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Gray: Molecules

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The formation of atoms from fundamental particles, interesting as this might be to the physicist, is far from being the ultimate stage in the organization of matter. As we mentioned earlier, when atoms are close enough to one another that the outer electrons of one atom can interact with the other atoms, then attractions can be set up between atoms, strong enough to hold them together in what is termed a chemical bond. In the simplest cases the bond arises from the sharing of two electrons between a pair of atoms, with one electron provided by each of the bonded atoms. Bonds based on electron sharing are known as covalent bonds, and two or more atoms held together as a unit by covalent bonds are known as a molecule. One of the principal triumphs of the theory of quantum mechanics in chemistry (see Chapter 8) has been its ability to predict the kinds of atoms that will bond together, and the three-dimensional structures and reactivities of the molecules that result. In molecular diagrams, a covalent, electron-sharing bond is represented by a straight line connecting the bonded atoms. In the water molecule, one atom of oxygen (O) is bonded to two hydrogen (H) atoms. The diagram for the molecule can be drawn two ways: 


The second version acknowledges the fact that a water molecule is not linear; the two H—O bonds make an angle of 104.5° with one another. Molecules of hydrogen gas, hydrogen sulfide, ammonia, methane, and methyl alcohol (methanol) have the following bond structures:



Figure 1.2: Shapes and relative sizes of some sample molecules. Two bonded atoms appear to interpenetrate because their electron clouds overlap. By convention, a tapered bond in a drawing represents a bond pointing out toward the observer, with the wide end or the taper closest, and a dashed line is used for a bond that points back behind the plane of the page.

These diagrams show only the connections between atoms in the molecules. They do not show the three-dimensional geometries (or shapes) of the
molecules. Figure 1-2 shows the shapes and the relative bulk of several molecules. Note that the bond angle in molecules having more than two atoms can vary. The angle in the water molecule is 104.5°, and the angle in hydrogen sulfide is 92.1°; the four atoms connected to the central carbon in
methane and methyl alcohol are directed to the four corners of a tetrahedron. The bond structure in straight-chain octane, one of the components of
gasoline, is   


Each of the molecular diagrams shown can be condensed to a molecular formula, which tells how many atoms of each element are in the molecule,
but provides little or no information as to how the atoms are connected. The molecular formula for hydrogen is H2; water, H2O, hydrogen sulfide,
H2S; ammonia, NH3; methane, CH4; methyl alcohol, CH2OH; and octane, C8H18. The formula for octane can also be written


The sum of the atomic weights of all the atoms in a molecule is its molecular weight. Using the atomic weights from the periodic table, we can calculate molecular weights. The molecular weight of hydrogen, H2, is

2 x 1.0080 u = 2.0160 u

A water molecule, H2O, has two atoms of hydrogen and one atom of oxygen, so:

(2 X 1.0080 amu) + [15.9994 amu) = 18.0154 amu



Calculate the molecular weight of methyl alcohol.
Solution The molecular formula is CH3OH or CH4O. Then:

1 carbon:         1 x 12.011 amu  = 12.011 amu
4 hydrogens:   4 x 1.008 amu    =   4.032 amu
1 oxygen:        1 x 15.999 amu  = 15.999 amu
Total molecular weight:                  32.04  amu

(If you wonder why the last figure has been dropped, see the discussion of significant figures in Appendix 4.)

In Example 7 (example above, but not numbered) notice that the natural atomic weight of carbon is not 12.000 amu but 12.011 amu, since carbon occurs as a mixture of 98.89% carbon-12 and 1.11% carbon-13, with trace amounts of carbon-14.


What is the molecular weight of pure octane?
Solution Since the molecular formula is C8H18, the molecular weight is

(8 x 12.011) + (18 x 1.008) = 114.23 amu



  • Dickerson, Richard E. and Gray, Harry B. and Haight, Gilbert P (1979) Chemical principles.

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