Notice that a Lewis representation does not differentiate between the sigma and the pi bonds in the so-called “double bond.” It simply shows the two together as two equal dashes. In this case a pure line-angle formula for ethene would look awkward because it would resemble an equal sign (=). Notice that although C–H bonds are not usually shown in line-angle formulas, sometimes they are included for enhanced clarity. Below is a Lewis and a line-angle representation of ethene, which is sometimes informally called ethylene. By this definition, the simplest possible alkene must contain two carbon atoms. The principles of all this chemistry will be discussed later in the course.ĪLKENES ARE HYDROCARBONS THAT CONTAIN AT LEAST ONE PI BOND AS PART OF THEIR MOLECULAR STRUCTURE. It is relatively easy to break a pi bond compared to the sigma bond. Under certain conditions, they have the capability to become DELOCALIZED, that is to say, they can move in the molecular skeleton from one atom to another, or even become spread over several atoms, according to principles we’ll study later.Īt the same time, in chemical reactions where electrons are to be traded, the pi electrons are more readily available because they are more exposed and less tightly bound by the nucleus. The electrons in the pi bond (or pi electrons) are less tightly bound by the nucleus, and therefore they are relatively mobile. The electrons in the sigma bond (or sigma electrons) are more tightly bound to the nucleus and don’t move too much. This has some implications in the properties and chemical reactivity of sigma and pi bonds. Sideways overlap is less efficient than head to head overlap and results in formation of weaker bonds. The pi bond, on the other hand, is relatively long and diffuse. As a rule, head to head overlap is the most efficient way to bond and results in relatively strong and stable bonds. The illustration above tries to convey a basic feature of the pi bond as compared to the sigma bond. The bond formed by the sp 2 orbitals is a sigma bond, and the bond formed by the p orbitals is called a pi bond. When two sp 2 hybridized carbon atoms approach each other to bond, two sp 2 orbitals approach each other head to head, and two p orbitals approach each other sideways.

To see this arrangement clearly, we must switch to a side view of the orbital system. That is to say, it is positioned at right angles to those orbitals, with one lobe coming out of the plane of the page and the other going behind the page. In this top view, the unhybridized p orbital cannot be seen because it also arranges itself to be as far apart from the sp 2 orbitals as possible.

A top view of this arrangement is shown below. The ideal angle between sp 2 orbitals is therefore 120 o. That is to say, the carbon nucleus will be at the center of an equilateral triangle, and the three sp 2 orbitals will point to the corners of that triangle. Therefore, the three equivalent sp 2 orbitals will arrange themselves in a trigonal planar configuration. Again, according to VSEPR theory, equivalent orbitals will arrange themselves in 3-D space to be as far apart from each other as possible. It still retains its original energy and shape. The process is shown below.Īs shown, the three resulting sp 2 orbitals are equivalent in energy, but the remaining p orbital has not been affected. We will now reproduce the sp 3 hybridization process for carbon, but instead of taking one s and three p orbitals to make four equivalent sp 3 orbitals, this time we’ll take only one s and two p orbitals to make three equivalent sp 2 orbitals, leaving one p orbital untouched.

C2f6 sigma and pi bonds