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Home  / GENERAL CHEMISTRY Textbook / Chapter 4. Molecules

Chapter 4. Molecules

We know that molecules are made up of atoms. As we spend energy to break a molecule into atoms, (i.e., we heat the molecule to a temperature of 2,000° to 5,000° C) we say that the atoms are bonded into molecules. The bonds with which atoms are thus connected to each other to form molecules are known as chemical bonds.

Contemporary model of chemical bonding

Since atoms are made up of negatively charged electrons and positively charged nuclei, it is natural to suppose that chemical bonding occurred at the expense of the attraction of the negatively charged electrons of one atom to the positively charged nuclei of another atom.  

One of the questions that will undoubtedly help to understand the process of chemical bonding is as follows:

How many electrons take part in the formation of a chemical bond? In the case of a hydrogen molecule, it is clear that in order to form a chemical bond, 2 electrons are sufficient, since each of the hydrogen atoms bonded into a H2 molecule, have but one electron.

All the other atoms (besides those of hydrogen) contain more than one electron.  

If our supposition about the formation of molecules at the expense of the attraction of one atom's nucleus to another atom's electrons is correct, it was not clear why the helium atom (He), having 2 electrons, does not form stable molecules of the He2 type. 

Studies on the composition of molecules, including hydrogen molecules and atoms of elements of the II period: Li, Be, B, C, N, O, F, and Ne, have shown that the number of hydrogen atoms that each of these elements can bond amounts to 1, 2, 3, 4, 3, 2, 1, 0, respectively. That is, the atoms of Li and F form stable molecules LiH and HF, while B and N form stable molecules BH3 and NH3. Ne does not form any stable molecules with hydrogen.

Conditions of chemical bond

As indicated in the previous section, the amount of electrons in the outermost shells of Li, Be, B, and C atoms comprises 1, 2, 3, and 4 respectively. That is, the number of hydrogen atoms that can bond to the given atoms is equal to the number of electrons in the outermost shells of these atoms.

In the case of molecule formation of the H2 and Cl2 type, both atoms entering the bonding process are equivalent. Two electrons take part in bond formation— one from each of the atoms to be bonded.

We can draw two conclusions from this data:

  1. Only the electrons in the outermost shell of the given atoms take part in bond formation.
  2. Only one electron of an atom is spent on bond formation in a hydrogen atom to form a single bond.

According to the second conclusion, the number of hydrogen atoms that an atom can bond to (in the case of Li, Be, B, and C) is equal to the number of electrons in the outermost shell of the central atom.

On the other hand, atoms of nitrogen (N), oxygen (O), and fluorine bond 3, 2, and 1 hydrogen atoms respectively, while the neon (Ne) atom does not bond to a hydrogen atom at all.

From the data concerning the structure of electronic shells, we know that the number of electrons in the outermost electronic shell of elements of the 2nd period (including N, O, F, and Ne) cannot exceed 8.

The previously cited FIEs indicate that after the increase in the number of electrons up to 8, the atoms of the 2nd period begin forming a new outermost shell.

A comparison of this data with that of the amount of hydrogen atoms (with only one electron) that can bond to N, O, and F atoms, allows us to make following conclusions:

  1. During bond formation of the N-H, O-H, F-H type [here the dash (-) indicates chemical bonding] the electron of the hydrogen atom enters the outermost shell of the central atom.
  2. The number of hydrogen atoms, which an atom of the 2nd period can bond to, is limited to the maximal amount of electrons that the outermost shell of such atoms can contain. According to the data on FIEs, this amount is equal to 8. 

According to these conclusions, Ne, which already has 8 electrons in the outermost shell, cannot form stable molecules of the NeH type. In reality such molecules do not exist.

Thus, experimental data on the FIEs, reveals their comparison with the chemical contents of stable molecules, and we can draw following conclusions:

  1. Only electrons situated in the outermost electronic shell of the atoms being bonded take part in the formation of chemical bonds. 
  2. Only one electron of the outermost shell offers the possibility for one bond formation. 
  3. Two electrons - one from each atom - take part in chemical bond formation between two atoms. These two electrons are bonding electrons.  
  4. After bond formation, both bonding electrons enter the outermost shells of the atoms to be bonded. Therefore, in the course of bond formation, the number of electrons in the outermost shell of the atoms to be bonded increases by one unit.  
  5. The amount of bonds that an atom can form is limited at the bottom by the amount of electrons present in the outermost shell of the given atom. For atoms of the 2nd and 3rd periods, this limitation is applied to atoms with fewer than 5 electrons in the outermost shell, i.e., to atoms Li, Na, Be, Mg, B, Al, C and Si.  
  6. The number of bonds that an atom of the 2nd and 3rd periods can form is limited at the top by the number of electrons that can be present in the outermost shells of these atoms. The maximal number of electrons that can be present in the outermost shells of atoms of the 2nd and 3rd periods is equal to 8 according to the data on the FIEs. This limitation is applied to atoms with more than 4 electrons in the outermost shell, i.e., to atoms N, P; O, S; F, Cl; Ne, Ar.

The data on IEs, FIEs, and the composition of stable molecules - their true values and comparisons - in the case of free atoms and atoms bonded into molecules, have allowed us to understand how atoms bond into molecules. 

Examples and Illustrations

Now let's take a few examples to see how we can make use of our rules to define the number of covalent bonds that an atom can form if we know the amount of electrons in the outermost shell of the given atom and its nuclear charge.

The nuclear charge and the amount of electrons in the outermost shell are defined experimentally and included in the table of elements which shows the nuclear charge values and the amount of electrons in the outermost electronic shells of atoms.

For example, let's calculate the number of covalent bonds that Na, Al, P, and Cl can form. Na and Al have 1 and 3 electrons in the outermost shell respectively, and according to our first rule — one electron in the outermost shell is used to form a covalent bond — they can form: Na - 1 and Al - 3 covalent bonds. After the bond formation, the number of electrons in the outermost shells of Na and Al is equal to 2 and 6 respectively; i.e., less than the maximum (8) for these atoms.

P and Cl have 5 and 7 electrons in the outermost shell respectively, and, according to the second of the enumerated regularities, they could have formed 5 and 7 covalent bonds.       

According to the 4th regularity, during the formation of a covalent bond, the number of electrons in the outermost shell of these atoms increases by 1. According to the 6th regularity, when the covalent bond forms, the number of electrons in the outermost shell of the atoms being bonded cannot be more than 8. That is, P can form only 3 bonds (8-5=3), while Cl can form only 1 (8-7 =1).

The described regularities for the formation of covalent bonds allow us to foretell theoretically the molecular structure of substances on the basis of an elementary analysis.

For example: on the basis of analyses, we have found that a certain substance consists of Na and Cl atoms. Knowing the regularities of covalent bond formation, we can say that Na can form only 1 covalent bond. That is, we can expect that every Na atom is bonded with a Cl atom via a covalent bond in this substance, and that the substance is composed of molecules of NaCl. The structural formula for this molecule is Na - Cl. Here the dash (-) indicates the covalent bond. The electronic formula of this molecule can be illustrated thus:  


. .
Na  : Cl

. .

According to the electronic formula the outermost shell of a Na atom in NaCl contains 2 electrons, while that of a Cl atom contains 8.

In the given formula, the electrons (dots) between the Na and Cl atoms are bonding atoms. Since the FIE of Cl is equal to 13 eV and that of Na is equal to 5.14 eV, the bonding pair of electrons is situated much closer to the Cl atom than to the Na atom.

Let's observe another case. On the basis of analyses, we have found that a substance is made up of Al atoms and Cl atoms. Al has 3 electrons in its outermost shell; thus, it can form 3 covalent bonds, while Cl, as in the previous case, can form only 1. This substance is represented as AlCl3 and its electronic formula can be illustrated thus:

. . . .
: Cl : Al : Cl :
. . . . . .
: Cl :
. .

whose structural formula is:

Cl —  Al —  Cl
 |
Cl

This electronic formula shows that there are 8 electrons in the outermost shell of the Cl atoms in AlCl3, while there are 6 in the outermost shell of the Al atom. 

On the basis of the comparison of the data on the FIEs with the data on the structure of stable molecules, we have found that during bond formation, both bonding electrons (one from each atom) enter the outermost shells of the atoms to be bonded. This conclusion allows us to imagine how a dual-atomic molecule actually looks. For example, the simplest dual-atomic molecule is a hydrogen molecule.

Hydrogen atoms are all perfectly identical, that is, the bonding pair of electrons is not only in the outermost shells of both atoms to be bonded, but is also at the same distance between the nuclei. Thus, there is only one definite place for the electrons — between the nuclei at equal distances from them.

Before the bonds were formed (i.e., in divided atoms) the electrons rotated around their nuclei, that is, when the nuclei are united, the electrons continue rotating. And, as indicated above, during their rotation both electrons are at the same distance from the hydrogen nuclei being bonded by them.

During bond formation, the bonding electrons enter the outermost shells of both atoms to be bonded.

This conclusion allows us to introduce a scheme for chemical bonding. According to this scheme, say, in the case of a hydrogen molecule, the bonding pair of electrons rotates on a plane perpendicular to the axis connecting the nuclei of the hydrogen atoms. The center of the electrons' rotation is situated at equal distances from the nuclei to be bonded. It is only in this case that the electrons can simultaneously enter the outermost electronic shells of both atoms to be bonded.

Thus, the hydrogen molecule can be illustrated as follows:

hydrogen molecule
Figure 4.1

A hydrogen molecule is composed of 2 hydrogen atom nuclei located at a distance of about 0.7Å from each other. Two electrons in the hydrogen molecule rotate in a circular orbit on a plane perpendicular to the axis connecting the nuclei. This electronic pair draws the nuclei together at the expense of the electrostatic forces and the nuclei's attraction forces to the two electrons rotating between them.

The center of the circle, around which the electrons rotate in a hydrogen molecule, is located in the middle of the axis connecting the hydrogen nuclei. That is, the electrons are at identical distances from the nuclei are bonded by the electrons.

When the same kind of atoms are bonded chemically (like atoms of Na-Na, F-F, Cl-Cl, etc.) the bonding electrons, just as in the case of a hydrogen molecule, are at the same distance from the nuclei of the atoms being bonded.

When various atoms are chemically bonded (like Na-Cl, etc.), the center of the circle of the bonding electrons is closer to the nucleus of the Cl atom. That is, the bonding electrons are shifted towards the atom that attracts the electrons more readily, i.e., towards the atom with a higher FIE.

The bonding electrons are at about the same distance from the nuclei being bonded by them, as are the non-bonding electrons. That is, when the bonding electrons form chemical bonds, they enter the outermost shells of the atoms being bonded.

The inner electrons are often called core electrons (see Fig. 4.2).

electrons in hydrogen molecule
Figure 4.2

The dots on the circles are the electrons. The smaller circles are the cores of the atoms. The core of the atom includes, besides the nucleus, all the inner electronic shells, i.e., all the electrons in the inner layers around the nucleus. That is, in the Cl atom this core includes a nucleus with 17 protons and 10 electrons (2 electrons in the layer closest to the nucleus and 8 electrons in the next closest layer).  

These inner electrons, like the atom's nucleus, do not take part in chemical transformations; that is why, for convenience, the drawing shows them united with the nucleus. Thus, the Cl core carries a surplus positive charge equal to 17-10 = 7 proton units. Seven electrons located in the outermost shell and indicated on the drawing with dots compensate this positive charge. They are called valence electrons since only these electrons in the atom take part in chemical bond formation and chemical transformations.

When forming a Cl2 molecule, both Cl atoms take part in bond formation and they increase the number of electrons in their outermost shells by 1 unit; that is, both outermost shells contain 8 electrons. When forming a NaCl molecule, the Na atom contains 2 electrons in the outermost shell (1 previous one and 1 acquired one) while the Cl atom contains 8 electrons.

The described manner of bonding atoms together is known as covalent bonding. The bond formed between such atoms is known as a covalent bond.

The bond used to connect identical atoms is called a homoatomic covalent bond; if the atoms are different, the bond is called a heteroatomic covalent bond [The Greek prefixes homo and hetero mean identical and different respectively].  

When the atoms' first ionization energies (FIEs) differ greatly (as during bond formation between Na and Cl atoms), the formed bond in the NaCl molecule is known as a polar bond.  

Atoms with more than one electron in the outermost shell can form not one but several covalent bonds between themselves. Such bonds are called multiple bonds. Examples of such bonds are those of nitrogen (NΞN) and oxygen (O = O) molecules.

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