Lewis Electron Dot Models
Lewis electron dot structures are a useful and popular way in chemistry classes to visualize the bonding in molecules. Practice making Lewis electron dot structures using bingo chips and cut-out element symbols for the first 18 elements and then for a series of molecules.
- Molecular bonding
- Lewis electron dot structures
A chemical bond is a strong attractive force between atoms or ions in a compound. Electrons in the outermost shells or orbitals are often the electrons involved in chemical bonding. One type of chemical bond between two atoms is called a covalent bond. The covalent bond consists of a pair of shared electrons from the outermost orbitals. The electrons in this outermost shell, or highest occupied energy level, have a special importance and are referred to as valence electrons. The electrons in the inner orbitals of an atom are referred to as core electrons. Since the maximum number of electrons in an orbital is 2 and since there is one s orbital and three p orbitals for any given energy shell, the maximum number of valence electrons is 8. Note: This holds for elements in the first three rows of the periodic table. Beyond that, exceptions may apply.
How are the number of valence electrons in an atom determined? Scientists discovered that there is a relationship between the organization of the periodic table and the electron configurations of the elements. All elements in a particular group of the periodic table have similar properties and thus the same number of valence electrons. In general then, the number of valence electrons is equal to the group number for each group of representative elements. The representative elements are designated by different headings on different periodic tables. Some tables use the American system IA to VIIIA group heading designations while other tables have the IUPAC recommendation with the 1–18 group headings. For those who prefer the IUPAC method, the number of valence electrons in columns 14–18 is the group number minus ten. Thus, for example, potassium has 1 valence electron—its group number is IA or 1. Bromine has 7 valence electrons—its group number is VIIA or 17.
Molecules are formed when atoms come together through chemical bonding. Many different molecular models have been devised as a way to visualize the arrangement of atoms in molecules. In 1916, G. N. Lewis, an American chemist, developed a system of arranging dots, representing valence electrons, around the symbols of the elements to form electron dot symbols, or Lewis electron dot symbols. Using this Lewis representation for the atom, the element symbol represents the nucleus and the inner core electrons of the atom while the dots represent the outer valence electrons involved in the bonding. This model of using electron dot structures, while not all-encompassing due to the variety of molecules in the world, provides a simple representation of many molecules. It is a useful tool and can lead to an increase in understanding the properties of many different molecules. The procedures for writing Lewis electron dot structures for individual atoms as well as for a variety of molecules are described.
Writing Lewis Electron Dot Structures for Atoms
A Lewis electron dot structure for an atom includes the element symbol and the “electron” dots. The number of dots is equal to the number of valence electrons which, as explained previously, is equal to the group number of the periodic table where the element is located. The maximum number of “electron” dots on any atom is 8. Thus the Lewis electron dot structure for potassium, which is located in Group IA or 1 on the periodic table, contains only one dot (see below) while the Lewis electron dot structure for bromine, which is located in Group VII or 17 on the periodic table, contains seven dots:
The exact placement of the dots is arbitrary, so the single dot around K can be on any of the four sides. Writing Lewis Electron Dot Structures for Molecules
A Lewis electron dot structure for a molecule includes each atom in the molecule connected by bonds. Each bond may be drawn with two electron dots or, more commonly, with a single line. In order to predict the bonding arrangement that occurs between atoms in a molecule, the stable nature of noble gases must be recognized. Noble gases, group VIII of the periodic table, are unreactive and form very few compounds. This stability is due to the outermost s
orbitals being completely filled and containing two s
electrons and six p
electrons. Helium has a filled valence orbital with only two s
electrons while all of the other noble gases have eight electrons in their valence orbitals, two in the s
orbital and six in the p
orbital. From these observations of the noble gases, a rule called the octet rule
was developed. The octet rule assumes that atoms form bonds to achieve stability and thus a noble gas configuration, in which each atom has eight electrons (an octet) in its valence orbitals. Thus, through electron sharing, each atom in a molecule can attain the electron configuration of a noble gas.
To write a Lewis dot structure for a molecule, a simple process may be followed. Determine the number of valence electrons supplied by each atom involved. (See step-by-step procedure that follows.) Visualize each atom donating its valence electrons to a large electron “pool.” When the atoms come together to form a molecule, picture the electrons from the “pool” being distributed to the molecule so that each atom in the molecule attains a noble gas structure. Some electrons may need to be shared to form bonds while others may be left unshared. In General
- H atoms want to be surrounded by two valence electrons, to attain the stability of He.
- All other nonmetal atoms want to be surrounded by eight valence electrons, and thus should follow the “octet rule.” Note: There are many exceptions to the “octet rule” which are beyond the scope of this introductory activity.
- Count the total number of valence electrons available. To do this, determine the number of valence electrons supplied by each atom in the compound and add them up. Remember that:
- Group IA or 1 atoms (H, Li, ...) have 1 valence electron.
- Group IIA or 2 atoms (Be, Mg, ...) have 2 valence electrons.
- Group IIIA or 13 atoms (B, Al, ...) have 3 valence electrons.
- Group IVA or 14 atoms (C, Si, ...) have 4 valence electrons.
- Group VA or 15 atoms (N, P, ...) have 5 valence electrons.
- Group VIA or 16 atoms (O, S, ...) have 6 valence electrons.
- Group VIIA or 17 atoms (F, Cl, ...) have 7 valence electrons.
- Group VIIIA or 18 atoms (Ne, Ar, ...) have 8 valence electrons except He which has only 2 valence electrons.
- Polyatomic anions—add one electron for each unit of negative charge.
- Cations—subtract one electron for each unit of positive charge.
Example: CCl4—C is in group IVA or 14, so it has 4 valence electrons. Cl is in group VIIA or 17, so each Cl has 7 valence electrons x 4 chlorine atoms equals 28 electrons. Total valence electrons for CCl4 = 32.
- Draw a “skeleton structure” for the molecule, joining atoms by single bonds.
The skeleton structure for the molecule, X2, is:
The skeleton structure for the molecule, XY, is:
For more complex molecules, more than one skeleton structure is possible. In general, if there is a single atom of one element and several atoms of another element, the single atom is usually the central atom. Thus, the skeleton structure for the molecule, XY2, is most likely
- Nature tends toward symmetry, so it is best to make your skeleton structure symmetrical when possible.
- Hydrogen is never the central atom; it is always a surrounding atom.
- Group VII halogens tend to be surrounding atoms rather than central atoms.
- Carbon tends to be a central atom.
Example: CCl4—Carbon tends to be a central atom and halogens tend to be surrounding atoms. Since nature tends toward symmetry, the following structure is proposed:
- From the total number of valence electrons, subtract two for each single bond in the skeleton structure to determine how many valence electrons are left to distribute.
Example: CCl4 has 32 total valence electrons. Subtract 2 electrons for each single bond x 4 single bonds. This leaves 32 – 8 or 24 remaining valence electrons.
- Distribute the remaining valence electrons as unshared pairs around the atoms in the molecule. Follow the octet rule, distributing electrons to give each atom a total of 8 electrons (except H which should have 2).
Example: Carbon has 4 bonds, or 8 electrons already surrounding it. Thus, the 24 remaining valence electrons are distributed around the 4 chlorine atoms to give each chlorine 6 paired electrons and one bond, or 8 total electrons.
- If this point is reached and there are too few valence electrons to give each atom an octet, a multiple bond is most likely necessary. While a single bond involves the sharing of one pair (or 2) electrons, multiple bonds involve the sharing of more than one pair of electrons. A double bond involves the sharing of two pairs (or 4) electrons, and a triple bond involves the sharing of three pairs (or 6) electrons. Thus the electrons in multiple bonds are counted in the octets of both of the bonded atoms.
H is in group IA or 1 with 1 valence electron; C is in group IVA or 14 with 4 valence electrons; and N is in group VA or 15 with 5 valence electrons. Total valence electrons = 10. The proposed skeleton structure can be drawn as follows, since C tends to be a central atom and H is always a surrounding atom:
Proposed Skeleton Structure
From the total of 10 valence electrons, subtract 2 for each single bond, leaving 6 extra valence electrons to be distributed. Try distributing the electrons around the nitrogen atom:
This structure is incorrect because there are too few electrons to give each atom an octet. Remember that each bond counts as 2 electrons. Hydrogen is stable with 2 electrons and nitrogen is stable with 8 electrons surrounding it; however, carbon only has 4 electrons rather than the octet it needs to be stable. Therefore, a multiple bond is most likely necessary. Try a double bond:
The structure is still incorrect because, although nitrogen has 8 electrons surrounding it, carbon only has 6 electrons. This is not enough to complete its octet, so try a triple bond:
This structure is now complete and correct. Hydrogen is stable with 2 electrons (1 bond), carbon is stable with 8 electrons (4 bonds) and nitrogen is stable with 8 electrons (3 bonds plus 2 unpaired electrons).
Element sheets, 4
Working individually or in teams, cut out the 32 element squares from the four sheets of elements.
Part 1. Elements
Part 2. Covalent Compounds
- Locate Data Table 1.
- Record the family or group number for each element listed in Data Table 1. Refer to a periodic table to do this step.
- List the number of valence electrons for each element in the next column of the Data Table 1.
- Locate the element square for the first element in the table and set the square on a flat surface. Use bingo chips to place the correct number of valence electrons around the element. Follow the rules for writing Lewis Dot Structures outlined in the background section of this lab handout. This is the Lewis Electron Dot Structure for that element.
- Draw the Lewis Electron Dot Structure in the last column of Data Table 1.
- Repeat steps 4 and 5 for each element listed in Data Table 1.
- Answer Questions 1–4.
- Locate Data Table 2.
- Determine the number of valence electrons for each individual atom in the first compound. Write each separate number in Data Table 2. Refer to Example 1 in Data Table 2.
- Add up the valence electrons for each atom and record the total number of valence electrons for the compound in the next column of the Data Table 2.
- Locate element squares for each atom in the compound and set the squares on a flat surface. Gather enough colored bingo chips to equal the total number of valence electrons. Use a different color for each element in the compound.
- Use the bingo chips to place the correct number of valence electrons in the compound. Follow the rules for atom placement and electron placement described in the background section of this lab handout.
- Once the structure is made and all rules have been followed, draw the Lewis Electron Dot Structure in the last column of the data table.
- Repeat steps 8–12 for each compound listed in Data Table 2.