6.1: Ionic and Covalent Bonding
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- Scott Van Bramer
- Widener University
Introduction
Learning Objectives
- Explain the formation of cations, anions, and ionic compounds
- Predict the charge of common metallic and nonmetallic elements, and write their electron configurations
- Describe the formation of covalent bonds
- Define electronegativity and assess the polarity of covalent bonds
It has long been known that pure carbon occurs in different forms (allotropes) including graphite and diamonds. But it was not until 1985 that a new form of carbon was recognized: buckminsterfullerene, commonly known as a “buckyball” and is shown in Figure \(\PageIndex{1}\). This molecule was named after the architect and inventor R. Buckminster Fuller (1895–1983), whose signature architectural design was the geodesic dome, characterized by a lattice shell structure supporting a spherical surface. Experimental evidence revealed the formula, C 60 , and then scientists determined how 60 carbon atoms could form one symmetric, stable molecule. They were guided by bonding theory—the topic of this chapter—which explains how individual atoms connect to form more complex structures.
Ionic Bonding
As you have learned, ions are atoms or molecules bearing an electrical charge. A cation is a positive ion that forms when a neutral atom loses one or more electrons from its valence shell. An an anion is a negative ion that forms when a neutral atom gains one or more electrons in its valence shell. The opposite charges of a cation and an anion cause an attraction that forms an ionic bond.
Compounds composed of ions are called ionic compounds, or salts, and their constituent ions are held together by ionic bonds : electrostatic forces of attraction between oppositely charged cations and anions. The properties of ionic compounds shed some light on the nature of ionic bonds. Ionic solids exhibit a crystalline structure and tend to be rigid and brittle; they also tend to have high melting and boiling points, which suggests that ionic bonds are very strong. Ionic solids are also poor conductors of electricity for the same reason—the strength of ionic bonds prevents ions from moving freely in the solid state. Most ionic solids, however, dissolve readily in water. Once dissolved or melted, ionic compounds are excellent conductors of electricity and heat because the ions can move about freely.
Neutral atoms and charged ions have very different physical and chemical properties. Figure \(\PageIndex{2}\) shows sodium metal, a soft, silvery-white metal that burns vigorously in air and reacts explosively with water, which is made of of sodium atoms. It also shows chlorine gas, a yellow-green gas that is extremely corrosive to most metals and very poisonous to animals and plants, which is made up of chlorine atoms combined into molecules of chlorine or Cl 2 . Sodium metal reacts vigerously with chlorine gas to form sodium chloride, a white crystalline compound called table salt. Sodium chloride contains sodium cations and chloride anions and has properties entirely different from sodium metal and chlorine gas. Sodium chloride is essential to life and dissolves in water while sodium atoms react vigorously with water and chlorine gas is poisonous.
The Formation of Ionic Compounds
Binary ionic compounds are composed of just two elements: a metal (which forms the cations) and a nonmetal (which forms the anions). For example, NaCl is a binary ionic compound. We can think about the formation of such compounds in terms of the periodic properties of the elements. Many metallic elements have relatively low ionization potentials and lose electrons easily. These elements lie to the left in a period or near the bottom of a group on the periodic table. Nonmetal atoms have relatively high electron affinities and thus readily gain electrons lost by metal atoms, thereby filling their valence shells. Nonmetallic elements are found in the upper-right corner of the periodic table.
Since all substances must be electrically neutral, the total number of positive charges on the cations of an ionic compound must equal the total number of negative charges on its anions. The formula of an ionic compound represents the simplest ratio of the numbers of ions necessary to give identical numbers of positive and negative charges. For example, the formula for aluminum oxide, Al 2 O 3 , indicates that this ionic compound contains two aluminum cations, Al 3 + , for every three oxide anions, O 2− [thus, (2 × +3) + (3 × –2) = 0].
It is important to note that the formula for an ionic compound does not represent the physical arrangement of its ions. Sodium chloride (NaCl) is not a molecule with one sodium atom connected to one chlorine atom. Instead each sodium ion is attracted to several chloride ions in different directions. This results in the ions arranging themselves into a tightly bound, three-dimensional structure called a lattice. The structure of this lattice depends upon the size and charge if the ions. Figure \(\PageIndex{3}\) shows the arragement for sodium chloride, which is a regular arrangement of equal numbers of Na + cations and Cl – anions.
The strong electrostatic attraction between Na + and Cl – ions holds them tightly together in solid NaCl. It requires 769 kJ of energy to dissociate one mole of solid NaCl into separate gaseous Na + and Cl – ions:
\[\ce{NaCl}(s)⟶\ce{Na+}(g)+\ce{Cl-}(g)\hspace{20px}ΔH=\mathrm{769\:kJ} \nonumber \]
Electronic Structures of Cations
When forming a cation, an atom of a main group element tends to lose all of its valence electrons, thus assuming the electronic structure of the noble gas that precedes it in the periodic table. For groups 1 (the alkali metals) and 2 (the alkaline earth metals), the group numbers are equal to the numbers of valence shell electrons and, consequently, to the charges of the cations formed from atoms of these elements when all valence shell electrons are removed. For example, calcium is a group 2 element whose neutral atoms have 20 electrons and a ground state electron configuration of 1 s 2 2 s 2 2 p 6 3 s 2 3 p 6 4 s 2 . When a Ca atom loses both of its valence electrons, the result is a cation with 18 electrons, a 2+ charge, and an electron configuration of 1 s 2 2 s 2 2 p 6 3 s 2 3 p 6 . The Ca 2 + ion is therefore isoelectronic with the noble gas Ar.
Example \(\PageIndex{1}\): Determining the Electronic Structures of Cations
Potassium is required in our diet. Write the electron configuration for a potassium ion.
Solution
First, write the electron configuration for the neutral atom:
- K: [Ar]4 s 1
Next, remove electrons from the highest energy orbital. Potassium is a member of group 1, so it should have a charge of 1+, and thus loses one electron from its s orbital. This gives the following electron configuration for the ion:
- K 2 + : [Ar]
Exercise \(\PageIndex{1}\)
Magnesium is also required in our diet. Write the electron configurations of the ion.
- Answer
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Mg 2 + : [Ne]
Electronic Structures of Anions
Most monatomic anions form when a neutral nonmetal atom gains enough electrons to completely fill its outer s and p orbitals, thereby reaching the electron configuration of the next noble gas. Thus, it is simple to determine the charge on such a negative ion: The charge is equal to the number of electrons that must be gained to fill the s and p orbitals of the parent atom. Oxygen, for example, has the electron configuration 1 s 2 2 s 2 2 p 4 , whereas the oxygen anion has the electron configuration of the noble gas neon (Ne), 1 s 2 2 s 2 2 p 6 . The two additional electrons required to fill the valence orbitals give the oxide ion the charge of 2– (O 2– ).
Example \(\PageIndex{2}\): Determining the Electronic Structure of Anions
Selenium and iodine are two essential trace elements that form anions. Write the electron configurations of the anions.
Solution
Se 2 – : [Ar]3 d 10 4 s 2 4 p 6
I – : [Kr]4 d 10 5 s 2 5 p 6
Exercise \(\PageIndex{2}\)
Write the electron configurations of a phosphorus atom and its negative ion. Give the charge on the anion.
- Answer
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P: [Ne]3 s 2 3 p 3
P 3– : [Ne]3 s 2 3 p 6
Covalent Bonds
In ionic compounds, electrons are transferred between atoms of different elements to form ions. But this is not the only way that compounds can be formed. Atoms can also make chemical bonds by sharing electrons between each other. Such bonds are called covalent bonds . Covalent bonds are formed between two atoms when both have similar tendencies to attract electrons to themselves (i.e., when both atoms have identical or fairly similar ionization energies and electron affinities). For example, two hydrogen atoms bond covalently to form an H 2 molecule; each hydrogen atom in the H 2 molecule shares an electron with the other hydrogen atom. This sharing of electrons forms a covalent bond.
Compounds that contain covalent bonds exhibit different physical properties than ionic compounds. In ionic compounds all the ions are strongly bound to each other in the lattice structure. With covalent bonds, the atoms are held together in distinct molecules. The attraction between these atoms is strong but the attraction between atoms in different molecules is much weaker because the molecules have a neutral charge. This is why molecular compounds often have lower melting and boiling points than ionic compounds. Many covalent compounds are liquids or gases at room temperature and as solids they are usually softer than ionic solids. Ionic compounds are good conductors of electricity when dissolved in water but most covalent compounds are poor conductors of electricity in any state since the molecules do not have an overall charge.
Formation of Covalent Bonds
Nonmetal atoms frequently form covalent bonds with other nonmetal atoms. For example, the hydrogen molecule, H 2 , contains a covalent bond between its two hydrogen atoms. Figure \(\PageIndex{4}\) illustrates why this bond is formed. Starting on the far right, we have two separate hydrogen atoms with a particular potential energy, indicated by the red line. Along the x -axis is the distance between the two atoms. As the two atoms approach each other (moving left along the x -axis), their valence orbitals (1 s ) begin to overlap. The single electrons on each hydrogen atom then interact with both atomic nuclei, occupying the space around both atoms. The strong attraction of each shared electron to both nuclei stabilizes the system, and the potential energy decreases as the bond distance decreases. If the atoms continue to approach each other, the positive charges in the two nuclei begin to repel each other, and the potential energy increases. The bond length is determined by the distance where these opposing forces balance and the molecule has the lowest potential energy.
It is essential to remember that energy must be added to break chemical bonds (an endothermic process), whereas forming chemical bonds releases energy (an exothermic process). In the case of H 2 , the covalent bond is very strong; a large amount of energy, 436 kJ, must be added to break the bonds in one mole of hydrogen molecules and cause the atoms to separate:
\[\ce{H2}(g)⟶\ce{2H}(g)\hspace{20px}ΔH=\mathrm{436\:kJ} \nonumber \]
Conversely, the same amount of energy is released when one mole of H 2 molecules forms from two moles of H atoms:
\[\ce{2H}(g)⟶\ce{H2}(g)\hspace{20px}ΔH=\mathrm{−436\:kJ} \nonumber \]
Pure vs. Polar Covalent Bonds
If the atoms that form a covalent bond are identical, as in H 2 , Cl 2 , and other diatomic molecules, then the electrons in the bond must be shared equally. We refer to this as a pure covalent bond . Electrons shared in pure covalent bonds have an equal probability of being near each nucleus. In the case of Cl 2 , each atom starts off with seven valence electrons, and each Cl shares one electron with the other, forming one covalent bond:
\[\ce{Cl + Cl⟶Cl2} \nonumber \]
The total number of electrons around each individual atom consists of six nonbonding electrons and two shared (i.e., bonding) electrons for eight total electrons, matching the number of valence electrons in the noble gas argon. Since the bonding atoms are identical, Cl 2 also features a pure covalent bond.
When the atoms linked by a covalent bond are different, the bonding electrons are shared, but no longer equally. Instead, the bonding electrons are more attracted to one atom than the other, giving rise to a shift of electron density toward that atom. This unequal distribution of electrons is known as a polar covalent bond , characterized by a partial positive charge on one atom and a partial negative charge on the other. The atom that attracts the electrons more strongly acquires the partial negative charge and vice versa. For example, the electrons in the H–Cl bond of a hydrogen chloride molecule spend more time near the chlorine atom than near the hydrogen atom. Thus, in an HCl molecule, the chlorine atom carries a partial negative charge and the hydrogen atom has a partial positive charge. Figure \(\PageIndex{5}\) shows the distribution of electrons in the H–Cl bond. Note that the shaded area around Cl is much larger than it is around H. Compare this to Figure \(\PageIndex{4}\), which shows the even distribution of electrons in the H 2 nonpolar bond.
We sometimes designate the positive and negative atoms in a polar covalent bond using a lowercase Greek letter “delta,” δ, with a plus sign or minus sign to indicate whether the atom has a partial positive charge (δ+) or a partial negative charge (δ–). This symbolism is shown for the H–Cl molecule in Figure \(\PageIndex{5b}\).
Electronegativity
Whether a bond is nonpolar or polar covalent is determined by a property of the bonding atoms called electronegativity . Electronegativity is a measure of the tendency of an atom to attract electrons (or electron density) towards itself. It determines how the shared electrons are distributed between the two atoms in a bond. The more strongly an atom attracts the electrons in its bonds, the larger its electronegativity. Electrons in a polar covalent bond are shifted toward the more electronegative atom; thus, the more electronegative atom is the one with the partial negative charge. The greater the difference in electronegativity, the more polarized the electron distribution and the larger the partial charges of the atoms.
Figure \(\PageIndex{6}\) shows the electronegativity values of the elements as proposed by one of the most famous chemists of the twentieth century: Linus Pauling. In general, electronegativity increases from left to right across a period in the periodic table and decreases down a group. Thus, the nonmetals, which lie in the upper right, tend to have the highest electronegativities, with fluorine the most electronegative element of all (EN = 4.0). Metals tend to be less electronegative elements, and the group 1 metals have the lowest electronegativities. Note that noble gases are excluded from this figure because these atoms usually do not share electrons with others atoms since they have a full valence shell. (While noble gas compounds such as XeO 2 do exist, they can only be formed under extreme conditions, and thus they do not fit neatly into the general model of electronegativity.)
Linus Pauling
Linus Pauling is the only person to have received two unshared (individual) Nobel Prizes: one for chemistry in 1954 for his work on the nature of chemical bonds and one for peace in 1962 for his fight against the nuclear arms race between East and West. He developed many of the theories and concepts that are foundational to our current understanding of chemistry, including electronegativity and resonance structures.
Pauling also contributed to many other fields besides chemistry. His research on sickle cell anemia revealed the cause of the disease—the presence of a genetically inherited abnormal protein in the blood—and paved the way for the field of molecular genetics. His work was also pivotal in curbing the testing of nuclear weapons; he proved that radioactive fallout from nuclear testing posed a public health risk.
Electronegativity versus Electron Affinity
We must be careful not to confuse electronegativity and electron affinity. The electron affinity of an element is a measurable physical quantity, namely, the energy released or absorbed when an isolated gas-phase atom acquires an electron, measured in kJ/mol. Electronegativity, on the other hand, describes how tightly an atom attracts electrons in a bond. It is a dimensionless quantity that is calculated, not measured. Pauling derived the first electronegativity values by comparing the amounts of energy required to break different types of bonds. He chose an arbitrary relative scale ranging from 0 to 4.
Electronegativity and Bond Type
The absolute value of the difference in electronegativity (ΔEN) of two bonded atoms provides a rough measure of the polarity to be expected in the bond and, thus, the bond type. When the difference is very small or zero, the bond is covalent and nonpolar. When it is large, the bond is polar covalent or ionic. The absolute values of the electronegativity differences between the atoms in the bonds H–H, H–Cl, and Na–Cl are 0 (nonpolar), 0.9 (polar covalent), and 2.1 (ionic), respectively. The degree to which electrons are shared between atoms varies from completely equal (pure covalent bonding) to not at all (ionic bonding).
The simplist guide to the covalent or ionic character of a bond is to consider the types of atoms involved and their relative positions in the periodic table. Bonds between two nonmetals are generally covalent; bonding between a metal and a nonmetal is often ionic. This idea is refined in Figure \(\PageIndex{7}\), which shows a rough approximation of the electronegativity differences associated with covalent, polar covalent, and ionic bonds. This table is a general guide but there are many exceptions. For example, the H and F atoms in HF have an electronegativity difference of 1.9, and the N and H atoms in NH 3 a difference of 0.9, yet both of these compounds form bonds that are considered polar covalent. Likewise, the Na and Cl atoms in NaCl have an electronegativity difference of 2.1, and the Mn and I atoms in MnI 2 have a difference of 1.0, yet both of these substances form ionic compounds.
Some compounds contain both covalent and ionic bonds. The atoms in polyatomic ions, such as OH – , \(\ce{NO3-}\), and \(\ce{NH4+}\), are held together by polar covalent bonds. However, these polyatomic ions form ionic compounds by combining with ions of opposite charge. For example, potassium nitrate, KNO 3 , contains the K + cation and the polyatomic \(\ce{NO3-}\) anion. Bonding in potassium nitrate is both ionic, from the electrostatic attraction between the ions K + and \(\ce{NO3-}\), and covalent, from the bonding between nitrogen and the oxygen atoms in \(\ce{NO3-}\).
Example \(\PageIndex{3}\): Ele ctronegativity and Bond Polarity
Bond polarities play an important role in determining the structure of proteins. Using the electronegativity values in Table A2, arrange the following covalent bonds—all commonly found in amino acids—in order of increasing polarity. Then designate the positive and negative atoms using the symbols δ+ and δ–:
C–H, C–N, C–O, N–H, O–H, S–H
Solution
The polarity of these bonds increases as the absolute value of the electronegativity difference increases. The atom with the δ– designation is the more electronegative of the two. Table \(\PageIndex{1}\) shows these bonds in order of increasing polarity.
| Bond | ΔEN | Polarity |
|---|---|---|
| C–H | 0.4 | \(\overset{δ−}{\ce C}−\overset{δ+}{\ce H}\) |
| S–H | 0.4 | \(\overset{δ−}{\ce S}−\overset{δ+}{\ce H}\) |
| C–N | 0.5 | \(\overset{δ+}{\ce C}−\overset{δ−}{\ce N}\) |
| N–H | 0.9 | \(\overset{δ−}{\ce N}−\overset{δ+}{\ce H}\) |
| C–O | 1.0 | \(\overset{δ+}{\ce C}−\overset{δ−}{\ce O}\) |
| O–H | 1.4 | \(\overset{δ−}{\ce O}−\overset{δ+}{\ce H}\) |
Exercise \(\PageIndex{3}\)
Silicones are polymeric compounds containing, among others, the following types of covalent bonds: Si–O, Si–C, C–H, and C–C. Using the electronegativity values in Figure \(\PageIndex{3}\), arrange the bonds in order of increasing polarity and designate the positive and negative atoms using the symbols δ+ and δ–.
Answer
| Bond | Electronegativity Difference | Polarity |
|---|---|---|
| C–C | 0.0 | nonpolar |
| C–H | 0.4 | \(\overset{δ−}{\ce C}−\overset{δ+}{\ce H}\) |
| Si–C | 0.7 | \(\overset{δ+}{\ce{Si}}−\overset{δ−}{\ce C}\) |
| Si–O | 1.7 | \(\overset{δ+}{\ce{Si}}−\overset{δ−}{\ce O}\) |
Summary
Covalent bonds form when electrons are shared between atoms and are attracted by the nuclei of both atoms. In pure covalent bonds, the electrons are shared equally. In polar covalent bonds, the electrons are shared unequally, as one atom exerts a stronger force of attraction on the electrons than the other. The ability of an atom to attract a pair of electrons in a chemical bond is called its electronegativity. The difference in electronegativity between two atoms determines how polar a bond will be. In a diatomic molecule with two identical atoms, there is no difference in electronegativity, so the bond is nonpolar or pure covalent. When the electronegativity difference is very large, as is the case between metals and nonmetals, the bonding is characterized as ionic.
Glossary
- bond length
- distance between the nuclei of two bonded atoms at which the lowest potential energy is achieved
- covalent bond
- bond formed when electrons are shared between atoms
- electronegativity
- tendency of an atom to attract electrons in a bond to itself
- polar covalent bond
- covalent bond between atoms of different electronegativities; a covalent bond with a positive end and a negative end
- pure covalent bond
- (also, nonpolar covalent bond) covalent bond between atoms of identical electronegativities
Summary
Atoms gain or lose electrons to form ions with particularly stable electron configurations. The charges of cations formed by the representative metals may be determined readily because, with few exceptions, the electronic structures of these ions have either a noble gas configuration or a completely filled electron shell. The charges of anions formed by the nonmetals may also be readily determined because these ions form when nonmetal atoms gain enough electrons to fill their valence shells.
Glossary
- inert pair effect
- tendency of heavy atoms to form ions in which their valence s electrons are not lost
- ionic bond
- strong electrostatic force of attraction between cations and anions in an ionic compound