Unit 3: Chemical Bonding

Section 2: Molecular Structures

Introduction

Electron Dot Structures

The Octet Rule

Exceptions to the Octet Rule

Bond Energy

Introduction

Have you tried picturing the chemical bonds between water moleucles? This can prove to be difficult. Therefore, scientists have devised ways to represent chemical bonds on paper. In this section, we will examine some common structural models that scientists use to describe chemical bonds. These models can help to better visualize chemical bonds.

Electron Dot Structures

An American chemist, G.N. Lewis, developed a system used to describe chemical bonds in 1916. This system is called electron dot structures or Lewis structures. In this system, the symbol of the element represents the nucleus of the atom. Dots are then placed around the the symbol. The dots represents the valence electrons (the outermost electrons in the atom that are used for chemical bonding) of an element. However, only two dots may be placed on each of the four sides of the symbol (top, left, bottom, right). This represents the electron pairs. A maximum of eight dots can fit one symbol. Each element in a family have the same number of dots around it. For example, sodium (Na) has one electron around it, potassium (K) has one valence electron (one dot), and the same holds for the remaining elements in the family.

The Octet Rule

When forming a compound, a metal would like to lose electrons, and a nonmetal would like to gain electrons. Why is this? Looking at the hydrogen electron dot structure, we see that hydrogen (H) has one dot (which side the dot is placed has no significance). When hydrogen combines with fluorine (F) it forms hydrogen fluoride. Fluorine (F) has seven dots in its electron dot structure. Since the compound is an polar molecule, we also know that H will tend to lose its electron, and F will pull on it stronger. According to the octet rule, nonmetals tend have eight electrons in a chemical bond. This is due to the fact that nonmetals try to have eight valence electrons like noble gases (eight dots) which are the most stable elements. To represent HF in terms of Lewis dot structures, we would write H with a dot on the right side, and F with a dot on the left side, two dots on the top side, two dots on the right side, and two dots on the bottom. The common link between the right side dot on the H atom and the left side dot of the F atom signifies the chemical bond. The octet rule tells us that this compound is possible. However, the octet rule does not accurately predict every stable configuration of all molecules and compounds.

Exceptions to the Octet Rule

The Lewis dot structure can serve as the underlying concept in understanding the chemical bonds in molecules. However, the octet rule, which governs the writing of lewis dot structure, has expections.

Not every nonmetal, nor metal, can form compounds in which each element satisifies the octet rule. For example in BF₃, boron has only three valence electrons which can bind to three fluorine atoms. As a result of covalent bonding, boron shares three electron pairs with fluorine. Furthermore, boron ends up with six valence electrons and not with eight electrons, as stated by the octet rule. Boron is an example as to why the octet rule does not apply to all metals. As for the nonmetals, exceptions also exist. In SF₆, sulfur forms six bonds with fluorine, resulting in 12 shared electrons (not eight). Experimental data shows that BF₃ and SF₆ form stable molecules. Thus the octet rule should not be the only rule in determining the nature of bonds and shapes of molecules and compounds.

Bond Energy

Bonds formed between elements have energy. The energy required to break these bonds is called bond energy. For strong bonds, more energy is required, and in weak bonds, less energy is required. Similarly, the energy released when a strong bond is broken is more than the energy released by a weak bond. To measure the bond energy between atoms and molecules, scientists rely on greater quantities than single atoms or molecules. Instead, they use 6.02*10²³ molecules or atoms. This number is very important in chemistry, and it represents one mole of a substance. It would prove much easier to measure 6.02*10²³ atoms or molecules than one single small atom or molecule!

To break a bond, energy must be put into the system. Thus, an endothermic reaction is required to measure the energy of a chemical bond. Experimental data has shown that it takes 436 kJ (in chemistry, the unit for energy most commonly used is the joule, abbreviated "J") to break one mole of H₂ molecules into H atoms. Thus the bond energy for H₂ is 436 kJ/mol. This is the energy of one mole of H₂ molecules. Bonds must also be broken to form molecules and compounds. If a reaction involves a molecule breaking up to form a new compound, the bonds that are broken release energy. However, energy is also required to form the bonds in the new compound.

In a molecule where a triple bond, double bond, or single bond is formed, the triple bond would be the strongest. As in the bonds between to two carbon atoms, a triple bond is the strongest of the possible single, double, and triple bonds that could be formed. A triple bond between two carbon atoms also satisify the octet rule. Experimental data has also shown that the atoms are closer together in a triple bond than in a single or double bond. In experiments with a compound containing a single, double, or triple carbon bond (i.e. C₂H₆, C₂H₄, and C₂H₂), the compound containing a double carbon bond (C₂H₄) was more reactive than the compound containing a single carbon bond (C₂H₆), and the triple carbon bond compound (C₂H₂) was more reactive than the double carbon bond. The compound with the triple bond released the most energy (highest bond energy) when reacting with a different compound because bonds were broken when the reaction was taking place; the stored energy of the triple carbon bond was more than a single or double carbon bond.