La no respuesta en las encuestas electorales

Bearing in mind that the whole world and we ourselves are all made up of substances, here are the main questions concerning substances that are answered by general chemistry:

1. What is the smallest particle of a substance?

2. At the expense of what forces do these particles bond with each other?

3. How do the bonds between particles break up in the course of transformation? 4. How is the likeness and difference of substances that surround us defined?

Physics Nobel Prize winner Richard Feynman asked: "What is the shortest phrase, concerning the most important scientific knowledge we possess, would you pass on to the next generation?" Feynman, himself answered: "Everything is made of atoms."

It is a known fact that all the substances around us (iron, stone, salt, bread, liquids, etc.) consist of molecules, which are the smallest particle of a substance consisting of atoms. An atom consists of a positively charged nucleus around whichnegatively charged electrons rotate. A nucleus consists of nucleons (protons and neutrons).

Protons are always positively charged, their charges (in absolute units) being equal to those of the electrons.

Neutrons do not carry any electrical charges. Protons and neutrons consist of quarks.

There are 6.02 · 1023_{molecules of water in 18 grams of water. A molecule of water consists of} an atom of oxygen (O) and two atoms of hydrogen (H), i.e., H2O.

An oxygen atom consists of a nucleus and 6 electrons. A hydrogen atom consists of a nucleus and one electron.

To break up water into molecules, we must heat the water to a temperature of 100° C. To further break up a water molecule into atoms, we must heat the steam of the water to a temperature of about 5,000° C.

To tear away the electrons of the oxygen atom from its nucleus, it is necessary to raise the temperature to more than 10,000° C.

Thus, according to experimental data, the limit for breaking up particles into the smallest unit possible depends on the temperature (energy) that we use in order to break up the substance.

* Atomic structure

As already stated, the atom is made up of a positively charged nucleus around which negatively

charged electrons rotate.

The nuclear charge is equal to the sum of the proton charges in the nucleus. The positive charge of a proton is equal to the negative charge of an electron (in absolute units). The number of electrons rotating around the nucleus is equal to the number of protons. That is why the sum of the atom's charges is equal to zero.

The electrons are distributed around the nucleus in layers, like clouds around the earth. In the first layer, closest to the nucleus, there are always 2 electrons, starting with elements in the second period; in each of the outer layers there are 8 electrons, or less in the second and third period. The atom's outermost layer contains some number of electrons from 1 to 8. The main forces inside the atom are the electrostatic forces.

Electrons with identical negative charges repulse each other. They remain in the atom due to their attraction to the atom's nucleus, which has a positive charge. The electrons do not fall into the nucleus because they rotate at a great speed around it. During the electrons' rotation around

the nucleus, a force appears which is similar to that which we get if we rotate a spring with a metal ball at the end of it. This force is known as the centrifugal force. Centrifugal forces are directed in the opposite direction from the electrostatic forces, thereby drawing the electrons from the atom's nucleus.

If neutral atoms have an identical number of electrons, the distribution of electrons in the layers of these atoms is also identical. That is, the structure of the electronic shells of these atoms is the same. Substances formed of identical atoms (same nuclear charges or same number of protons and electrons) are known as elements.

Currently, 109 elements are known. All these elements are organized into a table of elements, a copy of which usually hangs on the walls of chemistry classes. Each box in the table of elements contains an element's symbol and name.

Above, we have already used some of the symbols and names of elements when giving the formula of water (H2O), oxygen (O), hydrogen (H) and other elements.

The element's number (given in the box) coincides with the amount of protons that the atom's nucleus of the given element contains. Each column is called a group and is marked with Roman numerals, and each row is called a period and is marked with Arabic numerals.

Elements of one group all have close physical and chemical properties. The elements of main group I, are Lithium (Li), Sodium (Na), Potassium (K), Rubidium (Rb), Cezium (Cs) and Francium (Fr), known as alkali metals. They are light and soft, with low boiling and melting temperatures and have high heat conductivity and good electrical conductivity. These elements can readily take part in reactions.

The elements of main group VII (halogens) are Fluorine (F), Chlorine (Cl), Bromine (Br) and Iodine (I). These are also highly reactive elements, but unlike the elements of group I, they are all non-metallic. They have low heat conductivity and do not conduct electricity.

The elements of main group VIII include inert gases Helium (He) and Neon (Ne) and noble gases Argon (Ar), Krypton (Kr), Xenon (Xe) and Radon ( Rn). Unlike other elements, these do not form stable molecules of the He2type and are mainly chemically inert elements.

Definition of First Ionization Energy

All atoms, when they have sufficient energy, can lose their electrons and turn into positively charged ions.

The energy necessary to tear an electron away from the atom is known as ionization energy. The energy necessary to tear the first electron from the atom is known as the first ionization

energy (FIE). The more energy we need to tear away the electron, the more stable the given

Most of the neutral atoms are capable of connecting to an electron spontaneously (with energy discharge). This property is called the affinity of the atom to the electron. The energy discharged during this process is called the atom's affinity energy to the electron.

The main obstacle for electrons that strive to get into the shell of these atoms is the inter- electronic repulsion force of the electrons already inside the atom, which hinder additional electrons from entering the outermost shell

Let's take the atomic structure of a hydrogen atom as an example. This atom is made up of a nucleus and only one electron. The electron has a negative electric charge, while the nucleus has a positive charge equal to that of the electron. The total value of the atom's charge is zero.

The mass of the electron is 1,840 times smaller than that of the nucleus. That is, the atom's mass is almost completely situated in the atom's nucleus, around which the electron constantly rotates, at a distance of 0.53 · 10-10_{m and a speed of 3 · 10}8_cm/sec.

The electron is attracted to the nucleus by electric forces. The attractive force of the electron to the nucleus (bonding force F) is proportional to the charge (Z) of the nucleus and the electron; it is inversely proportional to the square of the distance (R) between them. That is:

F = Ze/R2 _(3.0)

As we see, the smaller the distance between the electron and the nucleus (R), and the greater the nuclear charge (Z), the greater the attractive force of the electron to the nucleus. Therefore, we need more energy to tear the electron away from the nucleus.

The more energy is required to break the bonding, the more stable the system. That is, if the breaking of the bond (the separation of the electron from the nucleus) in one system requires more energy than in another system, then the first system is more stable. The energy required for the breaking of one gram of a hydrogen atom has been defined experimentally. It is equal to 13.6 eV (electron volts).

The energy necessary to break the electron away from the nucleus in atoms composed of one electron and having a positive nuclear charge twice that of a hydrogen atom was also experimentally defined. It was found that in this case it was necessary to spend four times more energy (54.4 eV). That is, an atom with a two times greater charge is four times more stable. According to electrostatics, the energy (T) that is necessary to break two bonds by charges with opposite signs (Z and e) situated at a distance of R from each other is defined by the equation:

T = Ze/R (3.1)

The energy necessary to break the bonds is proportional to the charge values and inversely proportional to the distance between the charges.

Such a correlation is quite logical; the greater the charges, the greater their attraction to each other. Therefore, more energy must be spent on breaking the charges. The smaller the distance between them, the more energy must be spent on breaking them. This makes it clear why an atomic system with a nuclear charge two times greater than another’s is more stable and requires more energy to be broken up.

However, the following question requires additional explanation:

Why is it that when the nuclear charge is doubled, the amount of energy necessary for breaking the bond between the nucleus and the electron increases by four times or is equal to the square value of the doubled nuclear charge (54.4 /13.6 = 4)?

This is especially incomprehensible if we return to equation 3.1, according to which if one of the charges is doubled, the energy necessary for breaking the bond should also be doubled but not squared.

This discrepancy is due to the following:

In systems where charges Z and e are at rest relative to each other, energy T is, indeed, proportional to Z. That is, when charge Z increases, energy T increases proportionally. Unlike these systems, in atomic systems where the electron with charge e rotates around a nucleus with charge Z, and charge Z is increased, the radius of rotation R is proportionally decreased. This is because the electron is attracted to the nucleus with a greater force.

Graphs related to FIE

Now let's see how the FIE changes in regard to the change of the nuclear charge.

In figure 3.1, we see the dependence of the FIE on the nuclear charge for the first 20 elements of the table of elements.

Figure 3.1

According to experimental data, when nuclear charge changes with a simultaneous increase in the number of electrons in the outermost electronic shell, the FIE increases in five cases and decreases in two during the period.

Thus, for example, the FIE of lithium (Li), with a nuclear charge of three proton units, is equal to 5.4 eV; the FIE of beryllium (Be), with a nuclear charge of four proton units and whose outermost shell contains 2 electrons, is equal to 9.3 eV. That is, when the nuclear charge of a Li atom increases by one proton unit with a simultaneous increase in the number of electrons in the outermost shell by one unit, the FIE increases by 3.9 eV (i.e., 9.3 - 5.4 = 3.9) as we transit from Li to Be.

Unlike the transition from lithium (Li) to beryllium (Be), the transition from beryllium (Be) to boron (B) shows that the FIE decreases. Thus, if Be has a FIE of 9.3 eV, then B has a nuclear charge of four proton units, with three electrons in the outermost electronic layer, and the FIE is equal to 8.3 eV, i.e., by 1 eV less (9.3 - 8.3 = 1).

When transiting from boron (B) to carbon (C), from nitrogen (N) to oxygen (O), from oxygen (O) to fluorine (F), and from fluorine (F) to neon (Ne), the FIE increases by 3.1 (11.4 - 8.3 = 3.1, 14.5 - 11.4 = 3.1), by 3.8: (17.4 - 13.6 = 3.8) and by 5.2 (21.6 - 17.4 = 5.2), respectively. That is, here we see regularity like that seen during the transition from Li to Be.

When transiting from nitrogen (N) to oxygen (O), the FIE decreases by 0.9 eV (14.5 - 13.6 = 0.9); this is a dependence analogous to the one we had when transiting from beryllium (Be) to boron (B).

When transiting from neon (Ne) to sodium (Na), the FIE decreases by 16.46 eV (21.6 - 5.14 = 16.46). In the course of this transition, the nuclear charge also increases by one proton unit and one electron is added.

Thanks to experimental data, we know the IEs for all the electrons in the elements. In figures 3.2-3.5 (next two pages), we see the logarithmic dependence of the energy required for the consequent extraction of electrons from the Be, B, N and Ca. atoms.

The logarithmic scale is used so that the graphic does not stretch out vertically.

Figure 3.2

Figure 3.4

Figure 3.5

The electrons situated on one straight line belong to the same layer (shell). The electrons of one shell are situated at approximately equal distance from the nucleus. This is why during successive breakaways of the electrons (as a result of ionization), the ionization energies of the consecutive electrons gradually increase, while the repulsive forces between the electrons the same layer decrease. This is in accordance with electrostatics.

During the transition to the next layer, the difference between the IEs of the consecutive electrons abruptly increases. This is clearly seen via the difference of the FIEs in the elements

that have been mentioned above. The difference in the FIEs between the previous and consecutive elements, located in the same period, comprises less than 4 eV.

When forming a new shell, the comparison of the FIEs of the last element of the second period of neon (Ne) and the first element in the third period of sodium (Na), this difference increases to 16.46 eV. That is, according to the data on the FIEs, Li, Be, B, C, N, O, F, Ne and Na contain 1, 2, 3, 4, 5, 6, 7, 8, and 1 electron respectively in their outermost shells.

The study of ionization energies (necessary to tear an electron off the atom) of various atoms has shown that the electrons in the atoms are situated in layers. There are two electrons in the first shell nearest to the nucleus, and there are 8 electrons in each of the other inner shells.

According to experiments, the number of electrons in the outermost layers of the atoms changes periodically when the nuclear charge increases. For elements with fewer than 20 electrons, the maximal number of electrons in the outermost shell (layer) is eight. That is, the number of electrons in the outermost shell changes periodically from one to eight when the nuclear charge increases (see Figure 3.1).

The layered structure of the electronic cloud surrounding the nucleus and the periodic change in the amount of electrons in the outermost shell are explained by the fact that, during the gradual increase of the number of electrons (when the shell is being filled), inter-electronic repulsive forces begin to exceed the electron's attraction to the nucleus, and the joining of the electrons to the outermost shell requires additional energy.

Chapter 4. Molecules

We know that molecules are made up of atoms. As we spend energy to break a molecule into atoms (as 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 occurs 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 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, two electrons are sufficient, since each of the hydrogen atoms bonded into a H2molecule has only one electron.

All other atoms 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 is not clear why the helium atom (He), which has two electrons, does not form stable molecules of the He2type.

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 to amounts to 1, 2, 3, 4, 3, 2, 1 and 0, respectively. That is, the atoms of Li and F form stable molecules LiH and HF, while B and N form stable molecules BH3and NH3. Ne does not form any stable molecules with hydrogen.

Conditions of chemical bond

As indicated in the previous section, the number of electrons in the outermost shells of the Li, Be, B, and C atoms comprises one, two, three, and four 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 H2and Cl2type, 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 (F) bond to 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 to 8 of the number of electrons, the atoms of the 2nd_{period begin to form a new outermost shell.}

A comparison of this data with that of the number 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 an atom of the 2nd _{period can bond to is limited to the} maximum number of electrons that the outermost shell of tbjs atom can contain. According to the data on FIEs, this number is equal to 8.

According to these conclusions, Ne, which already has 8 electrons in its 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 the 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 a possibility of 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 minimum number of bonds that an atom can form is determined by the amount of electrons present in the outermost shell of the given atom. For atoms of the 2nd and 3rd