Unit One

Day 7: Covalent Molecular Substances and Lewis Structures

As you work through this section, if you find that you need a bit more background material to help you understand the topics at hand, you can consult “Chemistry: The Molecular Science” (5th ed. Moore and Stanitski) Chapter 6-2 through 6-4 and Chapter 2-7, and/or Chapter 4.5-4.7 in the Additional Reading Materials section.

D7.1 Covalent Molecular Substances

A covalent molecular substance consists of molecules. A molecule is a group of atoms held together by covalent bonding. Examples are water, sucrose (table sugar), carbon dioxide (dry ice), and iodine (See Figure 1).

Two images are shown and labeled “carbon dioxide” and “iodine.” The carbon dioxide structure is composed of molecules, each made up of one gray and two red atoms, stacked together into a cube. The image of iodine shows pairs of purple atoms arranged near one another, but not touching.
Figure 1. Carbon dioxide (CO2) forms a covalent molecular solid with a melting point of −78 °C. Iodine (I2) forms a covalent molecular solid that melts at 114 °C. Notice the CO2 and I2 molecules in each solid.

Activity 1: Reflection

In your notebook, make a heading for Molecules and Lewis Structures. Then write a few sentences describing your understanding of covalently bonded molecules and their structures. As you work through this section, make additional notes to review later.

Covalent molecular substances have properties that are significantly different from the properties of metals (Day 4, Section 6) or ionic compounds (Day 5, Section 6) and therefore it is useful to be able to classify a substance within one of these three categories. Typically, a covalent molecular substance consists of non-metals or metalloids—elements from the upper right of the periodic table (excepting noble gases but including H). Whereas metals and ionic compounds are usually solids at room temperature, many covalent molecular substances are liquids or gases; that is, they melt (and some boil) below room temperature or not too far above. Covalent molecular substances do not conduct electricity well as solids or liquids, the solids are weak and brittle or soft and waxy, and many are insoluble in water.

Activity 2: Types of Substances

In your notebook, make a table that lists characteristic properties of covalent molecular substances, ionic compounds, and metals. Include enough properties so that another student could use the table to distinguish each kind of substance based on a list of the properties. When you complete your table, click here to see an example table.

For each type of substance, write in your course notebook an explanation for each property based on  atomic-level structure.

In a molecule, atoms are connected by covalent bonds that result from overlap of atomic orbitals. There are several important aspects of molecular structure:

  • What kinds of atoms and how many of each are in the molecule?
  • Which atoms are bonded to which other atoms?
  • What are the distances between atoms that result in lowest energy? (Day 6, Section 1, Figure 4 shows the distance between two H atoms where energy is lowest—the bond length.)
  • How strong are the bonds?
  • What are the angles between the bonds?
  • What is the arrangement of the atoms in three-dimensional space?

These factors determine molecular structure; molecular structure determines the properties of covalent molecular substances.

For example, in a water molecule the two H and one O atom are arranged with the H atoms each bonded to the O atom. The bonds are strong—separating the atoms is difficult. The angle between the two O–H bonds is 104.5°, somewhat more than a right angle. Angles larger or smaller than 104.5° result in higher energy (lower stability). The properties of water would be quite different if the three atoms were all on a line (bond angle of 180°).

Ball and stick and space-filling models of a water molecule showing the bond angle of 104.5°.

Exercise 1: Classifying Substances

D7.2 Structures of Covalent Molecules

As we noted at the end of Day 6, molecular orbitals for molecules with three or more atoms are complicated and hard to draw. Thus, although MOs would convey a more descriptive and accurate picture of electron distribution within a molecule, chemists often rely on simpler diagrams to depict the atoms and bonds. Keep in mind, however, that there are aspects of molecular structure that such simpler diagrams do not represent.

Activity 3: Lewis Diagrams

Consider each element listed below. In your course notebook write the electron configuration for an atom of each element, determine the number of valence electrons, and write a Lewis diagram (see section D4.4). How are the Lewis diagrams related to the position of each element in the periodic table?

N          C          S          As         O          Br          F          Si          H

Lewis diagrams for the constituent atoms can be combined to make a Lewis structure for a molecule. A Lewis structure shows each single bond in a molecule as a pair of electrons shared between two adjacent atoms. Valence electrons that are not in a bond are shown as dots associated with individual atoms. Here are Lewis diagrams for two chlorine atoms and a Lewis structure for a Cl2 molecule:

A Lewis dot diagram shows a reaction. Two chlorine symbols, each surrounded by seven dots are separated by a plus sign. The dots on the first atom are all black and the dots on the second atom are all read. The phrase, “Chlorine atoms” is written below. A right-facing arrow points to two chlorine symbols, each with six dots surrounding their outer edges and a shared pair of dots in between. One of the shared dots is black and one is red. The phrase, “Chlorine molecule” is written below.

In Cl2 each Cl atom has three lone pairs (unshared electrons not used in bonding ) and shares one pair of electrons with the other Cl atom, forming a bond pair. Hence, each Cl atom is surrounded by eight valence electrons. Instead of a pair of dots, a bond pair is typically represented by a line—a single bond:

Two Lewis structures are shown. The left-hand structure shows two H atoms connected by a single bond. The right-hand structure shows two C l atoms connected by a single bond and each surrounded by six dots.

Activity 4: Lewis Structure and Electron Sharing

The Octet Rule

The term octet rule refers to the tendency of  atoms of main-group elements to gain, lose, or share enough electrons to form an octet: eight valence electrons (a noble gas electron configuration).

The Lewis diagram for an atom can be used to predict the number of bonds the atom will form. For example, a carbon atom has four valence electrons and therefore requires four more electrons to reach an octet:

Because a hydrogen atom needs only two electrons to fill its valence shell, H is an exception to the octet rule and forms only one bond. The transition elements also do not follow the octet rule.

Exercise 2: Number of Bonds

Double and Triple Bonds

Two atoms may need to share more than one pair of electrons to achieve the requisite octet. A double bond consists of two pairs of electrons are shared between two atoms. For example:
Two pairs of Lewis structures are shown. The left pair of structures shows a carbon atom forming single bonds to two hydrogen atoms. There are four electrons between the C atom and an O atom. The O atom also has two pairs of dots. The word “or” separates this structure from the same diagram, except this time there are two bond lines between the C atom and O atom. The name, “Formaldehyde” is written below these structures. Two more structures are on the right. The left shows two C atoms with four dots in between them and each C atom forming single bonds to two H atoms. The word “or” precedes the second structure, which is the same except that the C atoms are connected by two bond lines. The name, “ethene (ethylene)” is written below these structures.

A triple bond forms when three pairs of electron are shared between two atoms, as in carbon monoxide and the cyanide ion:

Two pairs of Lewis structures are shown. The left pair of structures show a C atom and an O atom with six dots in between them and a lone pair on each. The word “or” and the same structure with a triple bond in between the C atom and O atom also are shown. The name “Carbon monoxide” is written below these structures. The right pair of structures show a C atom and an N atom with six dots in between them, a lone pair on each, and a negative charge. The word “or” and the same structure with a triple bond in between the C atom and N atom also are shown. The name “Cyanide ion” is written below these structures.

Activity 5: Double and Triple Bonds

Write answers to these questions in your course notebook:

Write a Lewis structure for N2 and a Lewis structure for O2. Describe the type of bond in each case.

Do the N atoms in N2 and the O atoms in O2 follow the rule for number of bonds in Exercise 2?

Use the molecular-orbital energy-level diagram in Day 6 to calculate the bond order for N2 and for O2. How do the bond orders relate to the Lewis structures?

D7.3 Writing Lewis Structures with the Octet Rule

Here is a step-by-step procedure for drawing Lewis structures of molecules and polyatomic ions:

  1. Determine the total number of valence electrons. For cations, subtract one electron for each positive charge. For anions, add one electron for each negative charge.
  2. Choose one or more central atoms; a central atom bonds to several other atoms and is usually the atom the forms the greatest number of bonds. (Usually the central atom is written first in a chemical formula, such as P in PCl3.) If there are two or more central atoms, connect them using single bond lines.
  3. Draw a skeleton structure of the molecule by arranging the other atoms (which are called terminal atoms) around the central atom(s). Connect terminal atoms to the central atom(s) by single bond lines.
  4. Distribute the remaining electrons as lone pairs on the terminal atoms (except hydrogen), completing an octet around each atom.
  5. If there are still valence electrons available, place them on the central atom(s).
  6. If the number of electrons around a central atom is less than an octet, rearrange the electrons  to make multiple bonds with the central atom(s) until each atom has an octet.

Let’s apply these rules to a simple molecule, ammonia, NH3.

Here is a more complicated case: ethene (ethylene), C2H4.

Exercise 3: Lewis Structures and Valence Electrons

If you were to build a Lewis structure for nitrate ion, how many electrons would you need to allocate in your structure? (In other words, how many non-core electrons have to be in your structure?)

Exercise 4: Identifying Incorrect Lewis Structures

For each of the following Lewis structure, identify whether the structure is correct or not.  If the structure is incorrect, identify the error made in the representation.

D7.4 Hydrocarbons: Alkanes

Typical properties of covalent molecular substances are exemplified by hydrocarbons, compounds that contain only the elements carbon and hydrogen. Hydrocarbon molecules can have a variety of shapes, which chemists usually depict using Lewis structures: long chains of C atoms connected by single bonds; chains with branches; chains folded back on themselves to form rings; chains, branched chains, or rings that include double or triple bonds. Many hydrocarbons are found in plants, animals, and their fossils; other hydrocarbons have been prepared in the laboratory. We use hydrocarbons every day, mainly as fuels, such as natural gas (mainly methane in the U.S.), acetylene, propane, butane, gasoline, diesel fuel, and heating oil. The familiar plastics polyethylene, polypropylene, and polystyrene are also hydrocarbons.

Exercise 5: Combustion of Octane

When one mole of octane (C8H18) burns in air, how many moles of carbon dioxide (CO2) will form? How many moles of water (H2O) will form?

The simplest hydrocarbons are alkanes, which contain only single covalent bonds between carbon atoms. Each of the carbon atoms in an alkane is bonded to four other atoms, each of which is either carbon or hydrogen. Noncyclic alkanes have the general formula CnH2n+2, where the number of H atoms is two more than twice the number of C atoms. Alkanes are also called saturated hydrocarbons, because each C atom is saturated—bonded to the maximum possible number of H atoms.

Lewis structures indicate atomic connectivity (which atoms are connected to which other atoms) but this does not indicate the molecular shape in three dimensions. In a Lewis structure the carbon atoms are usually along a straight lines and unbranched alkanes are often called “straight-chain” alkanes. However, as the ball-and-stick and space-filling models of pentane in Figure 2 show, the C atoms do not lie in a straight line. The carbon-carbon bond angle is actually about 109°.

The figure illustrates four ways to represent molecules for molecules of methane, ethane, and pentane. In the first row of the figure, Lewis structural formulas show element symbols and bonds between atoms. Methane has a central C atom with four H atoms bonded to it. Ethane has a C atom with three H atoms bonded to it. The C atom is also bonded to another C atom with three H atoms bonded to it. Pentane has a C atom with three H atoms bonded to it. The C atom is bonded to another C atom with two H atoms bonded to it. The C atom is bonded to another C atom with two H atoms bonded to it. The C atom is bonded to another C atom with two H atoms bonded to it. The C atom is bonded to another C atom with three H atoms bonded to it. In the second row, ball-and-stick models are shown. In these representations, bonds are represented with sticks, and elements are represented with balls. Carbon atoms are black and hydrogen atoms are white in this image. In the third row, space-filling models are shown. In these models, atoms are enlarged and pushed together, without sticks to represent bonds. The molecule names and structural formulas are provided in the fourth row. Methane is named and represented with a condensed structural formula as C H subscript 4. Ethane is named and represented with two structural formulas C H subscript 3 C H subscript 3 and C subscript 2 H subscript 6. Pentane is named and represented as both C H subscript 3 C H subscript 2 C H subscript 2 C H subscript 2 C H subscript 3 and C subscript 5 H subscript 12.
Figure 2. Pictured are the Lewis structures, ball-and-stick models, and space-filling models for molecules of methane, ethane, and pentane.

The structures of alkanes may also be represented by condensed structural formulas, such as CH3CH3, shown below the names in Figure 2. Condensed structural formulas indicate how many H atoms are bonded to each C atom; they are related to Lewis structures, but all the bond symbols have been removed.

Recall from Day 6, Section 2, that the molecular orbital of a σ covalent bond has cylindrical symmetry along the internuclear axis. This means that regardless of how two σ-bonded atoms are rotated around the internuclear axis, a σ bond can form between them and the molecule’s energy changes very little. Hence, the atom at one end of a single bond can rotate easily relative to the atom at the other end. Rotations around C–C bonds in an alkane produce different molecular shapes. The various molecular shapes are called conformational isomers (or conformers). They have the same molecular formula and same atomic connectivity. Because rotation around single bonds occurs readily at room temperature, we cannot isolate one conformer from another. Conformers are the same chemical compound, with the same name and the same physical properties.

Figure 3. Two conformational isomers (conformers) of ethane.  Click on each image to see a rotatable 3D view of each conformer.
Figure 4. Move the slider at the bottom of the figure to see different conformations of nonane, C9H20. Each different conformation is the result of rotations around single C–C bonds.

D7.5 Constitutional Isomers

Compounds with the same molecular formula but different atomic connectivity  are called constitutional isomers (or structural isomers). For example, there are two alkanes with the formula C4H10:

The figure illustrates three ways to represent molecules of n dash butane and 2 dash methlylpropane. In the first row of the figure, Lewis structural formulas show element symbols and bonds between atoms. The n dash butane molecule shows 4 carbon atoms represented by the letter C bonded in a straight horizontal chain with hydrogen atoms represented by the letter H bonded above and below all carbon atoms. H atoms are bonded at the ends to the left and right of the left-most and right-most C atoms. In the second row, ball-and-stick models are shown. In these representations, bonds are represented with sticks, and elements are represented with balls. Carbon atoms are black and hydrogen atoms are white in this image. In the third row, space-filling models are shown. In these models, atoms are enlarged and pushed together, without sticks to represent bonds. The molecule names are provided in the fourth row.

Activity 6: Analyzing Constitutional Isomers

In your course notebook write an answer to each question and an explanation of your answer:

Describe the shape of the n-butane molecule and how it differs from the shape of the 2-methylpropane molecule.

Analyze the bonding in each molecule. Are all C atoms bonded to the same number of H atoms? If not, is the number of similarly bonded C atoms the same in each of the two structures ?

Could you convert one structure to the other solely by rotating around one or more C–C bonds? To convert one structure to the other, would one or more chemical bonds need to be broken so that atoms could be re-arranged?

Would you expect the physical properties (such as melting point and boiling point) to be different for the two substances?

The n– in n-butane stands for normal, an unbranched carbon chain. Typically the n– is omitted and the compound is just named “butane”; it is assumed that you know that no prefix refers to the unbranched chain. Because it is an isomer of butane, 2-methylpropane is sometimes called isobutane.

To change one constitutional isomer to another requires breaking and re-forming chemical bonds; breaking bonds requires significant input of energy and at room temperature very few molecules have that much energy. Therefore, structural isomers can be synthesized and separated from one another: they are different substances.

Because Lewis structures are not intended to indicate the 3D geometry, Lewis structures that look different may actually represent the same substance. For example, the three structures below all represent the same molecule, butane, and hence are not different isomers.


One way to tell whether a Lewis structure is the same as another is to work out the name of the compound. Names are designed so that the same structure always has the same name; if two structures are different, naming them correctly will result in different names. You will not be explicitly tested on naming compounds (nomenclature) in this course, but it is useful to know how to name common molecules that you will encounter.  You can read more about alkane nomenclature in the appendix, as well as Appendix E in “Chemistry: The Molecular Science” (5th ed. Moore and Stanitski).

Exercise 6: Naming Alkanes

Activity 7: Wrap-up

In your notebook, refer back to what you wrote about Molecules and Lewis Structures at the beginning of this day’s work. Update what you wrote and make a list of the main things you learned as you studied the unit. Your list should provide a summary you can use to review later for an exam.

Podia Question

Compare each Lewis structure to the initial structure. Classify each structure as the same as the initial structure, an isomer of  the initial structure, or neither the same as nor an isomer of the initial structure. Explain your reasoning using scientifically appropriate language.

 

Two days before the next whole-class session, this Podia question will become live on Podia, where you can submit your answer.

 

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Chemistry 109 Copyright © by John Moore; Jia Zhou; and Etienne Garand is licensed under a Creative Commons Attribution-NonCommercial-ShareAlike 4.0 International License, except where otherwise noted.