1.3 Atomic Structure and Ground State Electron Configurations

In Section 1.2, we saw that carbon’s bonding characteristics are what give rise to the large variety of organic molecules. Those bonding characteristics, and the bonding characteristics of all atoms, are governed by the electrons that the atom has.

Basic structure of the atom, which includes a nucleus that is surrounded by a cloud of electrons. The nucleus includes protons and neutrons. The caption reads, Basic structure of the atom: Atoms are composed of a nucleus surrounded by a cloud of electrons. Protons, in white, and neutrons, in gray, make up the nucleus. This figure is not to scale. If it were, the size of the electron cloud, which is much larger than the size of the nucleus, would have a radius on the order of 500 meters!
FIGURE 1-4 Basic structure of the atom Atoms are composed of a nucleus surrounded by a cloud of electrons. Protons (white) and neutrons (gray) make up the nucleus. (This figure is not to scale. If it were, the size of the electron cloud, which is much larger than the size of the nucleus, would have a radius on the order of 500 meters!)

This section, then, is devoted to the nature of electrons in atoms. We first review the basic structure of an atom, followed by a discussion of orbitals and shells. Finally, we review electron configurations, distinguishing between valence electrons—electrons that can be used for bonding—and core electrons.

1.3a The Structure of the Atom

At the center of an atom (Fig. 1-4) is a positively charged nucleus, composed of protons and neutrons. Surrounding the nucleus is a cloud of negatively charged electrons, attracted to the nucleus by simple electrostatic forces (the forces by which opposite charges attract one another and like charges repel one another). Individual electrons are incredibly small, even much smaller than the nucleus, but the space that electrons occupy (i.e., the electron cloud) is much larger than the nucleus. In other words:

Chemistry with Chicken Wire

Even though carbon takes center stage in organic chemistry, organic molecules invariably include other atoms as well, such as hydrogen, nitrogen, oxygen, and halogen atoms. Some of the most exciting chemistry today, however, involves extended frameworks of only carbon. A single flat sheet of such a framework is called graphene, and resembles molecular chicken wire. Wrapped around to form a cylinder, a graphene sheet forms what is called a carbon nanotube. Pure carbon can even take the form of a soccer ball—the so-called buckminsterfullerene.

Three ball-and-stick models of different structures formed only with carbon: grapheme, a carbon nanotube, and Buckminsterfullerene. The first illustration shows the structure of graphene, which has a flat, sheet-like framework formed by hexagonal rings made up of six carbon atoms connected together. The second shows the structure of a carbon nanotube, which is cylindrical in shape and is also formed by hexagonal rings made up of six carbon atoms connected together. The third illustration shows the spherical structure of Buckminsterfullerene formed by pentagonal rings made up of five carbon atoms and hexagonal rings made up of six carbon rings connected together.

These structures themselves have quite interesting electronic properties, giving them a bright future in nanoelectronics. Carbon nanotubes and buckminsterfullerenes have high tensile strength, moreover, giving them potential use for structural reinforcement in concrete, sports equipment, and body armor. Chemical modification gives these structures an even wider variety of potential uses. Graphene oxide, for example, has promising antimicrobial activity, and attaching certain molecular groups to the surface of a carbon nanotube or buckminsterfullerene has potential for use as drug carriers for cancer therapeutics.

●  The size of an atom is essentially defined by the size of its electron cloud.

●  The vast majority of an electron cloud (and thus the vast majority of an atom) is empty space.

Table 1-1 Charges and Masses of Subatomic Particles

Particle

Charge (e)a

Mass (u)b

Proton

+ 1

~1        

Neutron

   0

~1        

Electron

–1

~0.0005

ae = Elementary charge.

bu = Unified atomic mass unit.

Table 1-1 lists the mass and charge of each of these elementary particles. Notice that the masses of the proton and neutron are significantly greater than that of the electron, so the mass of an atom is essentially the mass of just the nucleus.

An atom, by definition, has no net charge. Consequently, the number of electrons in an atom must equal the number of protons. The number of protons in the nucleus, called the atomic number (Z), defines the element. For example, a nucleus that has six protons has an atomic number of 6, and can only be a carbon nucleus.

If the number of protons and the number of electrons are unequal, then the entire species (that particular combination of protons, neutrons, and electrons) bears a net charge, and is called an ion. A negatively charged ion, an anion (pronounced AN-eye-on), results from an excess of electrons. A positively charged ion, a cation (pronounced CAT-eye-on), results from a deficiency of electrons.

Solved Problem 1.1

How many protons and electrons does a cation of the carbon atom have if its net charge is +1?

Think
SHOW SECTION

How many protons are there in the nucleus of a carbon atom? Does a cation have more protons than electrons, or vice versa? How many more, given the net charge of the species?

Solve
SHOW SECTION

A carbon atom’s nucleus has six protons. A cation with a + 1 charge should have one more proton than it has electrons, so this species must have five electrons.

problem 1.2 (a) How many protons and electrons does an anion of the carbon atom have if its net charge is –1? (b) How many protons and electrons does a cation of the oxygen atom have if its net charge is +1? (c) How many protons and electrons does an anion of the oxygen atom have if its net charge is –1?

1.3b Atomic Orbitals and Shells

Electrons in an isolated atom reside in atomic orbitals. As we shall see, the exact location of an electron can never be pinpointed. An orbital, however, specifies the region of space where the probability of finding a given electron is high. More simplistically, we can view orbitals as “rooms” that house electrons. Atomic orbitals are examined in greater detail in Chapter 3; for now, it will suffice to review some of their more basic concepts.

An illustration shows the space-filling models of a spherical s orbital and a dumbbell-shaped p orbital. The caption reads, �Orbitals: Orbitals represent regions in space where an electron is likely to be. An s orbital is spherical, and a p orbital is a dumbbell.�
FIGURE 1-5 Orbitals Orbitals represent regions in space where an electron is likely to be. An s orbital is spherical, and a p orbital is a dumbbell.

●  Atomic orbitals have different shapes. An s orbital, for example, is a sphere, whereas a p orbital has a dumbbell shape with two lobes (Fig. 1-5). Each orbital is centered on the nucleus of its atom or ion.

●  Atomic orbitals are organized in shells (also known as energy levels). A shell is defined by the principal quantum number, n. There are an infinite number of shells in an atom, given that n can assume any integer value from 1 to infinity.

●  The first shell (n = 1) contains only an s orbital, called 1s.

●  The second shell (n = 2) contains one s orbital and three p orbitals, called 2s, 2px, 2py, and 2pz.

●  The third shell (n = 3) contains one s orbital, three p orbitals, and five d orbitals.

●  Up to two electrons are allowed in any orbital.

●  Therefore, the first shell can contain up to two electrons (a duet).

●  The second shell can contain up to eight electrons (an octet).

●  The third shell can contain up to 18 electrons.

●  With increasing shell number, the size and energy of the atomic orbital increase. For example, comparing s orbitals in the first three shells, the size and energy increase in the order 1s < 2s < 3s, as shown in Figure 1-6. Similarly, a 2p orbital is smaller in size and lower in energy than a 3p orbital.

●  Within a given shell, an atomic orbital’s energy increases in the following order: s < p < d, etc. In the second shell, for example, the 2s orbital is lower in energy than the 2p.

An illustration shows space-filling models of the 1s, 2s, and 3s orbitals in order of increasing orbital size and orbital energy. The 1s orbital is the smallest in size and has the least orbital energy. The 2s orbital is larger and has more orbital energy than the 1s orbital. The 3s orbital, which is the largest of these three, has the most orbital energy. The caption reads, �Relationship between principal quantum number, orbital size, and orbital energy: As the shell number of an orbital increases, its size and energy increase, too. The horizontal black lines indicate each orbital�s energy.�
FIGURE 1-6 Relationship between principal quantum number, orbital size, and orbital energy As the shell number of an orbital increases, its size and energy increase, too. The horizontal black lines indicate each orbital’s energy.

1.3c Ground State Electron Configurations: Valence Electrons and Core Electrons

The way in which electrons are arranged in atomic orbitals is called the atom’s electron configuration. The most stable (i.e., the lowest energy) electron configuration is called the ground state configuration. Knowing an atom’s ground state configuration provides insight into the atom’s chemical behavior, as we will see.

With the relative energies of atomic orbitals established, an atom’s ground state electron configuration can be obtained by applying the following three rules:

An energy diagram shows the order in which the first eighteen electrons fill into the s orbitals and p orbitals. An upward arrow indicates increasing energy levels. Beside the arrow is a block representing the s orbitals in order of increasing energy, with the 1s orbital at the lowest end, the 2s orbital above it, and the 3s orbital further above. Beside this block is another representing the p orbitals, with three 2p orbitals located slightly above the level of the 2s orbital, and three 3p orbitals located slightly above the level of the 3s orbital. The first two electrons occupy the 1s orbital and the next two occupy the 2s orbital. The fifth, sixth, and seventh electrons occupy each of the three 2p orbitals in order, and the eighth, ninth, and tenth electrons pair up with the three previous electrons respectively. The eleventh and twelfth electrons occupy the 3s orbital, and the thirteenth, fourteenth, and fifteenth electrons occupy each of the three 3p orbitals in order. The sixteenth, seventeenth, and eighteenth electrons pair up with the three previous electrons respectively. Each electron pair is represented by a pair of opposing vertical arrows. The caption reads, �Energy diagram of atomic orbitals for the first 18 electrons: The order of electron filling is indicated in parentheses. Each horizontal black line represents a single orbital. Each successive electron fills the lowest energy orbital available. Notice in the 2p and 3p sets of orbitals that no electrons are paired up until the addition of the fourth electron.�
FIGURE 1-7 Energy diagram of atomic orbitals for the first 18 electrons The order of electron filling is indicated in parentheses. Each horizontal black line represents a single orbital. Each successive electron fills the lowest energy orbital available. Notice in the 2p and 3p sets of orbitals that no electrons are paired up until the addition of the fourth electron.

1. Pauli’s exclusion principle: No more than two electrons (i.e., zero, one, or two electrons) can occupy a single orbital; two electrons in the same orbital must have opposite spins.

2. Aufbau principle: Each successive electron must fill the lowest energy orbital available.

3. Hund’s rule: Before a second electron can be paired in the same orbital, all other orbitals at the same energy must contain a single electron.

According to these three rules, the first 18 electrons fill orbitals as indicated in Figure 1-7. Each arrow represents an electron, and the direction of the arrow—up or down—represents the electron’s spin.

YOUR TURN 1.1
SHOW ANSWERS

In Figure 1-7, place a box around all of the orbitals in the second shell and label them.

The 2s and 2p orbitals are in the second shell. Electrons (3)–(10) in Figure 1-7 are in the second shell.

An energy diagram shows how the six electrons in a carbon atom fill the s orbitals and p orbitals. An upward arrow indicates increasing energy levels. Beside the arrow is a block representing the s orbitals in order of increasing energy, with the 1s orbital at the lowest end and the 2s orbital above it. Beside this block is another representing the p orbitals, with three 2p orbitals located slightly above the level of the 2s orbital. The first two electrons occupy the 1s orbital and the next two occupy the 2s orbital. The fifth and sixth electrons occupy two of the three 2p orbitals in order. Each electron pair is represented by a pair of opposing vertical arrows, and the fifth and sixth electrons are each represented by an upward arrow. The caption reads, �Energy diagram for the ground state electron configuration of the carbon atom: This configuration is abbreviated 1s2 2s2 2p2.�
FIGURE 1-8 Energy diagram for the ground state electron configuration of the carbon atom This configuration is abbreviated 1s22s22p2.

In the ground state, the six electrons found in a carbon atom would fill the orbitals as shown in Figure 1-8, with two electrons in the 1s orbital, two electrons in the 2s orbital, and one electron in each of two different 2p orbitals (it doesn’t matter which two). The shorthand notation for this electron configuration is 1s22s22p2.

Knowing the ground state electron configuration of an atom, we can distinguish valence electrons from core electrons.

● Valence electrons are those occupying the highest energy (i.e., valence) shell. For the carbon atom, the valence shell is the n = 2 shell.

● Core electrons occupy the remaining lower energy shells of the atom. For the carbon atom, the core electrons occupy the n = 1 shell.

Valence electrons are important because, as we discuss in Section 1.5, they participate in covalent bonds. As we can see in Figure 1-8, for example, carbon has four valence electrons and two core electrons, so bonding involving carbon is governed by those four valence electrons.

YOUR TURN 1.2
SHOW ANSWERS

In Figure 1-8, place a circle around the valence electrons and label them. Place a box around all of the core electrons and label them.

The valence electrons are in the second shell, 2s2 and 2p2, and the core electrons are in the first shell, 1s2.

We can use the periodic table to quickly determine how many valence electrons an atom has (a copy of the periodic table appears inside the book’s front cover).

The number of valence electrons in an atom is the same as the atom’s group number.

Carbon is located in group 4A, consistent with its four valence electrons, whereas chlorine (group 7A) has seven. According to its ground state electron configuration (1s22s22p63s23p5), chlorine’s valence electrons occupy the third shell.

Atoms are especially stable when they have completely filled valence shells. This is exemplified by the noble gases (group 8A), such as helium and neon, because they have completely filled valence shells and they do not form bonds to make compounds. Although the specific origin of this “extra” stability is beyond the scope of this book, the consequences are the basis for the octet and duet rules we routinely use when drawing Lewis structures (Section 1.5).

Solved Problem 1.3

Write the ground state electron configuration of the nitrogen atom. How many valence electrons does it have? How many core electrons does it have?

Think
SHOW SECTION

How many total electrons are there in a nitrogen atom? What is the order in which the atomic orbitals should be filled (see Fig. 1-7)? What is the valence shell and where do the core electrons reside?

Solve
SHOW SECTION

There are seven total electrons (Z = 7 for N). The first two are placed in the 1s orbital and the next two in the 2s orbital, leaving one electron for each of the three 2p orbitals. The electron configuration is 1s22s22p3. The valence shell is the second shell, so there are five valence electrons and two core electrons.

problem 1.4 Write the ground state electron configuration of the oxygen atom. How many valence electrons and how many core electrons are there?