|Chapter Three, Section One|
I. Lewis Bonding
In 1916 G. N. Lewis proposed a very simple theory to explain some stability of atoms, ions, and molecules. This theory is still a good starting place for chemistry students today. He noted that the Noble gases (formerly called the inert gases because people did not think that they would react with anything) all were very stable atoms. Also he noted that the ions tended to form a Noble gas electronic structure. We write the electronic structure for He (1s2), Ne (1s22s22p6), Ar (1s22s22p63s23p6) and we see that all of the subshells are filled. We write the electronic structure for some ions such as Na+ (1s22s22p6) and we see that this is the same electronic structure as Ne. Similarly for some other ions. You see that the outer electrons all have all 8 subshell filled so this principle of stability is called the octet rule. The transition elements do not obey this octet rule. Table 3.1 lists some common ions that we will be using. You need to memorize these ions. Ions are very different species than the parent atom as discussed in your book.
Bonds can be formed in two different ways:
Electronegativity is a measure of the ability of an atom to attract electrons to itself when it is bonded to some other atom. Pauling set up an electronegativity scale as seen in your text. The greater the difference in electronegativity between two atoms, the more ionic the bond. If the difference is greater than 1.9, the bond is called ionic. You do not have to memorize this table! However you should be a bit familiar with it, knowing that electronegativity increases from left to right and from bottom to top so that you could predict without a table which of two atoms would have the greater electronegativity.
II. Ionic Bonds
The difference in electronegativity between Na and Cl is 2.6 so we know that the bond between the two atoms will be ionic. The sodium atom will lose an electron to form the Na+ ion (note the Noble gas structure) and the Cl atom will gain the electron to form the Cl- ion (note again the Noble gas structure). Now a lot of these ions will attract each other to form the NaCl crystal (remember no such thing as a single NaCl molecule). In the crystal each Na ion is surrounded by 6 Cl ions and each Cl ion is surrounded by 6 Na ions. The structure is thus electrically neutral. Even though there is no such thing as a NaCl molecule you see that the ratio (1-1) of ions is the true lowest whole number ratio of the ions involved in the bonding. Ionic compounds such as BaCl2 have the ratio (1-2) of ions involved in the bonding. Below is a depiction of the crystal structure for the ionic compound NaCl:
III. Covalent Bonds
If the electronegativity difference is <1.9 then the bond is covalent. What is the EN difference between the two H atoms? (that is a tough one!) The answer is zero so the bond between the H atoms is definitely covalent. When different atoms are involved in the bonding the EN difference may be <1.9 but close enough to 1.9 so that the bond is really also a bit polar. Then we further divide the covalent bonding categorization into polar covalent and nonpolar covalent. We will make the following arbitrary distinction:
|EN Difference||Bond Type|
|0.5 to 1.9||polar covalent|
You need to memorize this table and be able to state the type of bonding between any pairs of atoms. The atom in a polar covalent bond with the greater EN will have the more negative charge. We indicate the charge with a Greek letter delta with delta+ for the atom with the lesser EN and delta- for the atom with the greater EN.
The ability to draw a Lewis structure for a particular molecule is very important for it gives us more of an understanding of how the molecule is constructed and, as we will see, it helps us to predict the three-dimensional geometry of the molecule. A Lewis structure shows the atoms of a molecule in their proper orientation to each other and it shows the arrangement of valence electrons around the molecule. There are some simple rules that will help you to draw the Lewis structure of some simple molecules. In drawing such molecules we need to consider some simple rules such as a hydrogen atom can have only one single bond to it, a carbon atom can have a total of only 4 bonds, an oxygen atom can have only two bonds, halogens can only have one bond, and nitrogen can have only 3 bonds. When you have several atoms involved, it is usually the atom with the lowest electronegativity which goes in the center. These rules will allow you to construct a lot of molecules.
For example, consider HCN. Since the hydrogen can have only one bond it will have to go on an end. The carbon atom has a lower electronegativity than N so it goes in the middle (H-C///N, where /// means a triple bond). Each bond represents 2 electrons, so a single bond (-) represents two electrons, a double bond(=) represents 4 electrons, and a triple bond (///) represents 6 electrons. Below is a better representation of HCN:
We can determine how many electrons we have to distribute about a molecule by counting the total number of valence electrons in the molecule. For HCN the H contributes 1, the C contributes 4, and the N contributes 5 for a total of 10 valence electrons. From the HCN bond diagram above, we count 8 electrons (two from the H-C bond and 6 from the C///N bond). That means that we have 2 electrons left over after the bonding is complete. What do we do with them? We put them with the most elecronegative atom, which is nitrogen. These extra electrons are unpaired and are called nonbonding electrons, for obvious reasons.
Sometimes we can draw more than one type of Lewis structure for the molecule. Such a molecule is SO2. We can satisfy the octet rule in two different ways with SO2 as indicated below:
We call each of these possibilities a "resonance structure" and usually place a double arrow between each. But which is correct? The answer is neither for this is just a model. Actually experiment tells us that each bond is exactly the same type, so they are not single and double bonds. A better way to write the structure is indicated below where you see that each bond is now the indicated to be the same and is somewhere between a single and a double bond:
Now you need to practice writing some Lewis structures.
After you have studied this material and practiced some problems, take quiz one. Note that there are no Lewis structure questions due to the difficulty of answering and stating such questions in this format. Be sure that you understand how to draw Lewis structures for simple molecules. If you score at least 80 on the test then you are ready to continue to the next section.
Web Author: Dr. Leon L. Combs
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