1.1.2 Chemical bonds

Unit 1 models can be approximated by the model viewing applet Jmol. This powerful visualization tool allows the user to move a molecular structure in any way desired. Atom distances and angles are easily determined. To measure a distance, double-click on two atoms. To measure a bond angle, do a double-click, single-click, double-click on three atoms. To measure a torsion angle, do a double-click, single-click, single-click, doubleclick on four atoms. A pop-up menu of commands may be accessed by the right button on a PC or a controlclick on a Mac while the cursor is inside the display frame. One way in which the shapes of molecules manifest themselves experimentally is through molecular dipole moments. A molecule which has one or more polar covalent bonds may have a dipole moment as a result of the accumulated bond dipoles. In the case of water, we know that the O-H covalent bond is polar, due to the different electronegativities of hydrogen and oxygen. Since there are two O-H bonds in water, their bond dipoles will interact and may result in a molecular dipole which can be measured. The following diagram shows four possible orientations of the O-H bonds. In the linear configuration (bond angle 180º) the bond dipoles cancel, and the molecular dipole is zero. For other bond angles (120 to 90º) the molecular dipole would vary in size, being largest for the 90º configuration. In a similar manner the configurations of methane (CH4) and carbon dioxide (CO2) may be deduced from their zero molecular dipole moments. Since the bond dipoles have cancelled, the configurations of these molecules must be tetrahedral (or squareplanar) and linear respectively. The case of methane provides insight to other arguments that have been used to confirm its tetrahedral configuration. For purposes of discussion we shall consider three other configurations for CH4, square-planar, square-pyramidal and triangular-pyramidal. Substitution of one hydrogen by a chlorine atom gives a CH3Cl compound. Since the tetrahedral, squareplanar and square-pyramidal configurations have structurally equivalent hydrogen atoms, they would each give a single substitution product. However, in the trigonal-pyramidal configuration one hydrogen (the apex) is structurally different from the other three (the pyramid base). Substitution in this case should give two different CH3Cl compounds if all the hydrogen react. In the case of di-substitution, the tetrahedral configuration of methane would lead to a single CH2Cl2 product, but the other configurations would give two different CH2Cl2 compounds. These substitution possibilities are shown in the above insert.

1.1.2 Chemical bonds

A chemical bond is an attraction between atoms. This attraction may be seen as the result of different behaviours of the outermost electrons of atoms. Although all of these behaviours merge into each other seamlessly in various bonding situations so that there is no clear line to be drawn between them, nevertheless behaviours of atoms become so qualitatively different as the character of the bond changes quantitatively, that it remains useful and customary to differentiate between the bonds that cause these different properties of condensed matter.

Table 1.2 Structure of different molecules according to VSEPR theory Figure 1.6 The four possible orientations of the O-H bonds

Unit 1  'Covalent' bond, one or more electrons (often a pair of electrons) are drawn into the space between the two atomic nuclei.  Here the negatively charged electrons are attracted to the positive charges of both nuclei, instead of just their own. This overcomes the repulsion between the two positively charged nuclei of the two atoms, and so this overwhelming attraction holds the two nuclei in a fixed configuration of equilibrium, even though they will still vibrate at equilibrium position. Thus, covalent bonding involves sharing of electrons in which the positively charged nuclei of two or more atoms simultaneously attract the negatively charged electrons that are being shared between them.  These bonds exist between two particular identifiable atoms, and have a direction in space, allowing them to be shown as single connecting lines between atoms in drawings, or modelled as sticks between spheres in models. In a polar covalent bond, one or more electrons are unequally shared between two nuclei.  Covalent bonds often result in the formation of small collections of better-connected atoms called molecules, which in solids and liquids are bound to other molecules by forces that are often much weaker than the covalent bonds that hold the molecules internally together. Such weak intermolecular bonds give organic molecular substances, such as waxes and oils, their soft bulk character, and their low melting points (in liquids, molecules must cease most structured or oriented contact with each other). When covalent bonds link long chains of atoms in large molecules, however (as in polymers such as nylon), or when covalent bonds extend in networks though solids that are not composed of discrete molecules (such as diamond or quartz or the silicate minerals in many types of rock) then the structures that result may be both strong and tough, at least in the direction oriented correctly with networks of covalent bonds. Also, the melting points of such covalent polymers and networks increase greatly.  Ionic bond, the bonding electron is not shared at all, but transferred. In this type of bond, the outer atomic orbital of one atom has a vacancy which allows addition of one or more electrons. These newly added electrons potentially occupy a lower energy-state (effectively closer to more nuclear charge) than they experience in a different atom. Thus, one nucleus offers a more tightly bound position to an electron than does another nucleus, with the result that one atom may transfer an electron to the other. This transfer causes one atom to assume a net positive charge, and the other to assume a net negative charge. The bond then results from electrostatic attraction between atoms, and the atoms become positive or negatively charged ions. Ionic bonds may be seen as extreme examples of polarization in covalent bonds. Often, such bonds have no particular orientation in space, since they

Understanding Biochem

There are two common (and equivalent) ways to describe molecular mass; both are used in this text. The first is molecular weight, or relative molecular mass, denoted Mr. The molecular weight of a substance is defined as the ratio of the mass of a molecule of that substance to one-twelfth the mass of carbon-12 (12C). Since Mr is a ratio, it is dimensionless—it has no associated units. The second is molecular mass, denoted m. This is simply the mass of one molecule, or the molar mass divided by Avogadro’s number. The molecular mass, m, is expressed in daltons (abbreviated Da). One dalton is equivalent to one-twelfth the mass of carbon-12; a kilodalton (kDa) is 1,000 daltons; a megadalton (MDa) is 1 million daltons. Consider, for example, a molecule with a mass 1,000 times that of water. We can say of this molecule either Mr =18,000 or m = 18,000 daltons. We can also describe it as an “18 kDa molecule.” However, the expression Mr 18,000 daltons is incorrect. Another convenient unit for describing the mass of a single atom or molecule is the atomic mass unit (formerly amu, now commonly denoted u). One atomic mass unit (1 u) is defined as onetwelfth the mass of an atom of carbon-12. Since the experimentally measured mass of an atom of carbon-12 is 1.9926 x 1023 g, 1 u = 1.6606 x 10-24 g. The atomic mass unit is convenient for describing the mass of a peak observed by mass spectrometry

result from equal electrostatic attraction of each ion to all ions around them. Ionic bonds are strong (and thus ionic substances require high temperatures to melt) but also brittle, since the forces between ions are short-range, and do not easily bridge cracks and fractures. This type of bond gives a charactistic physical character to crystals of classic mineral salts, such as table salt.

 Metallic bond. In this type of bonding, each atom in a metal donates one or more electrons to a "sea" of electrons that reside between many metal atoms. In this sea, each electron is free (by virtue of its wave nature) to be associated with a great many atoms at once.

The bond results because the metal atoms become somewhat positively charged due to loss of their electrons, while the electrons remain attracted to many atoms, without being part of any given atom. Metallic bonding may be seen as an extreme example of delocalization of electrons over a large system of covalent bonds, in which every atom participates. This type of bonding is often very strong (resulting in the tensile strength of metals). However, metallic bonds are more collective in nature than other types, and so they allow metal crystals to more easily deform, because they are composed of atoms attracted to each other, but not in any particularly-oriented ways. This results in the malleability of metals. The sea of electrons in metallic bonds causes the characteristically good electrical and thermal conductivity of metals, and also their "shiny" reflection of most frequencies of white light.  Aromatic (aryl) compounds  An aromatic (or aryl) compound contains a set of covalently bound atoms with specific characteristics  A delocalized conjugated π system, most commonly an arrangement of alternating single and double bonds  Coplanar structure, with all the contributing atoms in the same plane  Contributing atoms arranged in one or more rings  A number of π delocalized electrons that is even, but not a multiple of 4. That is, 4n + 2 number of π electrons, where n=0, 1, 2, 3, and so on. This is known as Hückel's Rule.

All bonds can be explained by quantum theory, but, in practice, simplification rules allow chemists to predict the strength, directionality, and polarity of bonds. The octet rule and VSEPR theory are two examples. More sophisticated theories are valence bond theory which includes orbital hybridization and resonance, and the linear combination of atomic orbitals molecular orbital method which includes ligand field theory. Electrostatics are used to describe bond polarities and the effects they have on chemical substances.

Table 1.3 Typical bond lengths in pm and bond energies in kJ/mol. Bond lengths can be converted to Å by division by 100 (1 Å = 100 pm).

H–H H–O

Bond Length (pm) H — Hydrogen 74 96

Energy (kJ/mol)

436 366

H–F 92

568 H–Cl 127 432 C — Carbon C–H 109 413 C–C 154 348 C–C= 151 =C–C≡ 147 =C–C= 148 C=C 134 614 C≡C 120 839 C–N 147 308 C–O 143 360 C–F 134 488 C–Cl 177 330 N — Nitrogen N–H 101 391 N–N 145 170 N≡N 110 945 O — Oxygen O–O 148 145 O=O 121 498 F, Cl, Br, I — Halogens F–F 142 158 Cl–Cl 199 243 Br–H 141 366 Br–Br 228 193 I–H 161 298 I–I 267 151

Figure 1.7 Benzene, the most widely recognized aromatic compound with six (4n + 2, n = 1) delocalized electrons.