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Home  / GENERAL CHEMISTRY Textbook / Chapter 11. Physical and chemical properties of substances

Chapter 11. Physical and chemical properties of substances

PHYSICAL  PROPERTIES  OF  SUBSTANCES 

During the development of chemistry, new terms were introduced: atom - the smallest species of an element; molecule - the smallest species of a substance. Transformations of substances proceeding without any change of the molecule structure are called physical transformations, while transformations that proceed with a change of the molecule structure, are called chemical transformations.  

It is not difficult to see the difference between the physical and chemical transformations. Physical transformations include melting and evaporation of substances; physical properties include glitter, color, and electro-conductivity. Chemical transformations deal with chemical properties like chemical reaction capabilities. Solubility and dissociation occupy an intermediate position.   

According to our theory of chemical bonding, bonds that break in the course of physical and chemical processes have identical electrostatic natures but they differ relative to their energies.   

There are 3 main types of bonds:

  1. Type I includes bonds in which the energy's affinity to the electrons of both bonding atoms is consumed (covalent, homopolar and heteropolar bonds).
  2. Type II includes bonds where the affinity energy of only one of the bonding atoms is consumed [donor-acceptor bonds (DABs)].
  3. Type III includes bonds in which the affinity energy is not consumed in either of the atoms [Van der Waals bonds (VWBs)]. 

These three types of bonds have one and the same electrical nature. The main energy gain during bond formation is conditioned by the concentration of the positive charge (70%-90%) and the atoms' affinity towards the electron (10%-30%). In spite of the fact that the atom's affinity to the electron is less than 30%, the elimination of the latter can decrease the bonding energy by ten times. This is because 30% of the attraction energy (attraction of the electrons to the molecule's nuclei) is defined by the value of 300 kJ/mol, which is commensurable with the bonding energy.     

The above-mentioned correlations between contributions have to do with the attraction power between electrons and nuclei. In fact, such a distribution of energy is but conditional. Likewise, the whole energy gain is achieved at the expense of the increase in the effective charge of both nuclei during molecule formation.

Because the nature of all the three types of bonds is identical, these bonds differ only in the energy applied to break them.  In order to break the bond of type I, energy of 50 to 500 kJ/mol is required; bonds of type II require 25 to 200 kJ/mol; bonds of type III - 5 to 30 kJ/mol.

According to thermodynamics, and considering the electronic heat capacity, the energy of the system can be defined by the equation: E = 10 RT (where is the Boltzmann constant  and T is the absolute temperature). This equation defines the temperatures of three types of bond breaking: types I, II, and III; respectively: 4,000 K; 1,500 K; and 400 K. 

As already indicated, an atom is usually bound to other atoms by various types of bonds.  For example, in solid metals each atom is bonded with other MEB, dynamic, single, double, or triple covalent bonds, or with DABs.  Historically, this type of bond has been singled out into a separate kind called metallic bond.   

Among the special bonds are conjugated bonds (combinations of double and single bonds) and aromatic bonds, like those in benzene molecules (where single and double bonds alternate with each other). Each of these bonds is manifested in special physical and chemical properties of substances. 

If an atom is connected to other atoms by various bonds, the breaking energy of the weak bonds increases, while that of the strong bonds decreases. For example, the bond-breaking energy in benzene comprises 520 kJ/mol, while the breaking energy of a single bond is equal to 350 kJ/mol and the same of a double bond is equal to 600 kJ/mol. As a rule, we can assume that the energies of the weak bonds increase as do many of the strong bonds and their bonding energies. 

Most of the strong covalent bonds have elements in which the outermost electronic shell is half-filled. When the number of electrons in this shell changes, as compared to the half-filled layers, the bonds break more readily. When the number of electrons decreases, the quantity of DABs increases at the expense of a decrease in the number of covalent bonds. When the number of electrons increases, the covalent bonds decrease in number. Thus, instead of DABs, we get very weak VWBs. 

Carbon, in the form of a diamond, (carbon forms 4 identical bonds) is the hardest substance ever. A diamond's compression coefficient is equal to 0.16 · 10-6 cm2 kg-1, its melting-point is higher than 3,550° C, and its boiling temperature is 4,827° C. Carbon crystal has the highest evaporation heat - 718 kJ/mol; and melting heat of 104 kJ/mol. Its C-C bond is the strongest covalent bond. The physical properties of carbon are in agreement with our new theory of chemical bonding.

In the case where one electron is situated on the atom's outermost shell and the FIE of this atom is less than 500 kJ/mol, we should expect an abrupt decrease in the bonding energy of metals formed from such atoms; this is followed by an abrupt decrease in their melting, boiling heat and temperatures. The properties of alkaline metals (Li, Na, K, etc.) differ greatly from the properties of other metals.

For example: melting temperatures vary from 28° C (Cs) to 179°  C (Li); their melting heat amounts to 2 kJ/mol (Cs) and 3 kJ/mol (Li); their evaporation heat is equal to 67 kJ/mol (Cs) and 148 kJ/mol (Li). These metals are extremely soft.  According to our theory, the distance between the nuclei in these metals is greater than in other metals, which is noticeable in their very low density; indeed, lower than that of water. 

Physical properties depend not only on the number of electrons in the outermost shell, they also depend on the kind of bond the given atom forms. Atoms can form not only single bonds, but also double and triple ones. For example, the outermost shells of nitrogen (N) and oxygen (O) atoms are filled to the limit by 8 electrons. In liquid and solid states, both N2 and O2 have multiple bonds between the atoms, and VWBs, between the molecules.   

The melting temperature for solid nitrogen is -209° C, the melting heat is 0.72 kJ/mol, the boiling temperature is  -195° C, and the evaporation heat is 5.38 kJ/mol. Oxygen has data close to this, the boiling temperature being -218° C, and the melting temperature being -183° C.   

It is known that the physical properties of elements are influenced not only by electronic structure, but also by the atoms' masses of these elements. The dependence of the atomic masses on the density of the substances is understood; the influence of the nuclei's mass on the evaporation heat and on the boiling point is evident from the kinetic gas theory, according to which — the greater the element's atomic mass, the smaller its kinetic energy at the same temperature.  

Atoms and molecules with a greater mass, when their potential energies are identical, achieve their kinetic energies at a much higher temperature. Besides depending on the mass of the nuclei, the physical properties of substances also depend on spatial structure of substances, which is described in the chapter devoted to this issue. 

Now let us return once more to the explanation of the regularities concerning the physical properties of substances in the framework of our theory of  chemical bonding. 

When the number of electrons in the outermost shell is equal to 4 (half of the maximal filling quota), the substances have the highest boiling temperature and evaporation heat. Any changes in the amount of electrons leads to a decrease in these values.  The expected dependence is obvious in the first 20 elements of the periodic table with a decrease in the number of electrons. This dependence is even more obvious when comparing the physical properties of alkaline and other metals, as previously indicated. 

As previously shown, the decrease of the evaporation heat and boiling temperature of nitrogen and oxygen is achieved by the formation of multiple bonds. The anomalous properties of halogens are likewise explained by the complete filling of the outermost electronic shell of halogen atoms during the formation of dual atomic molecules. The dependence of dual atomic bonding energy on the FIE proceeds via the maximum. The elements in the middle of the period have maximal energy bonds; i.e., atoms with 4 electrons in the outermost shell form not only a maximal number of covalent bonds, but also— the strongest ones. 

Thus, we can expect that the dependence of the physical properties of simple substances on the elements' FIEs have the same character as the dependence of homoatomic bonding energies in dual atomic molecules on the FIEs. A comparison of character dependencies of some elements' physical properties on the FIEs (table 11.1-1, figures 11.1-1 to 11.1- 4) with dependencies of homoatomic bonding energies on the FIEs (see figures from 6.4-5 to 6.4-6) shows that these dependencies are identical. Also, table 11.1-1, figures 11.1-1 to 11.1-4 show the data received from studying normal (not anomalous) metals. 

Additional confirmation comes from the anomalous behavior of zinc, cadmium, and mercury, which form weak covalent bonds. This is evident from their anomalous physical properties.  For example, mercury (Hg) has a melting point of -38° C, a melting heat of 2.33 kJ/mol, a boiling point of 356° C, and a boiling heat of 58 kJ/mol. While the element gold (Au) and the following element—thallium (Tl), have a melting point of 1,064° C & 302° C, a melting heat of 12.7 kJ/mol & 4.31 kJ/mol, a boiling point of 2,966° C & 1,457° C, and a boiling heat of 310 kJ/mol & 168 kJ/mol.  

The physical properties of elements following calcium are well explained on the basis of our theory with the supposition that the outermost electronic shells of atoms can contain more than 8 electrons. A comparison of the elements' FIEs from  19 to 28 shows that when the nuclear charge increases from 19 to 28, the number of electrons in the outermost shell increases from 1 to 12.  In the atoms of elements from  28 to 36 and from 47 to 54, 1 to 8 electrons fill the outermost shells.  

The dependence of these elements' FIEs on their placement in the rows and the rules of formation relative to both covalent and donor-acceptor bonds, are analogous to the first 20 elements of the table. The physical properties of these elements are explained by the same rules. 

For elements from #26 to #28, and from #44 to #46, the maximal amount of electrons in the outermost electronic shell can increase up to 18. Saturated shells appear in such cases. For example, stable metal compounds obey the rule of the 18 electrons: V, Cr, Mn, Fe, Ni, Nb, Mo, Ru, Rh, Pd. It is no wonder that the carbonyls of these metals, under usual conditions, exist not only as liquids: Ni(CO)4,  Fe(CO)5;  but also as gases: HCo(CO)4

The physical properties of substances, besides being influenced by bonding energy and the amount of bonds that an atom can form, are also greatly influenced by the spatial structure of substances.  

As already indicated in section 2.8, the angles between atoms are, as a rule, identical. The maximal distance between the bonds is conditioned by the mutual repulsion of the bonding electrons. When the distance is maximal, there is minimal repulsion energy, i.e., a minimal energy system.

However, the repulsion energy, which defines the spatial structure, slightly changes if there is a decrease in the angle between some bonds and, simultaneously, an increase in the angles between other bonds.

Thus, for example, carbon exists in the forms of diamonds and graphite. The angles between the bonds in a diamond are identical and comprise 109 degrees. In graphite the angles between three of the bonds are 120°, and there is a 90° angle between each of these three bonds and a fourth (on a perpendicular plane).

That is, the increase of the angle from 109 to 120 degrees between the three bonds, causing the decrease of the system's energy, compensates for the energy increase caused by the decrease of the angle from 109 to 90 degrees between the three bonds and the fourth. Graphite energy is smaller than diamond energy.

Interelectronic repulsion (between bonding electron pairs) leads to an abrupt decrease in the energy of the fourth bond, as compared to the bonding energy of the diamond. Graphite is a soft substance with a layer-like structure, while diamond is the hardest possible substance with an octahedron-like structure. 

      Table 11.1-1 

Melting Point, Fusion Heat, Boiling Point, Evaporation Heat of Simple Substances  vs. FIE of Elements 

  ELEMENT 

   FIE    

eV

MELTING Point

°C

FUSION HEAT 

kJ/mol

BOILING   Point

°C

EVAPORATION HEAT 

       kJ/mol

Na

5.1

97

2.6

882

90

Al

5.9

660

10.7

2467

291

Si

8.1

1414

 

2335

170

K

4.3

63

2.4

774

77

Sc

6.5

1539

16.1

2727

305

Ti

6.8

1657

18.6

3260

397

V

6.7

1890

17.5

3000

456

Cr

6.7

1890

14.6

2482

349

Cu

7.7

1084

13.3

2595

305

Ga

6

30

5.6

2403

267

Ge

9.7

958

 

2700

333

As

9.8

817

32

613

139

Se

9.7

220

5.2

684

14

Rb

4.17

39

2.2

688

75

Y

6.4

1495

17

2927

393

Zr

6.8

1852

5

3578

141

Nb

6.9

2468

26

3300

695

Mo

7.0

2610

27

5560

600

Te

7.3

2200

23

5000

500

Rh

7.5

1693

22

3700

500

Pd

8.3

1550

16

3200

360

Ag

7.6

960

11

220

254

In

5.8

156

3.3

2000

226

Sn

7.3

232

7

2270

290

Sb

8.6

530

20

1635

190

Te

9.0

449

17

1390

85

Cs

3.9

28

2

690

70

Nd

6.3

1024

 

3027

 

Eu

5.7

820

 

1440

 

Gd

 

1300

 

3000

 

W

 

3387

35

5900

799

Re

 

3180

38

5627

707

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11_1_1
Figure 11.1-1 

 


Figure 11.1-2

11_1_3
Figure 11.1-3


Figure 11.1-4

 

Chapter 11. Physical and chemical properties of substances    
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