Ionic bonding typically occurs when it is easy for one atom to lose one or more electrons and another atom to gain one or more electrons. However, some atoms won’t give up or gain electrons easily. Yet they still participate in compound formation. How?
There is another mechanism for obtaining a complete valence shell: sharing electrons. When electrons are shared between two atoms, they make a bond called a covalent bond.
You are watching: Covalent bonds are sometimes shown using electron dot diagrams
Let us illustrate a covalent bond by using H atoms, with the understanding that H atoms need only two electrons to fill the 1s subshell. Each H atom starts with a single electron in its valence shell:
Because two atoms are sharing one pair of electrons, this covalent bond is called a single bond.
As another example, consider fluorine. F atoms have seven electrons in their valence shell:
Covalent bonds can be made between different elements as well. One example is HF. Each atom starts out with an odd number of electrons in its valence shell:
Use Lewis electron dot diagrams to illustrate the covalent bond formation in HBr.
HBr is very similar to HF, except that it has Br instead of F. The atoms are as follows:
Use Lewis electron dot diagrams to illustrate the covalent bond formation in Cl2.
More than two atoms can participate in covalent bonding, although any given covalent bond will be between two atoms only. Consider H and O atoms:
Use a Lewis electron dot diagram to show the covalent bonding in NH3.
The N atom has the following Lewis electron dot diagram:
Use a Lewis electron dot diagram to show the covalent bonding in PCl3.
We also use Lewis symbols to indicate the formation of covalent bonds, which are shown in Lewis structures, drawings that describe the bonding in molecules and polyatomic ions. For example, when two chlorine atoms form a chlorine molecule, they share one pair of electrons:
The Lewis structure indicates that each Cl atom has three pairs of electrons that are not used in bonding (called lone pairs) and one shared pair of electrons (written between the atoms). A dash (or line) is sometimes used to indicate a shared pair of electrons:
A single shared pair of electrons is called a single bond. Each Cl atom interacts with eight valence electrons: the six in the lone pairs and the two in the single bond.
The Octet Rule
The other halogen molecules (F2, Br2, I2, and At2) form bonds like those in the chlorine molecule: one single bond between atoms and three lone pairs of electrons per atom. This allows each halogen atom to have a noble gas electron configuration. The tendency of main group atoms to form enough bonds to obtain eight valence electrons is known as the octet rule.
The number of bonds that an atom can form can often be predicted from the number of electrons needed to reach an octet (eight valence electrons); this is especially true of the nonmetals of the second period of the periodic table (C, N, O, and F). For example, each atom of a group 14 element has four electrons in its outermost shell and therefore requires four more electrons to reach an octet. These four electrons can be gained by forming four covalent bonds, as illustrated here for carbon in CCl4 (carbon tetrachloride) and silicon in SiH4 (silane). Because hydrogen only needs two electrons to fill its valence shell, it is an exception to the octet rule. The transition elements and inner transition elements also do not follow the octet rule:
Group 15 elements such as nitrogen have five valence electrons in the atomic Lewis symbol: one lone pair and three unpaired electrons. To obtain an octet, these atoms form three covalent bonds, as in NH3 (ammonia). Oxygen and other atoms in group 16 obtain an octet by forming two covalent bonds:
Double and Triple Bonds
As previously mentioned, when a pair of atoms shares one pair of electrons, we call this a single bond. However, a pair of atoms may need to share more than one pair of electrons in order to achieve the requisite octet. A double bond forms when two pairs of electrons are shared between a pair of atoms, as between the carbon and oxygen atoms in CH2O (formaldehyde) and between the two carbon atoms in C2H4 (ethylene):
A triple bond forms when three electron pairs are shared by a pair of atoms, as in carbon monoxide (CO) and the cyanide ion (CN–):
Writing Lewis Structures with the Octet Rule
For very simple molecules and molecular ions, we can write the Lewis structures by merely pairing up the unpaired electrons on the constituent atoms. See these examples:
For more complicated molecules and molecular ions, it is helpful to follow the step-by-step procedure outlined here:Determine the total number of valence (outer shell) electrons. For cations, subtract one electron for each positive charge. For anions, add one electron for each negative charge.Draw a skeleton structure of the molecule or ion, arranging the atoms around a central atom. (Generally, the least electronegative element should be placed in the center.) Connect each atom to the central atom with a single bond (one electron pair).Distribute the remaining electrons as lone pairs on the terminal atoms (except hydrogen), completing an octet around each atom.Place all remaining electrons on the central atom.Rearrange the electrons of the outer atoms to make multiple bonds with the central atom in order to obtain octets wherever possible.
Let us determine the Lewis structures of SiH4, CHO2−, NO+, and OF2 as examples in following this procedure:Determine the total number of valence (outer shell) electrons in the molecule or ion.For a molecule, we add the number of valence electrons on each atom in the molecule:
For a negative ion, such as CHO2−, we add the number of valence electrons on the atoms to the number of negative charges on the ion (one electron is gained for each single negative charge):
For a positive ion, such as NO+, we add the number of valence electrons on the atoms in the ion and then subtract the number of positive charges on the ion (one electron is lost for each single positive charge) from the total number of valence electrons:
Draw a skeleton structure of the molecule or ion, arranging the atoms around a central atom and connecting each atom to the central atom with a single (one electron pair) bond. (Note that we denote ions with brackets around the structure, indicating the charge outside the brackets:)
Polyatomic ions are bonded together with covalent bonds, as seen in the example of CHO2−. Because they are ions, however, they participate in ionic bonding with other ions. So both major types of bonding can occur at the same time.
NASA’s Cassini-Huygens mission detected a large cloud of toxic hydrogen cyanide (HCN) on Titan, one of Saturn’s moons. Titan also contains ethane (H3CCH3), acetylene (HCCH), and ammonia (NH3). What are the Lewis structures of these molecules?
SolutionCalculate the number of valence electrons.HCN: (1 × 1) + (4 × 1) + (5 × 1) = 10H3CCH3: (1 × 3) + (2 × 4) + (1 × 3) = 14HCCH: (1 × 1) + (2 × 4) + (1 × 1) = 10NH3: (5 × 1) + (3 × 1) = 8Draw a skeleton and connect the atoms with single bonds. Remember that H is never a central atom:
Test yourselfBoth carbon monoxide, CO, and carbon dioxide, CO2, are products of the combustion of fossil fuels. Both of these gases also cause problems: CO is toxic and CO2 has been implicated in global climate change. What are the Lewis structures of these two molecules?
What is the proper Lewis electron dot diagram for CO2?
The central atom is a C atom, with O atoms as surrounding atoms. We have a total of 4 + 6 + 6 = 16 valence electrons. Following the rules for Lewis electron dot diagrams for compounds gives us
What is the proper Lewis electron dot diagram for carbonyl sulfide (COS)?
Carbon soot has been known to man since prehistoric times, but it was not until fairly recently that the molecular structure of the main component of soot was discovered. In 1996, the Nobel Prize in Chemistry was awarded to Richard Smalley (Figure 1), Robert Curl, and Harold Kroto for their work in discovering a new form of carbon, the C60 buckminsterfullerene molecule (Figure 1 in Chapter 8 Introduction). An entire class of compounds, including spheres and tubes of various shapes, were discovered based on C60. This type of molecule, called a fullerene, shows promise in a variety of applications. Because of their size and shape, fullerenes can encapsulate other molecules, so they have shown potential in various applications from hydrogen storage to targeted drug delivery systems. They also possess unique electronic and optical properties that have been put to good use in solar powered devices and chemical sensors.
Exceptions to the Octet Rule
Many covalent molecules have central atoms that do not have eight electrons in their Lewis structures. These molecules fall into three categories:Odd-electron molecules have an odd number of valence electrons, and therefore have an unpaired electron.Electron-deficient molecules have a central atom that has fewer electrons than needed for a noble gas configuration.Hypervalent molecules have a central atom that has more electrons than needed for a noble gas configuration.Examples of these will be covered later chemistry courses.
Food and Drink App: Vitamins and Minerals
Vitamins are nutrients that our bodies need in small amounts but cannot synthesize; therefore, they must be obtained from the diet. The word vitamin comes from “vital amine” because it was once thought that all these compounds had an amine group (NH2) in it. This is not actually true, but the name stuck anyway.
All vitamins are covalently bonded molecules. Most of them are commonly named with a letter, although all of them also have formal chemical names. Thus vitamin A is also called retinol, vitamin C is called ascorbic acid, and vitamin E is called tocopherol. There is no single vitamin B; there is a group of substances called the B complex vitamins that are all water soluble and participate in cell metabolism. If a diet is lacking in a vitamin, diseases such as scurvy or rickets develop. Luckily, all vitamins are available as supplements, so any dietary deficiency in a vitamin can be easily corrected.
A mineral is any chemical element other than carbon, hydrogen, oxygen, or nitrogen that is needed by the body. Minerals that the body needs in quantity include sodium, potassium, magnesium, calcium, phosphorus, sulfur, and chlorine. Essential minerals that the body needs in tiny quantities (so-called trace elements) include manganese, iron, cobalt, nickel, copper, zinc, molybdenum, selenium, and iodine. Minerals are also obtained from the diet. Interestingly, most minerals are consumed in ionic form, rather than as elements or from covalent molecules. Like vitamins, most minerals are available in pill form, so any deficiency can be compensated for by taking supplements.
Figure 2. Vitamin and Mineral Supplements
Every entry down through pantothenic acid is a vitamin, and everything from calcium and below is a mineral.
Key Concepts and Summary
Valence electronic structures can be visualized by drawing Lewis symbols (for atoms and monatomic ions) and Lewis structures (for molecules and polyatomic ions). Lone pairs, unpaired electrons, and single, double, or triple bonds are used to indicate where the valence electrons are located around each atom in a Lewis structure. Most structures—especially those containing second row elements—obey the octet rule, in which every atom (except H) is surrounded by eight electrons. Exceptions to the octet rule occur for odd-electron molecules (free radicals), electron-deficient molecules, and hypervalent molecules.
1. Write the Lewis symbols for each of the following ions:
a) As3– b) I– c) Be2+ d) O2– e) Ga3+ f) Li+ g) N3–
2. Write the Lewis symbols of the ions in each of the following ionic compounds and the Lewis symbols of the atom from which they are formed:
a) MgS b) Al2O3 c) GaCl3 d) K2O e) Li3N f) KF
3. Write the Lewis structure for the diatomic molecule P2, an unstable form of phosphorus found in high-temperature phosphorus vapor.
4. Write Lewis structures for the following:
a) O2 b) H2CO c) AsF3 d) ClNO e) SiCl4
f) H3O+ g) NH4+ h) BF4− i) HCCH j) ClCN k) C22+
5. Write Lewis structures for the following:
a) SeCl3+ b) Cl2BBCl2 (contains a B–B bond)
6. Correct the following statement: “The bonds in solid PbCl2 are ionic; the bond in a HCl molecule is covalent. Thus, all of the valence electrons in PbCl2 are located on the Cl– ions, and all of the valence electrons in a HCl molecule are shared between the H and Cl atoms.”
7. Methanol, H3COH, is used as the fuel in some race cars. Ethanol, C2H5OH, is used extensively as motor fuel in Brazil. Both methanol and ethanol produce CO2 and H2O when they burn. Write the chemical equations for these combustion reactions using Lewis structures instead of chemical formulas.
8. Carbon tetrachloride was formerly used in fire extinguishers for electrical fires. It is no longer used for this purpose because of the formation of the toxic gas phosgene, Cl2CO. Write the Lewis structures for carbon tetrachloride and phosgene.
9. The arrangement of atoms in several biologically important molecules is given here. Complete the Lewis structures of these molecules by adding multiple bonds and lone pairs. Do not add any more atoms.
a) the amino acid serine:
c) pyruvic acid:
e) carbonic acid:
10. A compound with a molar mass of about 42 g/mol contains 85.7% carbon and 14.3% hydrogen by mass. Write the Lewis structure for a molecule of the compound.
11. How are single, double, and triple bonds similar? How do they differ?
12. How many electrons will be in the valence shell of H atoms when it makes a covalent bond?
13. What is the Lewis electron dot diagram of I2? Circle the electrons around each atom to verify that each valence shell is filled.
14. What is the Lewis electron dot diagram of NCl3? Circle the electrons around each atom to verify that each valence shell is filled.
15. Draw the Lewis electron dot diagram for each substance. a) SF2 b) BH4−
16. Draw the Lewis electron dot diagram for each substance. a) GeH4 b) ClF
17. Draw the Lewis electron dot diagram for each substance. Double or triple bonds may be needed.
a) SiO2 b) C2H4 (assume two central atoms)
18. Draw the Lewis electron dot diagram for each substance. Double or triple bonds may be needed.
1. a) eight electrons:
b) eight electrons:
c) no electrons: Be2+
d) eight electrons:
e) no electrons: Ga3+
f) no electrons: Li+
g) eight electrons:
In this case, the Lewis structure is inadequate to depict the fact that experimental studies have shown two unpaired electrons in each oxygen molecule.
5. a) SeCl3+:
6. Two valence electrons per Pb atom are transferred to Cl atoms; the resulting Pb2+ ion has a 6s2 valence shell configuration. Two of the valence electrons in the HCl molecule are shared, and the other six are located on the Cl atom as lone pairs of electrons.
11. Each bond includes a sharing of electrons between atoms. Two electrons are shared in a single bond; four electrons are shared in a double bond; and six electrons are shared in a triple bond.
double bond: covalent bond in which two pairs of electrons are shared between two atoms
free radical: molecule that contains an odd number of electrons
hypervalent molecule: molecule containing at least one main group element that has more than eight electrons in its valence shell
Lewis structure: diagram showing lone pairs and bonding pairs of electrons in a molecule or an ion
Lewis symbol: symbol for an element or monatomic ion that uses a dot to represent each valence electron in the element or ion
lone pair: two (a pair of) valence electrons that are not used to form a covalent bond
octet rule: guideline that states main group atoms will form structures in which eight valence electrons interact with each nucleus, counting bonding electrons as interacting with both atoms connected by the bond
single bond: bond in which a single pair of electrons is shared between two atoms
triple bond: bond in which three pairs of electrons are shared between two atoms