Explain different Characteristics of Periodic Properties. When we move vertically or horizontally through the groups in the periodic table, it can be seen that there is a relatively uniform change in the properties exhibited by the elements, which is why they are known as the periodic property.
The periodic law tells us that if we order the elements according to their increasing atomic number, we periodically find elements that have similar chemical and physical properties.
Elements in a group have similar properties, while the properties of elements in a period change progressively as you move through the table.
Explain different Characteristics of Periodic Properties
The observed trends in periodic properties, such as atomic radius, ionization energy, and electronic affinity, are related to the electronic configuration and effective nuclear charge of the elements.
They, therefore, allow the interpretation and prediction of chemical behaviour and physical properties.
which are in the same group or period. Atoms are currently seen as extremely small spherical particles, with a nucleus (with protons and neutrons) at its centre and surrounded by electrons.
Therefore, it can be said that the size of an atom is given by its atomic radius, which can be defined as the distance between the nucleus and the farthest electron from it.
However, because the electron cloud that surrounds the nucleus has no definite boundary, the size of an atom is determined by its interactions with the atoms around it.
One way to estimate the atomic radius is to consider the distance between the nuclei of two adjacent atoms.
Explain different Characteristics of Periodic Properties
Imagine a collection of argon atoms in gaseous form.
When two atoms collide, they bounce during their motion just like a billiard ball.
Let us stop when two atoms are at the point where they are close together during the collision, at which time the distance separating the nuclei is twice the radius of the atoms; That is, the radius of each atom is equal to half of that distance.
This type of radius can be called a non-binding atomic radius because the atoms are not forming molecules.
Atoms are so small that the unit of length used to express atomic radius is ngström or angstrom (Ǻ).
One ngström is equal to 1×10 -10 metres, which is an extremely small distance because 1 is one-millionth of a millimetre.
The radius of a caesium atom, which is practically the largest element in the periodic table (surpassed only by francium) is estimated at 2.65, while a helium atom, which is the smallest, measures only 0.31.
Scientists have developed a variety of experimental techniques to measure the distances separating nuclei in a molecule.
Based on the observation of these distances, in many molecules, it is possible to assign an atomic bond radius to each element.
For example, in the I 2 molecule, the distance separating the nuclei of iodine (I─I) is 2.66, so the atomic bond radius of iodine is 1.33.
As we move from top to bottom through the same group of the periodic table, the atomic radii of the elements increase, this is due to the fact that electrons occupy orbitals with energy levels that are close to that of the nucleus. further and further away.
That is, the valence electrons of the elements in period 1 occupy the orbitals of level 1 which are the orbitals of period 2 of level 2, the orbitals of period 3 of level 3.
And so on (for example, if we pass through group IA, the valence electron will be hydrogen’s 1 s orbital, lithium’s 2 s, sodium’s 3 s … francium 7 s).
Explain different Characteristics of Periodic Properties
On the other hand, as we move from the left to right period, the atomic radius of the representative elements (blocks s and p) decreases.
This may sound strange because as the atomic number (Z) increases, the number of electrons around the nucleus increases, however, the e- of the outermost shell fills orbitals with the same energy level, so it is said to be Maybe they are populating orbitals.
which are equidistant from the centroid. Keep in mind that as Z increases, so does the number of protons, so the attraction they exert on electrons will also increase, so that as we move through the period, the effective nuclear charge increases, Due to which the atomic radius decreases.
When atoms lose electrons they are no longer electrically neutral and remain positively charged (forming cations) because the number of protons is greater than that of electrons.
This disparity between protons and electrons causes the atomic radius to decrease as the effective nuclear charge increases.
That is, the attraction exerted by the proton on the remaining electrons is greater than the attraction on the electrons of the parent atom.
For example, a ground state Ca atom (Z = 20) has 20 protons (p +) that attract 20 electrons (e -) to the nucleus.
Losing two e-calcium ions (Ca 2+), 20 p+ should only attract the remaining 18 e-, so the force of attraction exerted by them on the latter is greater.
Therefore, it can be said that the radius of cations is always smaller than the radius of the neutral atoms from which they come.
Conversely, when an atom of an element gains electrons, forming an ion, the atomic radius increases, this is because the proton must attract a greater number of electrons to the nucleus.
Because they are attracted to the neutral atom. , so the intensity with which they attract them is less.
For example, we have, for example, a chlorine atom (Cl) with Z = 17 containing 17 p+ and 17 e-.
But when the electron is gained by forming a chloride ion (Cl 1 ), the same is 17p+. The nucleus should now attract 18 e-.
So they reduce the force of attraction on them.
Consequently, it is possible to say that the radius of ions is always greater than the radius of the neutral atoms from which they arise.
Electronic affinity, ionization energy and electronegativity give us information about the chemical behavior of elements.
Based on them, it can be predicted how these will be connected and what kind of link they will become.
Ionization energy (EI) is the energy that must be applied to an atom of an element in a ground state so that it can give up an electron and, as a result, form a cation with a 1+ charge.
For example, for calcium, the first ionization energy (EI 1 ) is 599 kJ/mol:
Explain different Characteristics of Periodic Properties
The second ionization energy is always higher than the first.
Because of the attraction exerted by protons in the nucleus on the remaining electrons, it is more difficult to separate an electron from a cation than from a ground-state element.
It is important that you understand and keep in mind that atoms are electrically neutral, which means that the number of protons (+) and electrons (-) are equal.
As you are well aware, this information is expressed by the atomic number (Z).
In the case of calcium (Ca), which has Z = 20, in the ground state, the 20 protons in the nucleus attract the 20 electrons found in the orbital with a certain intensity;
But when the Ca atom loses an electron, forming the Ca 1+ ion, the 20 protons will attract the remaining electrons to the nucleus with greater intensity because there are only 19 e – left.
As a result, more force (energy) must be applied to separate a second electron and form a cation with a 2+ charge.
For this reason, the second ionization energy (EI 2) is always higher than the first, so the EI for calcium is:
Ionization energy always increases in the following order:
In the table below you can see how the gradual ionization energy increases for the elements of period 3.
In general, the process by which cations are formed can be expressed as follows:
Returning to calcium, the first ionization energy can also be expressed as:
The fact that ionization energy has a positive sign means that energy must be applied to the atom to release an electron.
Therefore, cation formation is an endothermic process.
Metals have low ionization energy, so they can easily lose electrons and therefore form cations.
Explain different Characteristics of Periodic Properties
It can be said that electronic affinity is the opposite of ionization energy because it is the energy that is released when an atom of an element in the gaseous state gains an electron.
and forms an ion with a charge of 1-.
Thus for chlorine when the chloride ion is formed, the process can be expressed as:
Hence, the electron affinity of chlorine is -349 kJ/mol. The negative sign means that energy is released, that is, it is an exothermic process.
Nonmetals have high electronic affinities, so they can readily accept electrons and form ions.
If it is still not clear to you what are ionization energy and electron affinity.
And what is their difference, then click on the button and read the following example carefully.
Click on each button to view the information.
Finally, both ionization energy and electron affinity increase as we move through the groups from top to bottom and from left to right across the period of the periodic table.
Atomic radius, ionization energy, and electronic affinity are properties of individual atoms, however, by grouping elements into three large categories: metals, nonmetals and metals, it is possible to see that they each have similar properties. ,
It is possible to observe this classification in the following figure.
It is important to note that although hydrogen is in group IA, it belongs to nonmetals.
When an element exhibits more and more physical and chemical properties of metals, higher will be its metallic properties.
Explain different Characteristics of Periodic Properties
The metallic properties of elements increase as we move from right to left in a period and from top to bottom in a group of the periodic table.
Most metals have a brilliant lustre and conduct heat and electricity.
are deplorable and deplorable. All are solid at room temperature except mercury (melting point = -39 °C) which is a liquid.
Two metals melt at a temperature slightly above room temperature:
Caesium at 28.4 °C
and gallium at 29.8 °C. At the other extreme, many metals melt at very high temperatures, for example,
Chromium melts at 1900 °C.
As you will recall, due to their low ionization energy, metals tend to lose electrons, so they can easily form cations.
Compounds of metals with nonmetals are usually ionic substances, such as binary salts, for example, aluminium bromide (AlBr 3).
Nonmetals vary greatly in appearance, are not lustrous, and are generally poor conductors of heat and electricity.
Their melting points are generally lower than those of metals (although diamond, a form of carbon, melts at 3570 °C).
Seven nonmetals exist as diatomic molecules under normal conditions;
Five of them are gases (H2, N2, O2, F2 and Cl2), one is a liquid (Br2) and one is a volatile solid (I2).
The remaining nonmetals are solids that can be as hard as a diamond or as soft as sulphur.
It should be remembered that due to their high electronic affinity, nonmetals gain electrons when they react with metals.
For example, the reaction between aluminium and bromine forms aluminium bromide, a compound containing the aluminium ion Al 3+ and the bromide ion Br 1-.
Nonmetals usually acquire enough electrons to completely fill their outer p subshell to obtain a noble gas electron configuration.
For example, the bromine atom gains one electron to fill its 4p subshell:
Compounds composed entirely of nonmetals are molecular substances. For example, nonmetal oxides, halides, and hydrides are molecular substances that are usually low-melting liquid or solid gases at room temperature.
Metalloids have properties intermediate between those of metals and nonmetals; They may have some properties typical of metals, but lack others.
For example, silicon looks like metal but it is brittle rather than ductile and does not conduct heat and electricity as well as metals.
Many metalloids, especially silicon, are electrical semiconductors and are the main element used in the manufacture of integrated circuits and computer chips.
In short, it should be remembered that the atomic radius corresponds to half the distance between the nuclei of two adjacent atoms or the distance between the nucleus and the electron that is farthest from it.
Atomic radius is measured in ngström (Ǻ).
Ionization energy is the energy that must be applied to an atom of an element in a fundamental state in order for it to give up an electron and form a cation.
Whereas electronic affinity is the energy that is released when an atom of an element gains an electron and thus forms an ion.
As we move from top to bottom through the same group and from right to left through the period of the periodic table, the atomic radii of the elements begin to increase.
Whereas electron affinity and ionization energy increase inversely, i.e., from bottom to top and from left to right.
With regard to the radii of ions, it should be borne in mind that the radii of cations are always less than the radii of the atoms from which they come, whereas the radii of anions are greater.
While grouping the elements into metals, non-metals and metalloids, it is observed that they present similar properties.
In the case of metals, they have a brilliant lustre, are good conductors of heat and electricity, are malleable and ductile; Nonmetals vary greatly in appearance but are not lustrous.
and are generally poor conductors of heat and electricity, and metalloids exhibit properties intermediate between those of metals and nonmetals.
The metallic properties of elements increase as we move from top to bottom in the period of the periodic table and from right to left.
Because of their low ionization energy, metals easily lose electrons, so it is easier for them to form cations.
Non-metals, due to their high electron affinity, tend to accept electrons so that they can form anions easily.
Therefore, metals do not react with each other.
But when metals are chemically combined with non-metals, they form ionic compounds, such as binary salts. With respect to non-metals, they can form compounds among themselves, which are molecular substances.
For example, non-metal oxides, halides, hydrides of non-metals and diatomic gases such as H2 and N2.
