Do you know what the base of chemistry is? Yes, elements it is! Do you know every element has certain unique properties? These unique properties are referred to as the periodic properties of elements. So, let’s study in detail about the periodic properties of elements.
Periodic Properties of Elements
In the periodic table, the elements are arranged in the order of their increasing atomic number. All these elements present several different trends and we can make use of the table formation and the periodic law to understand and predict the chemical, physical, and atomic properties of the elements.
In order to understand the trend, you’ll have to analyze the element electron configuration. All the elements prefer an octet formation and tend to gain or lose the electrons in order to form that stable configuration. Now let’s learn some of the periodic properties of elements in detail and study the trends in these properties in the periodic table.
1) Atomic Radius
We can never determine the atomic radius of an atom. You know why? This is because there is never a zero probability of finding an electron, and thus there is never a distinct boundary to the atom. So, the only thing that we can really measure is the distance between two nuclei (internuclear distance).
A covalent radius is one-half the distance between the nuclei of two identical atoms. An ionic radius is one-half the distance between the nuclei of two ions in an ionic bond. The distance must be apportioned between the smaller cation and larger anion. A metallic radius is one-half the distance between the nuclei of two adjacent atoms in a crystalline structure.
The noble gases are left out of the trends in atomic radii because there is great debate over the experimental values of their atomic radii. The SI units for measuring atomic radii are the nanometer (nm) and the picometer (pm).
1 nm = 1 × 10-9 m; 1 pm = 1 × 10-12 m
Screening and Penetration
In order to help you understand the trend of atomic radius, you need to first understand the concept of screening and penetration. Penetration is the distance at which the electron is from the nucleus. Screening is the process in which the inner electrons block the outer electrons from the nuclear charge.
In this concept, let us assume that there is no screening between the outer electrons and that the inner electrons shield the outer electrons from the total positive charge of the nucleus. To comprehend the extent of screening and penetration within the atom, scientists came up with the effective nuclear charge, Zeff. The equation for calculating the effective nuclear charge is shown below:
Zeff = Z − S
where Z is the atomic number and S is the number of shielding electrons.
Understanding the Equation
In the equation above, S is the number of inner electrons that screen the outer electrons. Students can easily find S by using the atomic number of the noble gas that is one period above the element. For example, the S we would use for Chlorine would be 10 (the atomic number of Neon). Z is the total number of electrons in the atom.
Since we know that a neutral atom has an identical number of protons and electrons, we can use the atomic number to define Z. For example, Chlorine would have a Z value of 17 (the atomic number of Chlorine). Continuing to use Chlorine as an example, the 10 inner electrons (S) would screen out the positive charge of ten protons. Therefore, there would be an effective nuclear charge of 17+(-10) = +7.
The effective nuclear charge shows that the nucleus is pulling the outer electrons with a +7 charge and therefore the outer electrons are pulled closer to the nucleus and the atomic radii are smaller. In summary, the greater the nuclear charge, the greater pull the nucleus has on the outer electrons and the smaller the atomic radii.
In contrast, the smaller nuclear charge, the lesser pull the nucleus has on the outer electrons, and the larger atomic radii. Additionally, as the atomic number increases, the effective nuclear charge also increases.
Atomic Radius Trend in the Periodic Table
Now, let’s understand the atomic radius trend in the periodic table.
- Across the table, the atomic number increases moving left to right and so does the effective nuclear charge. In the nutshell, when moving left to right across a period the nucleus has a greater pull on the outer electrons and the atomic radii decreases.
- Now, when you move down a group in the table, the number of filled electron shells increases. However, the valence electrons keep the same effective nuclear charge, but now the orbitals are farther from the nucleus. As a result, the nucleus has less of a pull on the outer electrons and the atomic radii are larger.
2) Ionization Energy (Ionization Potential)
Ionization Energy is one of the most important periodic properties of elements. In order to remove an electron from an atom, you need sufficient energy to overcome the magnetic pull of the positive charge of the nucleus. So, the ionization energy (I.E. or I) is the energy that is required to completely expel an electron from a gaseous atom or an ion. The Ionization Energy is always positive.
The energy required to remove one valence electron is the first ionization energy, the second ionization energy is the energy required to remove a second valence electron, and so on. For example, consider the element Sodium(Na). Let’s give the equation for its first and second ionization energy.
- 1st Ionization energy: Na(g)→Na+(g) + e
- 2nd Ionization energy: Na+(g)→Na2+(g) + e
As you move left to right across a period, ionization energies increase. Also, when you go up a group, ionization energy increases. This is because there are fewer electrons shielding the outer electrons from the pull of the nucleus.
A Solved Question for You
Q: What is electron affinity?
Ans: Electron affinity (E.A.) is the energy change that occurs when an electron is added to a gaseous atom.