Induction and Polar Covalent Bonds

Chemists classify bonds into three categories: (1) covalent, (2) polar covalent, and (3) ionic. These categories emerge from the electronegativity values of the atoms sharing a bond. Electronegativity  is a measure of the ability of an atom to attract electrons. Table 1.1 gives the electronegativity values for elements commonly encountered in organic chemistry.

When two atoms form a bond, there is one critical question that allows us to classify the bond: What is the difference in the electronegativity values of the two atoms? Below are some rough guidelines:

If the difference in electronegativity is less than 0.5 , the electrons are considered to be equally shared between the two atoms, resulting in a covalent bond. Examples include C–C and C–H:

The C–C bond is clearly covalent, because there is no difference in electronegativity between the two atoms forming the bond. Even a C–H bond is considered to be covalent, because the difference in electronegativity between C and H is less than 0.5. If the difference in electronegativity is between 0.5 and 1.7,  the electrons are not shared equally between the atoms, resulting in a polar covalent bond . For example, consider a bond between carbon and oxygen (C–O). Oxygen is significantly more electronegative (3.5) than carbon (2.5), and therefore oxygen will more strongly attract the electrons of the bond.

The withdrawal of electrons toward oxygen is called induction , which is often indicated with an arrow like this:

Induction causes the formation of partial positive and partial negative charges, symbolized by the Greek symbol delta  . The partial charges that result from induction will be very important in upcoming chapters.

If the difference in electronegativity is greater than 1.7, the electrons are not shared at all. For example, consider the bond between sodium and oxygen in sodium hydroxide (NaOH):

The difference in electronegativity between O and Na is so great that both electrons of the bond are possessed solely by the oxygen atom, rendering the oxygen negatively charged and the sodium positively charged. The bond between the oxygen and sodium, called an ionic bond , is the result of the force of attraction between the two oppositely charged ions. The cutoff numbers (0.5 and 1.7) should be thought of as rough guidelines. Rather than viewing them as absolute, we must view the various types of bonds as belonging to a spectrum without clear cutoffs.

This spectrum has two extremes: covalent bonds on the left and ionic bonds on the right. Between these two extremes are the polar covalent bonds. Some bonds fit clearly into one category, such as C–C bonds (covalent), C–O bonds (polar covalent), or Na–O bonds (ionic).

However, there are many cases that are not so clear-cut. For example, a C–Li bond has a difference in electronegativity of 1.5, and this bond is often drawn either as polar covalent or as ionic. Both drawings are acceptable.

Another reason to avoid absolute cutoff numbers when comparing electronegativity values is that the electronegativity values shown above are obtained via a method developed by Linus Pauling. However, there are at least seven other methods for calculating electronegativity values, each of which provides slightly different values. Strict adherence to the Pauling scale would suggest that C–Br and C–I bonds are covalent, but these bonds will be treated as polar covalent throughout this course.

Consider the structure of methanol. Identify all polar covalent bonds and show any partial charges that result from inductive effects.

First identify all polar covalent bonds. The C􀀗H bonds are considered to be covalent because the electronegativity values for C and H are fairly close. It is true that carbon is more electronegative than hydrogen, and therefore, there is a small inductive effect for each C–H bond. However, we will generally consider this effect to be negligible for C–H bonds. The C–O bond and the O–H bond are both polar covalent bonds:

Now determine the direction of the inductive effects. Oxygen is more electronegative than C or H, so the inductive effects are shown like this:

These inductive effects dictate the locations of the partial charges:

 

Identifying Formal Charges

A formal charge  is associated with any atom that does not exhibit the appropriate number of valence electrons. When such an atom is present in a Lewis structure, the formal charge must be drawn. Identifying a formal charge requires two discrete tasks:

1.  Determine the appropriate number of valence electrons for an atom.

2.  Determine whether the atom exhibits the appropriate number of electrons.

The first task can be accomplished by inspecting the periodic table. As mentioned earlier, the group number indicates the appropriate number of valence electrons for each atom. For example, carbon is in group 4A and therefore has four valence electrons. Oxygen is in group 6A and has six valence electrons.

After identifying the appropriate number of electrons for each atom in a Lewis structure, the next task is to determine if any of the atoms in the Lewis structure exhibit an unexpected number of electrons. For example, consider this structure:

 

Remember that each bond represents two shared electrons. We split each bond apart equally, and then count the number of electrons on each atom:

Each hydrogen atom exhibits one valence electron, as expected. The carbon atom also exhibits the appropriate number of valence electrons (four), but the oxygen atom does not. The oxygen atom in this structure exhibits seven valence electrons, but it should only have six. In this case, the oxygen atom has one extra electron, and it must therefore bear a negative formal charge, which is indicated like this:

 

Consider the nitrogen atom in the structure below and determine if it has a formal charge:

We begin by determining the appropriate number of valence electrons for a nitrogen atom. Nitrogen is in group 5A of the periodic table, and it should therefore have five valence electrons. Next, we count how many valence electrons are exhibited by the nitrogen atom in this particular example:

In this case, the nitrogen atom exhibits only four valence electrons. It is missing one electron, so it must bear a positive charge, which is shown like this:

Identify any formal charges in the structures below:

 

 

 

ATOM

—Everything in our world is built of atoms, from the smallest piece of paper to the  biggest , most complicated device.— There are many types of atoms , there are even some artificial man made atoms. Each atom is unique and has a different mass, size and properties in  comparison to other atoms.

Elements are made out of atoms. After many years of research, scientist have found that there are 92 natural elements in the world. As mentioned before , there are some synthetic elements, over 20. however they are too unstable to occur naturally on earth.