Electronegativity and Polarity
One can summarize organic chemistry by stating that it is the chemistry of electrons. This is because electrons are the principle components of bonds and elements can be characterized by their electronic behavior. Thus, a fundamental understanding of electrons and their environment is needed before we plunge into the breadth of organic chemistry. This material may seem like a review if you’ve recently taken general chemistry, but it is crucial that you understand the concepts within this section before moving on.
Remember that the periodic table is arranged via columns and rows and that elements in the same row are similar in size and elements within a column have similar electronic behavior. This leads to “typical” behaviors exhibited by elements within the same column (AKA periodic trends).
One of these periodic trends is the ability of an element to attract electrons toward itself. This concept is known as electronegativity. There are numerous ways to quantify electronegativity values. Below is an arbitrary system created by Linus Pauling in 1932. These are most likely the values you would see in an undergraduate class and/or textbook.￼
Electronegativity increases across a row of the periodic table as the number of protons increases. This is known as nuclear charge. It is the positive attraction of the nucleus for negatively charged electrons. The question is, "Why does this increase as one travels up and toward the right of the periodic table?" The answer to this question is simple, yet two-fold. It is best understood if you also understand why electronegativity values decrease while travelling down and to the left of the periodic table.
Electronegativity values decrease down a column of the periodic table as atoms increase in size (AKA atomic radius). Atomic radius is related to electronegativity because valence electrons are further from the nucleus in larger atoms, thus the nuclear charge is not fully experienced. Think of this concept as you would of two magnets: When two magnets are far apart, they do not attract each other as greatly as they do when they are closer to one another. Thus, one magnet is the positively charged nucleus and the other is an electron. And, of course, the opposite is true. If an atom is small, the nucleus will have a strong effect on an electron (a strong nuclear charge will be "felt" by the electron) simply because they are closer together.
As you can see, electronegativity increases up and to the right along the periodic table. Thus, the higher the associated electronegativity value, the more an element attracts electrons toward itself. This results in Fluorine (F) being the most electronegative atom, and Francium (Fr) the least electronegative atom (or most Electropositive atom). Consequently, bonds between two atoms that differ greatly in their respective electronegativities are more polar than those between two atoms with similar electronegativities. This has major consequences in the study of organic chemistry.
Electronegativity values are commonly used to indicate whether the electrons are being shared equally or unequally within a bond. For example, in the HF molecule, the bonding electrons spend more time on the fluorine atom than the hydrogen atom because fluorine is more electronegative than hydrogen. We indicate this slight excess of electrons with a partial negative charge on the fluorine with the symbol δ- (remember that this is not a full charge, as in formal charges, but a partial charge—we will discuss this soon):
As a result of this unequal sharing of bonding electrons the HF molecule has an electric dipole moment. The electric dipole moment is a vector quantity, is represented by the symbol μ, and is measured in units called Debyes. (Please note that most of this technical information is unnecessary to understand the concepts of electronegativity and bond polarity and is only provided to maintain consistency with the underlying science.) In the case of HF, μ will lie along the direction of the H-F bond as such:
￼A difference of < 0.5 is considered to be a nonpolar covalent bond (equal sharing). A C-C or C-H bond, for example. A difference of 0.5-1.9 is considered to be a polar covalent bond (unequal sharing). A C-Cl bond, for example. A difference of > 1.9 is considered to be an ionic bond (e- transfers). Na+Cl-, for example.
By combining these two figures, we arrive at:
We will explore these three bonding categories more in depth in the next section.) Remember from before that the most electronegative atom in a polar covalent bond is designated the delta minus (δ-) whereas the least electronegative atom is designated the delta plus (δ+). We use these designations to identify where the electrons within the bond spend most of their time. Also remember that these are partial charges, they are not full charges, as in ionic compounds. Up until this point we have only concerned ourselves with the polarity of one bond (HF). However, what happens if a molecule has more than one polar bond? To determine if a molecule has a net dipole moment, a two-step process must be used:
Use electronegativity values to calculate the differences between atoms and identify all polar bonds and the directions of the dipole moments. Determine the spatial geometry of an individual atom (easiest to do by determining the hybridization of the atoms and thus their geometry) and determine if the dipoles cancel or reinforce each other.
As you can imagine, these concepts (electronegativity and polarity) are essential to understanding organic chemistry. Bond polarity (which results from differences in electronegativities) is one of the most fundamental concepts of organic chemistry and having a firm grasp of it is crucial to your success with this subject. We will explore the consequences of bond polarity later in our discussions of reactions. For now, know that polarity is one of the reasons molecules undergo reactions at all.