | | | | | | | | | He |
| :---: | :---: | :---: | :---: | :---: | :---: | :---: | :---: | :---: | :---: | :---: | :---: | :---: | :---: | :---: | :---: | :---: | :---: |... |
69 | CHAPTER CONTENTS - 2.1 Polar Covalent Bonds and Electronegativity
| $\begin{aligned} & \mathrm{Na} \\ & 0.9 \end{aligned}$ | $\begin{aligned} & \hline \mathrm{Mg} \\ & 1.2 \end{aligned}$ | | | | | | | | | | | $\begin{aligned} & \mathrm{Al} \\ & 1.5 \end{aligned}$ | $\begin{gathered} \mathrm{Si} \\ 1.8 \end{ga... |
70 | CHAPTER CONTENTS - 2.1 Polar Covalent Bonds and Electronegativity
| $\begin{gathered} \mathrm{K} \\ 0.8 \end{gathered}$ | $\begin{aligned} & \hline \mathrm{Ca} \\ & 1.0 \end{aligned}$ | $\begin{aligned} & \mathrm{Sc} \\ & 1.3 \end{aligned}$ | Ti 1.5 | $\begin{gathered} \hline \mathrm{V} \\ 1.6 \end{gathered}$ | $\begin... |
71 | CHAPTER CONTENTS - 2.1 Polar Covalent Bonds and Electronegativity
| $\begin{aligned} & \hline \mathrm{Rb} \\ & 0.8 \end{aligned}$ | $\begin{gathered} \hline \mathrm{Sr} \\ 1.0 \end{gathered}$ | Y 1.2 | $\begin{aligned} & \mathrm{Zr} \\ & 1.4 \end{aligned}$ | $\begin{aligned} & \hline \mathrm{Nb} \\ & 1.6 \end{aligned}$... |
72 | CHAPTER CONTENTS - 2.1 Polar Covalent Bonds and Electronegativity
| $\begin{gathered} \hline \mathrm{Cs} \\ 0.7 \end{gathered}$ | $\begin{aligned} & \mathrm{Ba} \\ & 0.9 \end{aligned}$ | La 1.0 | Hf 1.3 | Ta 1.5 | W | Re 1.9 | $\begin{aligned} & \hline \text { Os } \\ & 2.2 \end{aligned}$ | Ir 2.2 | Pt 2.2 | Au 2.4 | H... |
73 | CHAPTER CONTENTS - 2.1 Polar Covalent Bonds and Electronegativity
FIGURE 2.3 Electronegativity values and trends. Electronegativity generally increases from left to right across the periodic table and decreases from top to bottom. The values are on an arbitrary scale, with $\mathrm{F}=4.0$ and $\mathrm{Cs}=0.7$. Elemen... |
74 | CHAPTER CONTENTS - 2.1 Polar Covalent Bonds and Electronegativity
(b)
Methyllithium
FIGURE 2.4 Polar covalent bonds. (a) Methanol, $\mathrm{CH}_{3} \mathrm{OH}$, has a polar covalent $\mathrm{C}-\mathrm{O}$ bond, and (b) methyllithium, $\mathrm{CH}_{3} \mathrm{Li}$, has a polar covalent C-Li bond. The computer-gene... |
75 | CHAPTER CONTENTS - 2.1 Polar Covalent Bonds and Electronegativity
PROBLEM Which element in each of the following pairs is more electronegative?
2-1 (a) Li or H
(b) B or Br
(c) Cl or I (d) C or H
PROBLEM Use the $\delta+/ \delta$ - convention to indicate the direction of expected polarity for each of the bonds 2-2 indi... |
76 | CHAPTER CONTENTS - 2.2 Polar Covalent Bonds and Dipole Moments
Just as individual bonds are often polar, molecules as a whole are often polar as well. Molecular polarity results from the vector summation of all individual bond polarities and lone-pair contributions in the molecule. As a practical matter, strongly polar... |
77 | CHAPTER CONTENTS - 2.2 Polar Covalent Bonds and Dipole Moments
Dipole moments for some common substances are given in TABLE 2.1. Of the compounds shown in the table, sodium chloride has the largest dipole moment $(9.00 \mathrm{D})$ because it is ionic. Even small molecules like water ( $\mu=1.85 \mathrm{D}$ ), methanol... |
78 | CHAPTER CONTENTS - 2.2 Polar Covalent Bonds and Dipole Moments
TABLE 2.1 Dipole Moments of Some Compounds
| Compound | Dipole moment (D) | Compound | Dipole moment (D) |
| :---: | :---: | :---: | :---: |
| $\mathrm{CH}_{3} \mathrm{CO}_{2} \mathrm{H}$ | 1.70 | <smiles>c1ccccc1</smiles> <br> Benzene | 0 |
| $\mathrm{CH}... |
79 | Predicting the Direction of a Dipole Moment -
Make a three-dimensional drawing of methylamine, $\mathrm{CH}_{3} \mathrm{NH}_{2}$, and show the direction of its dipole moment $(\mu=$ 1.31). |
80 | Strategy -
Look for any lone-pair electrons, and identify any atom with an electronegativity substantially different from that of carbon. (Usually, this means O, N, F, Cl, or Br.) Electron density will be displaced in the general direction of the electronegative atoms and the lone pairs. |
81 | Solution -
Methylamine contains an electronegative nitrogen atom with a lone pair of electrons. The dipole moment thus points generally from $-\mathrm{CH}_{3}$ toward the lone pair.
Methylamine
( $\mu=1.31$ )
PROBLEM Ethylene glycol, $\mathrm{HOCH}_{2} \mathrm{CH}_{2} \mathrm{OH}$, may look nonpolar when drawn, but... |
82 | Solution - 2.3 Formal Charges
Closely related to the ideas of bond polarity and dipole moment is the assignment of formal charges to specific atoms within a molecule, particularly atoms that have an apparently "abnormal" number of bonds. Look at dimethyl sulfoxide $\left(\mathrm{CH}_{3} \mathrm{SOCH}_{3}\right)$, for i... |
83 | Dimethyl sulfoxide -
Formal charges, as the name suggests, are a formalism and don't imply the presence of actual ionic charges in a molecule. Instead, they're a device for electron "bookkeeping" and can be thought of in the following way: A typical covalent bond is formed when each atom donates one electron. Although... |
84 | Dimethyl sulfoxide -
$$
\begin{aligned}
\text { Formal charge } & =\left(\begin{array}{c}
\text { Number of } \\
\text { valence electrons } \\
\text { in free atom }
\end{array}\right)-\left(\begin{array}{c}
\text { Number of } \\
\text { valence electrons } \\
\text { in bonded atom }
\end{array}\right) \\
& =\left(... |
85 | Dimethyl sulfoxide -
TABLE 2.2 A Summary of Common Formal Charges
| Atom | C | | | N | | 0 | | S | | P |
| :---: | :---: | :---: | :---: | :---: | :---: | :---: | :---: | :---: | :---: | :---: |
| Structure | <smiles>CC(C)(C)C</smiles> | <smiles>CC(C)(C)C</smiles> | <smiles>CC(C)(C)C</smiles> | | <smiles>CN+C</... |
86 | Dimethyl sulfoxide - 2.4 Resonance
Most substances can be represented unambiguously by the Kekulé line-bond structures we've been using up to this point, but an interesting problem sometimes arises. Look at the acetate ion, for instance. When we draw a line-bond structure for acetate, we need to show a double bond to o... |
87 | Dimethyl sulfoxide - 2.4 Resonance
Resonance is a very useful concept that we'll return to on numerous occasions throughout the rest of this book. We'll see in Chapter 15, for instance, that the six carbon-carbon bonds in aromatic compounds, such as benzene, are equivalent and that benzene is best represented as a hybr... |
88 | Dimethyl sulfoxide - 2.5 Rules for Resonance Forms
When first dealing with resonance forms, it's useful to have a set of guidelines that describe how to draw and interpret them. The following rules should be helpful: |
89 | RULE 1 -
Individual resonance forms are imaginary, not real. The real structure is a composite, or resonance hybrid, of the different forms. Species such as the acetate ion and benzene are no different from any other. They have single, unchanging structures, and they don't switch back and forth between resonance forms... |
90 | RULE 2 -
Resonance forms differ only in the placement of their $\boldsymbol{\pi}$ or nonbonding electrons. Neither the position nor the hybridization of any atom changes from one resonance form to another. In the acetate ion, for instance, the carbon atom is $s p^{2}$-hybridized and the oxygen atoms remain in exactly ... |
91 | RULE 2 -
RULE 5
The resonance hybrid is more stable than any individual resonance form. In other words, resonance leads to stability. Generally speaking, the larger the number of resonance forms a substance has, the more stable the substance is, because its electrons are spread out over a larger part of the molecule a... |
92 | RULE 2 - 2.6 Drawing Resonance Forms
Look back at the resonance forms of the acetate ion and the acetone anion shown in the previous section. The pattern seen there is a common one that leads to a useful technique for drawing resonance forms. In general, any three-atom grouping with a $p$ orbital on each atom has two r... |
93 | Drawing Resonance Forms for an Anion -
Draw three resonance structures for the carbonate ion, $\mathrm{CO}_{3}{ }^{2-}$. |
94 | Strategy -
Look for three-atom groupings that contain a multiple bond next to an atom with a $p$ orbital. Then exchange the positions of the multiple bond and the electrons in the $p$ orbital. In the carbonate ion, each singly bonded oxygen atom with three lone pairs and a negative charge is adjacent to the $\mathrm{C... |
95 | Solution -
Exchanging the position of the double bond and an electron lone pair in each grouping generates three resonance structures. |
96 | Drawing Resonance Forms for a Radical -
Draw three resonance forms for the pentadienyl radical, where a radical is a substance that contains a single, unpaired electron in one of its orbitals, denoted by a dot $(\cdot)$. |
97 | Strategy -
Find the three-atom groupings that contain a multiple bond next to an atom with a $p$ orbital. |
98 | Solution -
The unpaired electron is on a carbon atom next to a $\mathrm{C}=\mathrm{C}$ bond, giving a typical three-atom grouping that has two resonance forms.
In the second resonance form, the unpaired electron is next to another double bond, giving another three-atom grouping and leading to another resonance form.... |
99 | Solution - 2.7 Acids and Bases: The Brønsted-Lowry Definition
Perhaps the most important of all concepts related to electronegativity and polarity is that of acidity and
basicity. We'll soon see, in fact, that the acid-base behavior of organic molecules explains much of their chemistry. You may recall from a course in ... |
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