tructure* 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 yourtraining 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 beused
throughout your study of organic chemistry.
ATOMS, ELECTRONS, AND ORBITALS
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. A neutral atom has equal numbers of protons, which are positivelycharged, 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 asif 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 functions and 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.
Study Guide TOC
FIGURE 1.1 Probability distribution ( 2) for an electron
in a 1s orbital.
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
electronat 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 thenucleus, 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. Forconvenience, chemists use the term
“orbital” in several different ways. A drawing 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, anddirectional properties.
Spherically symmetrical ones such as shown in Figure 1.1 are called s orbitals. The letter s is preceded by the principal quantum number n (n
1, 2, 3, etc.) which speciﬁes the shell and is related to the energy of the orbital. An electron in a 1s orbital is
likely to be found closer to the nucleus, is lower in energy, and is more strongly held
than an electron in a 2s...
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