Carey - Organic Chemistry - chapt01

Carey - Organic Chemistry - chapt01

(Parte 1 de 7)


Structure*is the key to everything in chemistry. The properties of a substance depend on the atoms it contains and the way the atoms are connected. What is less obvious, but very powerful, is the idea that someone who is trained in chemistry can look at a structural formula of a substance and tell you a lot about its properties. This chapter begins your training toward understanding the relationship between structure and properties in organic compounds. It reviews some fundamental principles of molecular structure and chemical bonding.By applying these principles you will learn to recognize the structural patterns that are more stable than others and develop skills in communicating chemical information by way of structural formulas that will be used throughout your study of organic chemistry.


Before discussing bonding principles, let’s first review some fundamental relationships between atoms and electrons. Each element is characterized by a unique atomic number Z,which is equal to the number of protons in its nucleus. Aneutral atom has equal numbers of protons, which are positively charged, and electrons, which are negatively charged.

Electrons were believed to be particles from the time of their discovery in 1897 until 1924, when the French physicist Louis de Broglie suggested that they have wavelike properties as well. Two years later Erwin Schrödinger took the next step and calculated the energy of an electron in a hydrogen atom by using equations that treated the electron as if it were a wave. Instead of a single energy, Schrödinger obtained a series of energy levels,each of which corresponded to a different mathematical description of the electron wave. These mathematical descriptions are called wave functionsand are symbolized by the Greek letter (psi).

*A glossary of important terms may be found immediately before the index at the back of the book.

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According to the Heisenberg uncertainty principle, we can’t tell exactly where an electron is, but we can tell where it is most likely to be. The probability of finding an electron at a particular spot relative to an atom’s nucleus is given by the square of the wave function ( 2) at that point. Figure 1.1 illustrates the probability of finding an electron at various points in the lowest energy (most stable) state of a hydrogen atom. The darker the color in a region, the higher the probability. The probability of finding an electron at a particular point is greatest near the nucleus, and decreases with increasing distance from the nucleus but never becomes zero. We commonly describe Figure 1.1 as an “electron cloud” to call attention to the spread-out nature of the electron probability. Be careful, though. The “electron cloud” of a hydrogen atom, although drawn as a collection of many dots, represents only one electron.

Wave functions are also called orbitals.For convenience, chemists use the term “orbital” in several different ways. Adrawing such as Figure 1.1 is often said to represent an orbital. We will see other kinds of drawings in this chapter, use the word “orbital” to describe them too, and accept some imprecision in language as the price to be paid for simplicity of expression.

Orbitals are described by specifying their size, shape, and directional properties.

Spherically symmetrical ones such as shown in Figure 1.1 are called s orbitals.The letter sis preceded by the principal quantum numbern (n 1, 2, 3, etc.) which specifies the shelland is related to the energy of the orbital. An electron in a 1sorbital is likely to be found closer to the nucleus, is lower in energy, and is more strongly held than an electron in a 2sorbital.

Regions of a single orbital may be separated by nodal surfaceswhere the probability of finding an electron is zero. A1sorbital has no nodes; a 2sorbital has one. A 1sand a 2sorbital are shown in cross section in Figure 1.2. The 2swave function changes sign on passing through the nodal surface as indicated by the plus ( ) and minus ( ) signs in Figure 1.2. Do not confuse these signs with electric charges—they have nothing to do with electron or nuclear charge.Also, be aware that our “orbital” drawings are really representations of 2(which must be a positive number), whereas and refer to the sign of the wave function ( ) itself. These customs may seem confusing at first but turn out not to complicate things in practice. Indeed, most of the time we won’t

8CHAPTER ONEChemical Bonding x z

FIGURE 1.1Probability distribution ( 2) for an electron in a 1sorbital.


Nucleus x

FIGURE 1.2Cross sections of (a) a 1sorbital and (b) a 2sorbital. The wave function has the same sign over the entire 1sorbital. It is arbitrarily shown as , but could just as well have been designated as . The 2sorbital has a spherical node where the wave function changes sign.

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Instead of probability distributions, it is more common to represent orbitals by their boundary surfaces,as shown in Figure 1.3 for the 1sand 2sorbitals. The boundary surface encloses the region where the probability of finding an electron is high—on the order of 90–95%. Like the probability distribution plot from which it is derived, a picture of a boundary surface is usually described as a drawing of an orbital.

Ahydrogen atom (Z 1) has one electron; a helium atom (Z 2) has two. The single electron of hydrogen occupies a 1sorbital, as do the two electrons of helium. The respective electron configurations are described as:

In addition to being negatively charged, electrons possess the property of spin.The spin quantum numberof an electron can have a value of either 12or 12. According to the Pauli exclusion principle,two electrons may occupy the same orbital only when they have opposite, or “paired,” spins. For this reason, no orbital can contain more than two electrons. Since two electrons fill the 1sorbital, the third electron in lithium (Z 3) must occupy an orbital of higher energy. After 1s,the next higher energy orbital is 2s.The third electron in lithium therefore occupies the 2sorbital, and the electron configuration of lithium is

The period(or row) of the periodic table in which an element appears corresponds to the principal quantum number of the highest numbered occupied orbital (n 1 in the case of hydrogen and helium). Hydrogen and helium are first-row elements; lithium (n 2) is a second-row element. With beryllium (Z 4), the 2slevel becomes filled, and the next orbitals to be occupied in it and the remaining second-row elements are the 2px,2py,and 2pzorbitals. These orbitals, portrayed in Figure 1.4, have a boundary surface that is usually described as “dumbbell-shaped.” Each orbital consists of two “lobes,” that is, slightly flattened spheres that touch each other along a nodal plane passing through the nucleus. The 2px,

2py,and 2pzorbitals are equal in energy and mutually perpendicular. The electron configurations of the first 12 elements, hydrogen through magnesium, are given in Table 1.1. In filling the 2porbitals, notice that each is singly occupied before any one is doubly occupied. This is a general principle for orbitals of equal energy known y y

FIGURE 1.3Boundary surfaces of a 1sorbital and a 2sorbital. The boundary surfaces enclose the volume where there is a 90–95% probability of finding an electron.

A complete periodic table of the elements is presented on the inside back cover.

BackForwardMain MenuTOCStudy Guide TOCStudent OLCMHHE Website as Hund’s rule.Of particular importance in Table 1.1 are hydrogen, carbon, nitrogen, and oxygen.Countless organic compounds contain nitrogen, oxygen, or both in addition to carbon, the essential element of organic chemistry. Most of them also contain hydrogen.

It is often convenient to speak of the valence electronsof an atom. These are the outermost electrons, the ones most likely to be involved in chemical bonding and reactions. For second-row elements these are the 2sand 2pelectrons. Because four orbitals

(2s,2px,2py,2pz) are involved, the maximum number of electrons in the valence shell of any second-row element is 8. Neon, with all its 2sand 2porbitals doubly occupied, has eight valence electrons and completes the second row of the periodic table.

PROBLEM 1.1How many valence electrons does carbon have?

Once the 2sand 2porbitals are filled, the next level is the 3s,followed by the 3px,3py, and 3pzorbitals. Electrons in these orbitals are farther from the nucleus than those in the 2sand 2porbitals and are of higher energy.

10CHAPTER ONEChemical Bonding x x z y y z

FIGURE 1.4Boundary surfaces of the 2porbitals. The wave function changes sign at the nucleus. The yz-plane is a nodal surface for the 2pxorbital. The probability of finding a 2px electron in the yz-plane is zero. Analogously, the xz-plane is a nodal surface for the 2pyorbital, and the xy-plane is a nodal surface for the 2pzorbital.

TABLE 1.1Electron Configurations of the First Twelve Elements of the Periodic Table

Number of electrons in indicated orbital


Hydrogen Helium Lithium Beryllium Boron Carbon Nitrogen Oxygen Fluorine Neon Sodium Magnesium

Atomic number Z

Answers to all problems that appear within the body of a chapter are found in Appendix 2. A brief discussion of the problem and advice on how to do problems of the same type are offered in the Study Guide.

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PROBLEM 1.2Referring to the periodic table as needed, write electron configurations for all the elements in the third period.

SAMPLE SOLUTIONThe third period begins with sodium and ends with argon. The atomic number Zof sodium is 1, and so a sodium atom has 1 electrons. The maximum number of electrons in the 1s,2s,and 2porbitals is ten, and so the eleventh electron of sodium occupies a 3sorbital. The electron configuration of

Neon, in the second period, and argon, in the third, possess eight electrons in their valence shell; they are said to have a complete octetof electrons. Helium, neon, and argon belong to the class of elements known as noble gasesor rare gases.The noble gases are characterized by an extremely stable “closed-shell” electron configuration and are very unreactive.

Atoms combine with one another to give compoundshaving properties different from the atoms they contain. The attractive force between atoms in a compound is a chemical bond.One type of chemical bond, called an ionic bond,is the force of attraction between oppositely charged species (ions) (Figure 1.5). Ions that are positively charged are referred to as cations;those that are negatively charged are anions.

Whether an element is the source of the cation or anion in an ionic bond depends on several factors, for which the periodic table can serve as a guide. In forming ionic compounds, elements at the left of the periodic table typically lose electrons, forming a cation that has the same electron configuration as the nearest noble gas. Loss of an electron from sodium, for example, gives the species Na , which has the same electron configuration as neon.

Alarge amount of energy, called the ionization energy,must be added to any atom in order to dislodge one of its electrons. The ionization energy of sodium, for example, is 496 kJ/mol (119 kcal/mol). Processes that absorb energy are said to be endothermic. Compared with other elements, sodium and its relatives in group IAhave relatively low ionization energies. In general, ionization energy increases across a row in the periodic table.

Elements at the right of the periodic table tend to gain electrons to reach the electron configuration of the next higher noble gas. Adding an electron to chlorine, for example, gives the anion Cl , which has the same closed-shell electron configuration as the noble gas argon.

Energy is released when a chlorine atom captures an electron. Energy-releasing reactions are described as exothermic,and the energy change for an exothermic process has a negative sign. The energy change for addition of an electron to an atom is referred to as its electron affinityand is 349 kJ/mol ( 83.4 kcal/mol) for chlorine.

FIGURE 1.5An ionic bond is the force of electrostatic attraction between oppositely charged ions, illustrated in this case by Na (red) and Cl (green). In solid sodium chloride, each sodium ion is surrounded by six chloride ions and vice versa in a crystal lattice.

In-chapter problems that contain multiple parts are accompanied by a sample solution to part (a). Answers to the other parts of the problem are found in Appendix 2, and detailed solutions are presented in the Study Guide.

The SI (Système International d’Unites)unit of energy is thejoule(J). An older unit is the calorie(cal). Most organic chemists still express energy changes in units of kilocalories per mole (1 kcal/mol 4.184 kJ/mol).

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PROBLEM 1.3Which of the following ions possess a noble gas electron configuration?

Transfer of an electron from a sodium atom to a chlorine atom yields a sodium cation and a chloride anion, both of which have a noble gas electron configuration:

Were we to simply add the ionization energy of sodium (496 kJ/mol) and the electron affinity of chlorine ( 349 kJ/mol), we would conclude that the overall process is endothermic with H° 147 kJ/mol. The energy liberated by adding an electron to chlorine is insufficient to override the energy required to remove an electron from sodium. This analysis, however, fails to consider the force of attraction between the oppositely charged ions Na and Cl–, which exceeds 500 kJ/mol and is more than sufficient to make the overall process exothermic. Attractive forces between oppositely charged particles are termed electrostatic,or coulombic, attractionsand are what we mean by an ionic bondbetween two atoms.

PROBLEM 1.4What is the electron configuration of C ? Of C ? Does either one of these ions have a noble gas (closed-shell) electron configuration?

Ionic bonds are very common in inorganiccompounds, but rare in organicones.

The ionization energy of carbon is too large and the electron affinity too small for carbon to realistically form a C4 or C4 ion. What kinds of bonds, then, link carbon to other elements in millions of organic compounds? Instead of losing or gaining electrons, carbon shareselectrons with other elements (including other carbon atoms) to give what are called covalent bonds.


The covalent,or shared electron pair,model of chemical bonding was first suggested by G. N. Lewis of the University of California in 1916. Lewis proposed that a sharing of two electrons by two hydrogen atoms permits each one to have a stable closed-shell electron configuration analogous to helium.

Two hydrogen atoms, each with a single electron

Hydrogen molecule: covalent bonding by way of a shared electron pair

(Parte 1 de 7)