This week my students did the molecular model lab which is a bummer for them because we are one class period behind in lecture. So- if they didn't take the time to really read chapter 8 thoroughly before lab (most of them probably didn't) then the lab probably went in one ear and out the other.
I'm not sure if it makes more sense to try to explain ionic vs covalent bonding before or after Lewis structures/VSEPR shapes. By learning to use the Lewis Dot model students can predict shapes which influence the overall bonding characteristics of molecules. The larger concepts of ionic vs covalent bonding can be explained without knowledge of Lewis structures, however, the details seem much clearer once the model is explained.
I think people get bogged down in the Lewis model/VSEPR rules when they initially see these concepts. I'm wondering if a better explanation about overall bonding behavior might help them see the forest through the trees (forgive my cliche here) and help them understand the overall message that Lewis dot structures and VSEPR shapes try to convey. This is why I'm posting about ionic vs covalent bonding today.
Most people are told early on in chemistry that an ionic bond is that of a metal and a nonmetal and that electrons are transferred from one atom to the other to create charged species. We all know that opposite charges attract (+ charge attracts a - charge) so it makes sense that the ionic bond is held together primarily by the two individual charges of each ion respectively. The most common example is sodium (Na) and chlorine (Cl). Sodium sheds one electron to become Na+ while chlorine gains an electron to become Cl-. The ionic compound is then held together by electrostatic interactions between the positive and negative charges.
Covalent bonding is explained early on as the sharing of electrons. Elements like carbon don't really form charges. However, like a typical group 4 element, carbon has four valence electrons and bonds with 4 substituents. It forms a bond that doesn't involve the donation/acceptance of electrons. The carbon shares its 4 valence electrons with its substituents. (Somebody asked me once if we think of carbon as a partially positive +4 or as a partially negative -4 and I've done some reading on this- come see me if you are interested)
The issue of "shared" electrons in covalent bonding then brings up another issue in bonding: are the electrons shared equally? Does this influence any other characteristics of the molecule?
The answer to this question is that bonding models like the Lewis model along with charts of electronegativity (another trend of the periodic table) tell us a significant amount about the bonding behavior we can predict about certain compounds. This also influences whether a compound is a solid, liquid or gas at room temperature and other physical characteristics like solubility, boiling point, and melting point.
The concept of unequally shared electrons creates what we call polar molecules in chemistry. Polar molecules have a charge separation that is somewhat like ionic compounds but is mixed with a sharing, covalent characteristic of the bond. To determine if a covalent bond has polarity we look at the electronegativity of the atoms involved in the bond. Generally, the larger the difference in electronegativity, the greater the charge separation of the atoms and the more ionic characteristics the molecule obtains. You can think of a polar covalent bond as a "sticky" bond. We refer to it as having a "dipole moment" if there is a large separation of charge. And generally, like-minded dipole moments tend to affiliate themselves with like-minded dipole moments (sound strangely familiar to human behavior- eh?)
Electronegativity is the ability of an atom to hold electrons. Generally, electronegativity is greater for atoms on the right, top of the periodic table because these are small compact atoms that want to gain electrons to become like the group 8 noble gases. Fluorine is generally our best example of a highly electronegative atom. You can actually calculate the electronegativity difference between two atoms based on experimentally determined electronegativity values. (See figure in book). In general, I "eyeball" it based on where two atoms lie on the PT. The difference in EN between boron and fluorine, for example, would generally be larger than the difference between nitrogen and fluorine because the nitrogen is much closer in proximity to the fluorine than boron. The charge separation is larger and it has a greater dipole moment.
A great example is looking at the charge separation of molecules that dissolve in water. Water is a highly polar molecule. The electronegativity difference of the oxygen and hydrogen is large and creates a charge separation between the two atoms. Water has a fairly large dipole moment and is a polar molecule. Because of its overall polarity, water tends to dissolve compounds/molecules that have dipole moments and are also considered polar molecules. This is the reason that oil does not dissolve in water (as you can observe in your own kitchen); oil does not have a charge separation associated with it. The long hydrocarbon tails are largely nonpolar covalent bonds and won't complex with the partial + and partial - characteristics of the H2O bond.
Aside from electronegativity difference between two atoms, the other component that determines the overall polar characteristics of a molecule is symmetry. It is possible that a molecule could have polar bonds and yet exhibit no overall dipole moment. In this case, despite its polar bonds, it would not dissolve in water.
Let's look at the case of carbon tetrafluoride. There is definitely a charge difference between the carbon and the fluorine molecule. Each individual bond has a dipole moment. The tetrahedral structure, however, is a symmetrical structure. Therefore, the individual dipole moments cancel each other out and the molecule does not exhibit overall polarity. I would not expect CF4 to dissolve in water. It is considered a nonpolar covalent molecule overall because of the symmetry of the molecule.
This is where the importance of lewis dot structures enter the picture. It is only with an accurate picture of bonding (single, double, triple) with a Lewis dot structure that the correct shape of the molecule can be determined. It is only with this shape that symmetry can be determined. The overall symmetry of the molecule dictates whether or not the molecule contains polarity.
What about the case of CHF3. Is it the same as CF4 in that the symmetry of the molecule cancels the dipole moments created by each C-F bond? The answer is no because in this case the molecule does not exhibit overall symmetry. The H substituent is not equal to the 3 F substituents in terms of the dipole moment it creates. For this reason, the molecule exhibits an overall dipole moment and is considered a polar molecule.
