Covalent Compounds: (Hybridization)

This section will discuss compounds that contain 3 atoms in a compound that will be applicable to compounds with hundreds of atoms. When considering such a wide range of compounds there are some features common to many of them. These basic features are due to molecular orbitals.

When atoms of different elements react to form chemical bonds and new molecules, the electron atomic orbitals (s,p,d,f) of the atoms of one element and the electron atomic orbitals (s,p,d,f) of the atoms of the other element will influence each other due to the fact that their negative charges are repelling each other (like charges repel). These interactions result in the formation of molecular orbitals which are modifications of the atomic orbitals of the elements that react. Molecular orbitals have shapes and orientations characteristic of the molecule of the compound.

An example to introduce and illustrate molecular orbitals is the compound methane, CH4. Methane is a gas that occurs naturally on earth and is found in petroleum deposits, landfills, among other places.

Molecules of methane are symmetrical with regard to atom placement and charge placement. The C and H atoms in methane are arranged as a tetrahedron with C at the center of the molecule and the bond angle between each H-C-H is 109.5 degrees).  (See Tetrahedral Molecular Geometry)

Starting with atomic orbitals of H and C, H has a 1s valence electron which is spherical and C has 2s22p2 valence electrons. Open a second window in your browser. Copy and paste http://www.ptable.com/# in the address bar. Click on the “Orbital” tab and then click on C.

Now find letters s, p, d and f with the stacks of small squares on the right.

Move the pointer or arrow over the small rectangle box labeled 2s and click. Note the spherical shape that appears to the right. Next move the pointer over each of the 2p orbitals. Note the p orbitals are 90 degrees relative to each other. In molecules of CH4, if the four hydrogen atoms with spherical orbital electrons bond to the 2p 90 degree orbitals of the C atom and there is  no modification, the bond angles of the CH4 molecule would be 90 degree bond angles and not 109 degrees found in these molecules .

Hybridization
To explain bonding in molecules, the concept of hybridization was proposed. As mentioned in the introductory paragraph, there are numerous interactions between valence electrons of the H and C atoms. The rational given is that these interactions force a modification of the C orbitals to minimize the interactions resulting in lower energy states. The resultant modification of the C orbitals is the promotion of one 2s orbital electron to the empty 2p orbital to form four (4) hybrid sp3 molecular orbitals. The designation sp3 is a shorthand way to show that one 2s and three 2p orbitals combine to form four (4) hybrid sp3 molecular orbitals.

The concept of hybrids is common in botany where 2 different plants are “crossed” to yield a plant with characteristics of both plants.

In this discussion, the three p orbitals with 90 angles are “crossed” with one s orbital with a 360 degree angle (sphere). Since three of the combined orbitals are 90 degrees and 1 is 360 degrees, the result should be more like the p orbitals (3/4 p orbitals versus ¼ s orbital) with a wider angle. Hence 109 degree bond angles in methane.

It is important the student understand (or at least have a working knowledge of) hybridization because hybridization provides an explanation for the shapes and properties of compounds. To begin building your understanding of hybridization, comparing molecules of methane (CH4), water (H2O) and ammonia (NH3) is a good starting point.

Data for molecules of the three compounds is shown in the following table.

  CH4 NH3 H20
Molecular weight ( g/mol) 16 17 18
Hybridization sp3 sp3 sp3
Bond angle (deg) 109.0 107.3 104.4
Shape Tetraheral Pyramid Bent
Boiling pt (C at 760 mm HG) -161.7 -33.34 100

The molecular weights of the three compounds are essentially the same. The hybridization of the center atom in each molecule is sp3. The bond angle shapes are different. There is a very big difference in the boiling point of the molecules of the compounds. How can these differences be explained?

We start by comparing four representations of the three molecules.

Lewis Structures (diagrams)

CH4 NH3 H2O
methanelewisdot ammonialewisdot waterlewisdot

2D structures

CH4 NH3 H2O
 methanelines  ammonialines  waterlines

 

Model kit representation

CH4 NH3 H2O
 methanemodelred ammoniamodelred watermodelred

Dipole representation

CH4 NH3 H2O
 methanetetra  ammoniawithdelta  waterwithdelta

 

Starting with the Lewis dot structure, each molecule (methane, ammonia, and water) has an octet (4 pairs) of valence electrons around the central atom with sp3 hybridization. Comparing the three Lewis dot structures, methane has all four of the electron pairs bonded to H atoms. Ammonia has three of the electron pairs bonded to H atoms. Water has only two electron pairs bonded to H atoms.

Comparing the 2D representations, only the bonded electrons are shown as lines. Each line represents a bond consisting of a pair of electrons shared between the two atoms. In this representation, the non-bonded electron pairs are not shown for ammonia and water but the non-bonded electron pairs have a large impact on the shape and the properties of ammonia and water. The effect of the non-bonded electrons provides the answer asked above, “How can these differences be explained?” This representation shows the ratio of elements in a molecule of the compound, but has no value regarding the molecular shape of the molecule

Comparing the Model kit representations, methane molecules have four pairs of electrons bonded to H atoms and have attained their optimum arrangement of plus and minus charge interaction (lowest energy configuration) to form 109 degree bond angles. The four bonded H atoms are identical and the methane molecule is symmetrical.

In ammonia molecules, the one lone pair of electrons will repel the three bonded pairs and the bond angle will be lowered from 109 degrees to 107 degrees. The oval represents a pair of unbonded electrons.

In water molecules, the two lone pairs (represented by the two blue ovals) will have even more of a repulsive effect on the two bonded pairs and the bond angle will be lowered from 109 degrees to 104 degrees.

Comparing the dipole representations, charge is distributed equally over methane molecules which produces a nonpolar molecule. Nonpolar compounds have symmetrical molecules. Because the charge is evenly distributed, a molecule of methane has very little attraction for another methane molecule. The attraction between methane molecules is so weak that methane is a gas at room temperature and it has a very low boiling point (-161.7 C ).

Ammonia molecules have a permanent dipole with the lone pair of electrons being the location of negative charge and the H end of the molecule being the location of positive charge, a polar molecule. The hydrogens will be more positive because the N-H bond is between two atoms in which the N atom has a greater electronegativity than H. As discussed in Covlant compounds: Basics, in a covalent bond the atom with the greater electronegativity will have a greater attraction for the bonded electrons leaving the other atom more positive. Dipole representations are an attempt to show location of charges by placing a delta symbol (delta20w20h  ) symbol with either a + or – sign. Since each ammonia molecule is a dipole (polar molecule), the negative end of one ammonia molecule will be attracted to the positive end of another ammonia molecule. This type of attraction between molecules is referred to as intermolecular attraction.  Ammonia molecules have a larger attraction for each other as reflected in a much higher boiling point for ammonia (-33.34 C). The attraction is not very strong since the boiling point is still low compared to room temperature.

In water molecules, the two lone pairs will have even more of a repulsive effect on the two bonded pairs and the bond angle will be lowered from 109 degrees to 104 degrees. Water molecules have a permanent dipole due to the location of the two lone pairs at one end of the molecule (negative charge) and positive charge located at the end of the molecule where the hydrogens are located, a polar molecule. See dipole representation. Since each water molecule is a dipole, the negative end of one water molecule will be attracted to the positive end of another water molecule to produce intermolecular attraction. Water molecules have a very high attraction for each other as reflected in its boiling point of 100 C.

The best way to display the three dimensional structure of molecules is to use computer modeling programs that take into account the sizes and charges of atoms to generate the molecules. 3D representations of methane, ammonia, and water are available  at Methane Ammonia Water.

As you do note the differences listed below.  Use the mouse to move the molecule of each example to see the shape of each molecule.

Notice the following:
1. Methane has a tetrahedral shape.
2. Ammonia has a pyramid shape
3. Water has a bent shape

In the bottom right of each molecule display is the text “JSMole”.  Click on the text to open up a menu.  Scroll down to the Surface item and click on it.  For methane select Molecular Surface.  For ammonia and water select Molecular Electrostatic Potential (range 0.1 0.1)

Items to note: (Again use your mouse to rotate the model)
Methane has a surface that is the same

Ammonia: The N atom has a red surface by it indicating this space is negatively charge.  The hydrogens have a blue surface indicating this space is positively charged.

Water: The O atom has a red surface by it indicating this space is negatively charge.  The hydrogens have a blue surface indicating this space is positively charged.

This separation of charge within the molecules of ammonia and water is the dipole mentioned above and is responsible for these the molecules of ammonia and water being polar compounds.

The 3D perspective of molecules is the closest to the actual shape and location of charge. When molecules are discussed in the remainder of this course, the student should get in the practice of thinking of them in their 3D shape. This visual will help as we discuss interactions and reactions in the subsequent sections.

Summary and future application
The discussion above has shown methane, ammonia, and water molecules have an octet of electrons around the central atom of the molecules. This octet of electrons has sp3 hybridization. Methane molecules have four identical atoms bonded to the central atom resulting in a symmetrical (think non-polar) molecule with no charge separation and very little intermolecular attraction. Hence a very low boiling point.

Ammonia molecules have three identical atoms bonded to the central atom and a pair of electrons not bonded resulting in a nonsymmetrical (think polar) molecule with charge separation (dipole) and a larger intermolecular attraction. Hence a low boiling point.

Water molecules have two identical atoms bonded to the central atom and two pairs of electrons not bonded resulting in a nonsymmetrical (think polar) molecule with charge separation (dipole) and a much larger intermolecular attraction. Hence a high boiling point.

To emphasize the chemistry of nonpolar and polar molecules consider the following information. The amount of methane that will dissolve in water at 20o C is 0.024 g/1 kg of water. The amount of ammonia that will dissolve in water at 20o C is 520 g/kg water. Why?
Water molecules are intermolecular bonded (See Amination of intermolecular bonding) due to the dipole (polar) property of each water molecule. Ammonia molecules have dipoles (polar) and will interact with water molecules. An image is that the polar ammonia molecules will be attracted to water molecules and interrupt the intermolecular bonding between water molecules.

Methane molecules are nonpolar (no dipole). Since they do not have a charge separation they cannot interact with water molecules and will not be taken into the water structure.
The basics introduced in this discussion can be applied to all molecules.

The sections of Interactions and Reactions will reinforce these concepts.