Organic chemistry paula bruice 8th edition pdf free download

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As all download links are working fine from our end. When the single valence electron of lithium or sodium is removed, the species that is formed is called an ion because it carries a charge. Fluorine has seven valence electrons Table 1. Therefore, it readily acquires an electron to fill its outer shell.

Gaining the electron forms F - , a fluoride ion. A hydrogen atom has one valence electron. Therefore, it can achieve a completely empty shell by losing an electron, or a filled outer shell by gaining an electron.

A positively charged hydrogen ion is called a proton because when a hydrogen atom loses its valence electron, only the hydrogen nucleus—which consists of a single proton—remains.

When a hydrogen atom gains an electron, a negatively charged hydrogen ion—called a hydride ion—is formed. Find potassium K in the periodic table and predict how many valence electrons it has. What orbital does the unpaired electron occupy? For example, two fluorine atoms can each attain a filled second shell by sharing their unpaired valence electrons. A bond formed as a result of sharing electrons between two nuclei is called a covalent bond.

A covalent bond is commonly shown by a solid line rather than by a pair of dots. Oxygen has six valence elec- trons, so it needs to form two covalent bonds to achieve an outer shell of eight electrons. Nitrogen, with five valence electrons, must form three covalent bonds, and carbon, with four valence elec- trons, must form four covalent bonds.

Notice that all the atoms in water, ammonia, and methane have filled outer shells. Electronegativity is a measure of the ability of an atom to the same electronegativity. The electronegativities of some of the elements are shown in Table 1. A polar covalent bond is Notice that electronegativity increases from left to right across a row of the periodic table and from a covalent bond between atoms bottom to top in any of the columns.

As a result, there are several scales of electronegativities. The electronegativities listed here are from the scale devised by Linus Pauling. If the electronegativity difference between the bonded atoms is less than 0. That is, the atoms share the bonding electrons equally—the electrons represented by the bond are symmetrically distributed around each atom. Examples of nonpolar covalent bonds are shown below. The negative end of the sodium chloride crystals table salt bond is the end with the more electronegative atom.

The greater the difference in electronegativity between the bonded atoms, the more polar the bond. By convention, chemists draw the www. Thus, the head of the arrow is at the negative end of the bond; a short perpendicular line near the tail of the arrow marks the positive end of the bond.

Physicists draw the arrow in the opposite direction. Lithium If the electronegativity difference between the atoms is greater than 1. Sodium chloride is an example of an ionic compound also known as a salt.

Ionic disorder. Scientists do not yet know why compounds are formed when an element on the left side of the periodic table transfers one or more lithium salts have these therapeutic effects. Polar covalent bonds fall somewhere in between. The size of the dipole is indicated by the dipole moment. Because the charge on an electron is 4.

Thus, a dipole moment of 1. The dipole moments of bonds commonly found in organic compounds are listed in Table 1. For example, the dipole moment of hydrogen chloride HCl is 1. The dipole moment of a molecule with more than one covalent bond depends on the dipole moments of all the bonds in the molecule and the geometry of the molecule. We will look at this in Section 1. Because oxygen is more electronegative than carbon, oxygen has a partial negative charge and carbon has a partial positive charge.

H 3Cd. The bond length is 1. The charge on an electron is 4. Red, signifying the most negative electrostatic potential, is used for regions that attract electron-deficient species most strongly. Because the angstrom continues to be used by many organic chemists, we will use angstroms in this text book. Other colors indicate interme- diate levels of attraction. By comparing the three maps, we can tell that the hydrogen in LiH is more electron- rich than a hydrogen in H2, whereas the hydrogen in HF is less electron-rich than a hydrogen in H2.

A particular atom can have different sizes in different molecules, because the size of an atom in a potential map depends on its electron density. For example, the negatively charged hydrogen in LiH is bigger than a neutral hydrogen in H2, which is bigger than the positively charged hydrogen in HF.

Which compounds are polar? Why does LiH have the largest hydrogen? Which compound has the hydrogen that would be most apt to attract a negatively charged molecule? Then we will look at the kinds of structures that are used more commonly for organic compounds. Lewis Structures The chemical symbols we have been using, in which the valence electrons are represented as dots or solid lines, are called Lewis structures.

Lewis structures show us which atoms are bonded together and tell us whether any atoms possess lone-pair electrons or have a formal charge, two concepts described below. Therefore, www. Lone-Pair Electrons When you draw a Lewis structure, make sure hydrogen atoms are surrounded by two electrons and Lone-pair electrons are valence C, O, N, and halogen F, Cl, Br, I atoms are surrounded by eight electrons, in accordance with electrons that do not form bonds.

Valence electrons not used in bonding are called nonbonding electrons, lone-pair electrons, or simply, lone pairs. Formal Charge Once the atoms and the electrons are in place, you must examine each atom to see whether a formal charge should be assigned to it. Notice that half the bonding electrons is the same as the number of bonds. An oxygen atom has six valence electrons Table 1. Notice that one-half of the bonding electrons is the same as the number of bonds.

Which atom bears the formal negative charge in the hydroxide ion? Which atom has the greater electron density in the hydroxide ion? Which atom bears the formal positive charge in the hydronium ion? Which atom has the least electron density in the hydronium ion? Drawing Lewis Structures Nitrogen has five valence electrons Table 1.

Recall that a cation is a positively charged ion and an anion is a negatively charged ion. A species containing an atom with a single unpaired A carbanion is a species electron is called a radical often called a free radical. CH3 O CH3 b. These numbers are very important to remember when you are drawing structures of organic compounds because they provide a quick way to recognize when you have made a mistake.

Atoms with more bonds or fewer bonds than is required for a neutral atom must have either a formal charge or an unpaired electron. Each atom in the following Lewis structures has a filled outer shell.

Notice that because none of the molecules has a formal charge or an unpaired electron, H forms 1 bond, C forms 4 bonds, N forms 3 bonds, O forms 2 bonds, and Br forms 1 bond. Notice, too, that each N has 1 lone pair, each O has 2 lone pairs, and Br has 3 lone pairs. Draw the Lewis structure for CH4O. Draw the Lewis structure for HNO2. Distribute the atoms, remembering that C forms 4 bonds, O forms 2 bonds, and each H forms 1 bond. Always put the hydrogens on the outside of the molecule because H can form only 1 bond.

F orm bonds and fill octets with lone-pair electrons, using the number of valence electrons determined in 1. A ssign a formal charge to any atom whose number of valence electrons is not equal to the number of its lone-pair electrons plus the number of bonds.

None of the atoms in CH4O has a formal charge. Distribute the atoms, putting the hydrogen on the outside of the molecule. If a species has two or more oxygen atoms, avoid oxygen—oxygen single bonds.

These are weak bonds, and few compounds have them. Form bonds and fill octets with lone-pair electrons, using the number of valence electrons determined in 1. None of the atoms in HNO2 has a formal charge. NO3- c. CH3 NH3 g. HCO3- www. NaOH h. The total number of valence electrons is 23 5 for N and 6 for each of the three Os. Because the species has one negative charge, we must add 1 to the number of valence electrons, for a total of We then use the 24 electrons to form bonds and fill octets with lone-pair electrons.

O O N O incomplete octet. It does not make a difference which oxygen atom we choose. When we check each atom to see whether it has a formal charge, we find that two of the Os are negatively charged and that the N is positively charged, for an overall charge of - 1. The total number of valence electrons is 17 5 for N and 6 for each of the two Os.

Because the species has one positive charge, we must subtract 1 from the number of valence electrons, for a total of The 16 electrons are used to form bonds and fill octets with lone-pair electrons. Draw two Lewis structures for C2H6O. Draw three Lewis structures for C3H8O. Hint: The two Lewis structures in part a are constitutional isomers—molecules that have the same atoms but differ in the way the atoms are connected.

The three Lewis structures in part b are also constitutional isomers. Lone pairs are usually not shown, unless they are needed to draw attention to some chemical property of the molecule. Notice that because none of the molecules in Table 1. CH3CH2Cl b. CH3OCH3 f. CH3 2CHCl c. CH3 3CBr d. Each vertex in a skeletal structure represents a carbon, and each carbon is understood to be bonded to the appropriate number of hydrogens to give the carbon four bonds. Atoms other than carbons are shown, and hydrogens bonded to atoms other than carbon are also shown.

An atomic orbital is a three-dimensional region around the nucleus where an electron is most likely to be found. Because the Heisenberg uncertainty principle states that both the precise location and the exact momen- tum of an atomic particle cannot be simultaneously determined, we can never say precisely where an electron is—we can only describe its probable location.

Because the second shell lies farther from the nucleus than does the first shell Section 1. A 2s orbital, therefore, is represented by a larger sphere. Because of the greater size of a 2s orbital, its average electron density is less than the average electron density of a 1s orbital.

This is called a node. Nodes occur because electrons have both particle-like and wave-like properties. A node is a consequence of the wave-like properties of an electron.

There are two types of waves: traveling waves and standing waves. Traveling waves move through space. Light is an example of a traveling wave. A standing wave, on the other hand, is confined to a limited space.

The vibrating string of a guitar is a standing wave—the string moves up and down but does not travel through space. The region where the guitar string has no transverse displacement zero amplitude is a node. This means that the node of a 2s orbital is actually a spherical surface within the 2s orbital.

Because the electron wave has zero amplitude at the node, there is zero probability of finding an electron at the node. Generally, the lobes are depicted as teardrop shaped, but computer-generated representations reveal that they are shaped more like doorknobs as shown on the right on the top of p. The node of the p orbital is a plane—called a nodal plane—that passes through the center of the nucleus, between its two lobes.

There is zero probability of finding an electron in the nodal plane of the p orbital. This means that each p orbital is perpendicular to the other two p orbitals. The energy of a 2p orbital is slightly greater than that of a 2s orbital because the average location of an electron in a 2p orbital is farther away from the nucleus. The Lewis model, which shows www. A drawback of the model is that it treats electrons like particles and does not take into account their wave-like properties.

Molecular orbital MO theory combines the tendency of atoms to fill their octets by sharing electrons the Lewis model with their wave-like properties, assigning electrons to a volume of space called an orbital.

According to MO theory, covalent bonds result when atomic orbitals combine to form molecular orbitals. And like atomic orbitals, molecular orbitals, too, have specific sizes, shapes, and energies. An atomic orbital surrounds an atom.

A molecular orbital surrounds a molecule. Imagine a meeting of two separate H atoms. As the 1s atomic orbital of one hydrogen atom approaches the 1s atomic orbital of the other hydrogen atom, the orbitals begin to overlap.

The atoms continue to move closer, and the amount of overlap increases until the orbitals combine to form a molecular orbital. The covalent bond that is formed when the two s orbitals overlap is called a sigma S bond.

A S bond is cylin- drically symmetrical—the electrons in the bond are symmetrically distributed about an imaginary line connecting the nuclei of the two atoms joined by the bond. The more the orbitals overlap, the more the energy decreases, until the atoms are so close that their positively charged nuclei begin to repel each other.

This repulsion causes a large increase in energy. Figure 1. This distance is the bond length of the new to maximum stability. As Figure 1. Breaking the bond requires precisely the www. Thus, the bond dissociation energy—a measure of bond strength—is the energy required to break a bond or the energy released when a bond is formed.

Every covalent bond has a characteristic bond length and bond dissociation energy. Bonding and Antibonding Molecular Orbitals Orbitals are conserved. In other words, the number of molecular orbitals formed must equal the number of atomic orbitals combined.

Where is the other molecular orbital? It is the wave-like properties of the electrons that cause two atomic orbitals to form two molecu- lar orbitals. The two atomic orbitals can combine in an additive constructive manner, just as two light waves or two sound waves can reinforce each other Figure 1.

The constructive combina- tion of two s atomic orbitals is called a s sigma bonding molecular orbital. The two atomic orbitals can also combine in a destructive way, canceling each other. The cancel- lation is similar to the darkness that results when two light waves cancel each other or to the silence that results when two sound waves cancel each other Figure 1. In an MO diagram, the energies of both the atomic orbitals and the molecular orbitals are represented as horizontal lines, with the bottom line being the lowest energy level and the top line the highest energy level.

When two atomic orbitals overlap, two molecular orbitals are formed— one lower in energy and one higher in energy than the atomic orbitals. Before covalent bond formation, each atomic orbitals electron is in an atomic orbital. After covalent s bonding molecular orbital bond formation, both electrons are in the bonding MO.

The antibonding MO is empty. This increased electron density between the nuclei is what binds the atoms together Figure 1. Electrons in the antibond- ing molecular orbital, however, are most likely to be found anywhere except between the nuclei, because a nodal plane lies between the nuclei Figure 1. As a result, electrons in the antibonding orbital leave the positively charged nuclei more exposed to one another.

Therefore, electrons in the antibonding orbital detract from, rather than assist in, the formation of a bond. Electrons in a bonding MO assist in bonding. Electrons in an antibonding MO detract from bonding. The MO diagram shows that the bonding molecular orbital is lower in energy and is, therefore, more stable than the individual atomic orbitals.

This is because the more nuclei an electron senses, the more stable it is. The antibonding molecular orbital, with less electron density between the nuclei, is less stable—and, therefore, higher in energy—than the atomic orbitals. It is this electrostatic attraction that gives a covalent bond its strength. We can conclude, therefore, that the strength of the covalent bond increases as the overlap of the atomic orbitals increases. Covalent bond strength increases as atomic orbital overlap increases.

The MO diagram in Figure 1. Using the same diagram, we can also predict that He2 does not exist, because the four electrons of He2 two from each He atom would fill the lower energy bonding MO and the higher energy antibonding MO. The two electrons in the antibonding MO would cancel the advantage to bonding that is gained by the two electrons in the bonding MO. Forming a pi p Bond When two p atomic orbitals overlap, the side of one orbital overlaps the side of the other.

The www. In other Study Guide and Solutions Manual. Thus, a Lewis structure gives us a first approximation of the structure of a simple mol- ecule, and VSEPR gives us a first glance at the shape of the molecule. All other covalent bonds in organic molecules are S bonds.

Because organic chemists generally think of chemical reactions in terms of the changes that occur in the bonds of the reacting molecules, the VSEPR model often provides the easiest way to visualize chemical change. However, the model is inadequate for some molecules because it does not allow for antibonding molecular orbitals. Our choice will depend on which model provides the best description of the molecule under discussion.

Then we will examine the bonding in ethane, a compound with two carbons attached by a carbon—carbon single bond. Because all four bonds have the same length 1. Four different ways to represent a methane molecule are shown here. The potential map of methane shows that neither carbon nor hydrogen carries much of a charge: there are neither red areas, representing partially negatively charged atoms, nor blue areas, repre- senting partially positively charged atoms.

Compare this map with the potential map for water on p. The absence of partially charged atoms can be explained by the similar electronegativities of carbon and hydrogen, which cause them to share their bonding electrons relatively equally see p. Methane, therefore, is a nonpolar molecule.

You might be wondering how carbon can form four covalent bonds when it has only two unpaired valence electrons Table 1. Carbon has to form four covalent bonds otherwise it would not com- plete its octet. We need, therefore, to come up with an explanation that accounts for the fact that The blue colors of Uranus and Neptune carbon forms four covalent bonds when it has only two valence electrons.

How can they be identi- cal if carbon uses an s orbital and three p orbitals to form these four bonds? Hybrid Orbitals Hybrid orbitals result from Hybrid orbitals are mixed orbitals that result from combining atomic orbitals. The concept of www. If the one s and three p orbitals of the second shell are all combined and then apportioned into four equal orbitals, each of the four resulting orbitals will be one part s and three parts p. The four sp3 orbitals are degenerate—that is, they all have the same energy.

The lobes differ in size, however, because the s orbital adds to one lobe of the p orbital and subtracts from the other lobe Figure 1. The result is a hybrid orbital to form a hybrid sp3 orbital with two lobes that differ in size.

The larger lobe of the sp3 orbital is used to form covalent bonds. The stability of an sp3 orbital reflects its composition; it is more stable than a p orbital, but not as stable as an s orbital Figure 1. To simplify the orbital dipictions that follow, the phases of the orbitals will not be shown. An sp3 orbital is more stable lower in energy than a p orbital but less stable higher in energy than an s s orbital. Tetrahedral Carbon; Tetrahedral Bond Angle The four sp3 orbitals adopt a spatial arrangement that keeps them as far away from each other as Electron pairs stay as far possible.

They do this because electrons repel each other, and moving as far from each other as pos- from each other as possible. When four sp3 orbitals move as far from each other as possible, they point toward the corners of a regular tetrahedron—a pyramid with four faces, each an equilateral triangle Figure 1.

This H H arrangement allows the four orbitals to H be as far apart as possible. Therefore, the bond angles in methane are A carbon, such as the one in methane, that forms covalent bonds using four equivalent sp3 orbitals is called a tetrahedral carbon.

If you are thinking that hybrid orbital theory appears to have been contrived just to make things fit, then you are right. Nevertheless, it gives us a very good picture of the bonding in organic compounds. All the bonds in ethane are single bonds. To bond to four atoms, each carbon uses four sp3 orbitals as they do in methane. The smaller lobes of the sp3 orbitals are not shown.

Because both carbons are tetrahedral, each of the bond angles in ethane is nearly the tetrahedral bond angle of The potential map shows that ethane, like methane, is a nonpolar molecule.

H End-on overlap forms a cylindrically symmetrical bond—a sigma s bond Section 1. Thus, all the bonds in methane and ethane are sigma s bonds. Notice in Figure 1. This causes the back lobes the nonoverlapping green lobes to be quite small. PROBLEM 28 Explain why a s bond formed by overlap of an s orbital with an sp3 orbital of carbon is stronger than a s bond formed by overlap of an s orbital with a p orbital of carbon.

Two bonds connect- ing two atoms is called a double bond. Each of the carbons forms four bonds, but each carbon is bonded to only three atoms. H H C C H H ethene ethylene To bond to three atoms, each carbon hybridizes three atomic orbitals: an s orbital and two of the p orbitals. Because three orbitals are hybridized, three hybrid orbitals are formed.

These are called sp2 orbitals. After hybridization, each carbon atom has three degenerate sp2 orbitals and one unhybridized p orbital: an unhybridized p orbital p p p hybridization p sp2 sp2 sp2 s 3 hybrid orbitals are formed 3 orbitals are hybridized www. Therefore, the axes of the three orbitals lie in a plane, directed toward the corners of an equilateral triangle with the carbon nucleus at the center.

The smaller lobes of the sp2 orbitals are not shown. The unhybridized p orbital is perpendicular to the plane defined A double bond consists of one by the axes of the sp2 orbitals Figure 1. S bond and one P bond.

The two bonds in the double bond are not identical. One bond results from the overlap of an sp2 orbital of one carbon with an sp2 orbital of the other carbon; this is a sigma s bond because it is cylindrically symmetrical Figure 1. The second carbon—carbon bond results from side-to-side overlap of the two unhybridized p orbitals. Side-to-side overlap of p orbitals forms a pi p bond Figure 1. Thus, one of the bonds in a double bond is a s bond, and the other is a p bond. Remember that all single bonds in organic compounds are s bonds.

The two p orbitals are parallel to each other. This forces the triangle formed by one carbon and two hydrogens to lie in the same plane as the triangle formed by the other carbon and two hydrogens.

As a result, all six p bond s bond atoms of ethene lie in the same plane, and the electrons in the p orbitals occupy a volume of space above and below the plane Figure 1. Perpendicular negative charge the pale orange area above the two carbons.

If you could turn the potential map to that plane are the two parallel over, you would find a similar accumulation of negative charge on the other side.

This results in an accumulation of electron density above and below the plane containing Representations of Ethene the two carbons and four hydrogens. Diamond is the hardest of all substances, whereas graphite is a slippery, soft solid most familiar to us as the lead in pencils.

Both materials, in spite of their very different physical properties, contain only carbon atoms. The two substances differ solely in the hybridization of the carbon atoms. Diamond consists of a rigid three-dimensional network of carbon atoms, with each carbon bonded to four others via sp3 orbitals.

The carbon atoms in graphite, on the other hand, are sp2 hybridized, so each bonds to only three other carbons.