Ionic, Covalent, and Polar Bonds

In trying to explain why atoms form bonds, G. N. Lewis proposed that an atom is most stable if its outer shell is either filled or contains eight electrons and it has no electrons of higher energy. According to Lewis’s theory, an atom will give up, accept, or share electrons in order to achieve a filled outer shell or an outer shell that contains eight electrons. This theory has come to be called the octet rule.

Lithium (Li) has a single electron in its 2s atomic orbital. If it loses this electron, the lithium atom ends up with a filled outer shell—a stable configuration. Removing an electron from an atom takes energy—called the ionization energy. Lithium has a relatively low ionization energy—the drive to achieve a filled outer shell with no electrons of higher energy causes it to lose an electron relatively easily. Sodium (Na) has a single electron in its 3s atomic orbital. Consequently, sodium also has a relatively low ionization energy because, when it loses an electron, it is left with an outer shell of eight electrons. Elements (such as lithium and sodium) that have low ionization energies are said to be electropositive—they readily lose an electron and thereby become
positively charged. The elements in the first column of the periodic table are all electropositive—each readily loses an electron because each has a single electron in its outermost shell.

Electrons in inner shells (those below the outermost shell) are called core electrons. Core electrons do not participate in chemical bonding. Electrons in the outermost shell are called valence electrons, and the outermost shell is called the valence shell. Carbon, for example, has two core electrons and four valence electrons.

Lithium and sodium each have one valence electron. Elements in the same column of the periodic table have the same number of valence electrons, and because the number of valence electrons is the major factor determining an element’s chemical properties, elements in the same column of the periodic table have similar chemical properties. Thus, the chemical behavior of an element depends on its electronic configuration.

When we draw the electrons around an atom, as in the following equations, core electrons are not shown; only valence electrons are shown. Each valence electron is shown as a dot. Notice that when the single valence electron of lithium or sodium is removed, the resulting atom—now called an ion—carries a positive charge.

Fluorine has seven valence electrons (Table 1.2). Consequently, it readily acquires an electron in order to have an outer shell of eight electrons. When an atom acquires an electron, energy is released. Elements in the same column as fluorine (e.g., chlorine, bromine, and iodine) also need only one electron to have an outer shell of eight, so they, too, readily acquire an electron. Elements that readily acquire an electron are said
to be electronegative—they acquire an electron easily and thereby become negatively charged.

Ionic Bonds
Because sodium gives up an electron easily and chlorine acquires an electron readily, when sodium metal and chlorine gas are mixed, each sodium atom transfers an electron to a chlorine atom, and crystalline sodium chloride (table salt) is formed as a result. The positively charged sodium ions and negatively charged chloride ions are independent species held together by the attraction of opposite charges. A bond is an attractive force between two atoms. Attractive forces between opposite charges are called electrostatic attractions. A bond that is the result of only electrostatic attractions is called an ionic bond. Thus, an ionic bond is formed when there is a transfer of electrons, causing one atom to become a positively charged ion and the other to become a negatively charged ion.

Sodium chloride is an example of an ionic compound. Ionic compounds are formed when an element on the left side of the periodic table (an electropositive element) transfers one or more electrons to an element on the right side of the periodic table (an electronegative element).

Covalent Bonds
Instead of giving up or acquiring electrons, an atom can achieve a filled outer shell by sharing electrons. For example, two fluorine atoms can each attain a filled shell of eight electrons by sharing their unpaired valence electrons. A bond formed as a result of sharing electrons is called a covalent bond.

Two hydrogen atoms can form a covalent bond by sharing electrons. As a result of covalent bonding, each hydrogen acquires a stable, filled outer shell (with two electrons).

Similarly, hydrogen and chlorine can form a covalent bond by sharing electrons. In doing so, hydrogen fills its only shell and chlorine achieves an outer shell of eight electrons.

A hydrogen atom can achieve a completely empty shell by losing an electron. Loss of its sole electron results in a positively charged hydrogen ion. 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. A hydrogen atom can achieve a filled outer shell by gaining an electron, thereby forming a negatively charged hydrogen ion, called a hydride ion.

Because oxygen has six valence electrons, 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 electrons, must form four covalent bonds to achieve a filled outer shell.

Polar Covalent Bonds

In the F - F and H -H covalent bonds shown previously, the atoms that share the bonding electrons are identical. Therefore, they share the electrons equally; that is, each electron spends as much time in the vicinity of one atom as in the other. An even (nonpolar) distribution of charge results. Such a bond is called a nonpolar covalent bond. In contrast, the bonding electrons in hydrogen chloride, water, and ammonia are more attracted to one atom than another because the atoms that share the electrons in these molecules are different and have different electronegativities. Electronegativity is the tendency of an atom to pull bonding electrons  oward itself. The bonding electrons in hydrogen chloride, water, and ammonia molecules are more attracted  on the atom with the greater electronegativity. This results in a polar distribution of charge. A polar covalent bond is a covalent bond between atoms of different electronegativities. The electronegativities of some of the elements are shown in Table 1.3. Notice that electronegativity increases as you go from left to right across a row of the periodic table or up any of the columns. A polar covalent bond has a slight positive charge on one end and a slight negative charge on the other. Polarity in a covalent bond is indicated by the symbols (S+)  and which denote partial positive and partial negative charges, respectively. The negative end of the bond is the end that has the more electronegative atom. The greater the difference in electronegativity between the bonded atoms, the more polar the bond will be.

The direction of bond polarity can be indicated with an arrow. By convention, the
arrow points in the direction in which the electrons are pulled, so 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.

You can think of ionic bonds and nonpolar covalent bonds as being at the opposite ends of a continuum of bond types. An ionic bond involves no sharing of electrons. A nonpolar covalent bond involves equal sharing. Polar covalent bonds fall somewhere in between, and the greater the difference in electronegativity between the atoms forming the bond, the closer the bond is to the ionic end of the continuum. C - H bonds are
relatively nonpolar, because carbon and hydrogen have similar electronegativities (electronegativity difference = 0.4 see Table 1.3). N - H bonds are relatively polar (electronegativity differnece = 0.9 ), but not as polar as O - H bonds (electronegativity differnece = 1.4). The bond between sodium and chloride ions is closer to the
ionic end of the continuum (electronegativity difference = 2.1 ), but sodium chloride is not as ionic as potassium fluoride (electronegativity difference = 3.2).

Basic Concept

The Structure of an Atom
An atom consists of a tiny dense nucleus surrounded by electrons that are spread throughout a relatively large volume of space around the nucleus. The nucleus contains positively charged protons and neutral neutrons, so it is positively charged. The electrons are negatively charged. Because the amount of positive charge on a proton equals the amount of negative charge on an electron, a neutral atom has an equal number of protons and electrons. Atoms can gain electrons and thereby become negatively charged, or they can lose electrons and become positively charged. However, the number of protons in an atom does not change. Protons and neutrons have approximately the same mass and are about 1800 times more massive than an electron. This means that most of the mass of an atom is in its nucleus. However, most of the volume of an atom is occupied by its electrons, and that is where our focus will be because it is the electrons that form chemical bonds.

The atomic number of an atom equals the number of protons in its nucleus. The atomic number is also the  number of electrons that surround the nucleus of a neutral atom. For example, the atomic number of carbon is 6, which means that a neutral carbon atom has six protons and six electrons. Because the number of protons in an atom does not change, the atomic number of a particular element is always the same—all carbon atoms have an atomic number of 6.

The mass number of an atom is the sum of its protons and neutrons. Not all carbon atoms have the same mass number, because, even though they all have the same number of protons, they do not all have the same number of neutrons. For example, 98.89% of naturally occurring carbon atoms have six neutrons—giving them a mass number of 12—and 1.11% have seven neutrons—giving them a mass number of 13.
These two different kinds of carbon atoms (C-12 and C-13 )and are called isotopes. Isotopes have the same atomic number (i.e., the same number of protons), but different mass numbers because they have different numbers of neutrons. The chemical properties of isotopes of a given element are nearly identical.

The atomic weight of a naturally occurring element is the average weighted mass of its atoms. Because an atomic mass unit (amu) is defined as exactly of the mass of the 1/12th of C-12. Atomic mass of C-12 is 12.0000 amu;


The Distribution of Electrons in an Atom


Electrons are moving continuously. Like anything that moves, electrons have kinetic energy, and this energy is what counters the attractive force of the positively charged protons that would otherwise pull the negatively charged electrons into the nucleus.

For a long time, electrons were perceived to be particles—infinitesimal “planets” orbiting the nucleus of an atom. In 1924, however, a French physicist named Louis de-Broglie showed that electrons also have wavelike properties. He did this by combining a formula developed by Einstein that relates mass and energy with a formula developed by Planck relating frequency and energy. The realization that electrons have
wavelike properties spurred physicists to propose a mathematical concept known as quantum mechanics.
Quantum mechanics uses the same mathematical equations that describe the wave motion of a guitar string to characterize the motion of an electron around a nucleus. The version of quantum mechanics most useful to chemists was proposed by Erwin Schrödinger in 1926. According to Schrödinger, the behavior of each electron in an atom or a molecule can be described by a wave equation. The solutions to the Schrödinger equation are called wave functions or orbitals. They tell us the energy of the electron and the volume of space around the nucleus where an electron is most likely to be found.

According to quantum mechanics, the electrons in an atom can be thought of as occupying a set of concentric shells that surround the nucleus. The first shell is the one closest to the nucleus. The second shell lies farther from the nucleus, and even farther out lie the third and higher numbered shells. Each shell contains subshells known as atomic orbitals. Each atomic orbital has a characteristic shape and energy and occupies a characteristic volume of space, which is predicted by the Schrödinger equation. An important point to remember is that the closer the atomic orbital is to the nucleus, the lower is its energy.

The first shell consists of only an s atomic orbital; the second shell consists of s and p atomic orbitals; the third shell consists of s, p, and d atomic orbitals; and the fourth and higher shells consist of s, p, d, and atomic orbitals (Table 1.1).

Each shell contains one s atomic orbital. The second and higher shells—in addition to their s orbital—each contain three degenerate p atomic orbitals. Degenerate orbitals are orbitals that have the same energy. The third and higher shells—in addition to their s and p orbitals—also contain five degenerate d atomic orbitals, and the fourth and higher shells f also contain seven degenerate atomic orbitals. Because a maximum of two electrons can coexist in an atomic orbital the first shell, with only one atomic orbital, can contain no more than two electrons. The second shell, with four atomic orbitals—one s and three p— can have a total of eight electrons. Eighteen electrons can occupy the nine atomic orbitals—one s, three p, and five d—of the third shell, and 32 electrons can occupy the 16 atomic orbitals of the fourth shell. In studying organic chemistry, we will be concerned primarily with atoms that have electrons only in the first and second shells. The ground-state electronic configuration of an atom describes the orbitals occupied by the atom’s electrons when they are all in the available orbitals with the lowest energy.

If energy is applied to an atom in the ground state, one or more electrons can jump into a higher energy orbital. The atom then would be in an excited-state electronic configuration.
The following principles are used to determine which orbitals electrons occupy:
1. The aufbau principle (aufbau is German for “building up”) tells us the first thing we need to know to be able to assign electrons to the various atomic orbitals.
According to this principle, an electron always goes into the available orbital with the lowest energy. The relative energies of the atomic orbitals are as follows:
Because a 1s atomic orbital is closer to the nucleus, it is lower in energy than a 2s atomic orbital, which is lower in energy—and is closer to the nucleus—than a 3s atomic orbital. Comparing atomic orbitals in the same shell, we see that an s atomic orbital is lower in energy than a p atomic orbital, and a p atomic orbital is
lower in energy than a d atomic orbital.

2. The Pauli exclusion principle states that (a) no more than two electrons can occupy each atomic orbital, and (b) the two electrons must be of opposite spin. It is called an exclusion principle because it states that only so many electrons can occupy any particular shell. Notice in Table 1.2 that spin in one direction is designated by an upward-pointing arrow, and spin in the opposite direction by a downward-pointing arrow.

1s < 2s < 2p < 3s < 3p < 4s < 3d < 4p < 5s < 4d < 5p < 6s < 4f < 5d < 6p < 7s < 5f

Because a 1s atomic orbital is closer to the nucleus, it is lower in energy than a 2s atomic orbital, which is lower in energy—and is closer to the nucleus—than a 3s atomic orbital. Comparing atomic orbitals in the same shell, we see that an s atomic orbital is lower in energy than a p atomic orbital, and a p atomic orbital is
lower in energy than a d atomic orbital.

2. The Pauli exclusion principle states that (a) no more than two electrons can occupy each atomic orbital, and (b) the two electrons must be of opposite spin. It is called an exclusion principle because it states that only so many electrons can occupy any particular shell. Notice in Table 1.2 that spin in one direction is designated by an upward-pointing arrow, and spin in the opposite direction by a
downward-pointing arrow.

From these first two rules, we can assign electrons to atomic orbitals for atoms that contain one, two, three, four, or five electrons. The single electron of a hydrogen atom occupies a 1s atomic orbital, the second electron of a helium atom fills the 1s atomic orbital, the third electron of a lithium atom occupies a 2s atomic orbital, the fourth electron of a beryllium atom fills the 2s atomic orbital, and the fifth electron of a boron
atom occupies one of the 2p atomic orbitals. (The subscripts x, y, and z distinguish the three 2p atomic orbitals.) Because the three p orbitals are degenerate, the electron can be put into any one of them. Before we can continue to larger atoms—those containing six or more electrons—we need Hund’s rule:

3. Hund’s rule states that when there are degenerate orbitals—two or more orbitals with the same energy—an electron will occupy an empty orbital before it will pair up with another electron. In this way, electron repulsion is minimized. The sixth electron of a carbon atom, therefore, goes into an empty 2p atomic orbital,
rather than pairing up with the electron already occupying a 2p atomic orbital. (See Table 1.2.) The seventh electron of a nitrogen atom goes into an empty 2p atomic orbital, and the eighth electron of an oxygen atom pairs up with an electron occupying a 2p atomic orbital rather than going into a higher energy 3s orbital.