Periodicity of Properties
According to the fundamental law regulating the contemporary periodic table, an element’s properties are periodic functions of its atomic number. These characteristics repeat frequently or follow a specific pattern frequently. The periodicity of properties is the name given to this phenomenon.
EXPLANATION
Due to the repetition of comparable electronic configurations with the same amount of electrons in the outermost orbit, periodic characteristics of elements develop. The number of valence electrons in a given group stays constant. On the other hand, when we move from left to right throughout a period, the number of valence electrons rises. The quantity of electrons in the valence shell determines an element’s chemical properties.
Periodicity of properties examples
- Shielding effect
- Ionization energy
- Electron affinity
- electronegativity
- Atomic size
Shielding Effect
The shielding effect refers to the decrease in the force of attraction between the nucleus and the valence electrons caused by the electrons in the inner subshells.
Explanation
When there are more inner electrons, they shield the outermost electron from the nucleus, allowing it to neglect the nuclear pull. This is known as the shielding or screening effect. In a multielectron atom, the electrons in the valence shell are drawn toward the nucleus but are repelled by the electrons in the inner shells. The real force of attraction between the nucleus and the valence electrons is significantly diminished as a result of the repulsive forces acting in opposite directions. Due to the presence of electrons in the inner shells, a phenomenon known as the screening effect or shielding effect causes the nucleus’s force of attraction on valence electrons to weaken.
• Trends of shielding effect concerning groups
As we go from top to bottom in a group the influence of the shielding effect increases because the number of the shell of elements increases, as well as the number of inner shell electrons, also increases.
• Trends of shielding effect concerning periods
As we go from left to right in a periodic table the number of shells remains the same. Therefore the shielding effect remains unchanged or constant.
Atomic Size
The distance between an atom’s nucleus and valence shell is known as its atomic size.It is also called atomic radii.
Variation in atomic radii in period
In any particular period, the atomic radius decreases as you move through it. This is due to the fact that, within a period, you shift from one element to the next to its right. In the same valence shell, an additional electron is added. The nucleus’s positive charge likewise rises by 1 at the same moment. The nucleus’s attraction to the electron in the valence shell grows. As a result, the atomic radius and shell size both decrease. For instance, the atomic size decreases from lithium to beryllium. This is seen from the electronic arrangement of Li (2s1) and Be valence shell (2s2). The atomic number grows from 3 to 4 while the shell number n remains unchanged while traveling from Li to Be. As a result, the nucleus’s force on the electron in the valence shell increases. So, atomic radius or atomic size decreases.
Variation in atomic radii in group
Moving down the group of elements in any particular major group causes the atomic radius to increase. This is so because an atom’s valence shell size determines how big it is. With each subsequent lower element in the group, the atom gains an extra electron shell. This increases the atomic radius. For example, the atomic radius increases when moving from Li to Na.
Ionization Energy
The least amount of energy necessary to remove the outermost electron from a gaseous atom in isolation is known as “ionization energy.”
M(g) + ionization energy⎯⎯→M (g) + e
Explanation
The ionization energy measures the extent to which the nucleus attracts the outermost electron. The stronger attraction between the nucleus and the outermost electron is indicated by a high value of ionization energy. The nucleus and the outermost electron are more weakly attracted to one another when the ionization energy is low. The unit of ionization energy is Kj/Mol.
Trends in the values of ionization energy
- In Groups
In a group, the ionization energy value decreases from top to bottom. This is due to the fact that as you descend in the group, the shielding effect in atoms gets stronger. The nucleus’s attraction to the valence electrons is weaker due to the greater shielding effect. Valence electrons are therefore simpler to remove. This causes a decrease in ionization energy in a group from top to bottom. For example, the ionization energy of Li is greater than that of Na.
Ionization energy of Li= 520 KJ/mol
Ionization energy of Na= 496 KJ/mol
2. In Periods
The shielding effect is consistent as you move in a period from left to right. However, the nuclear charge increases. The nucleus’s and the valence electron’s attraction grows stronger. So, Ionization energy increases when moving from left to right in a period. For example, the ionization energy of Be is less than B.
Ionization energy of C = I086 Kj/mol
Ionization energy of N = 1402 Kj/mol
Electron Affinity
The amount of energy released when an electron adds up in the valence shell of an isolated atom to produce a uni-negative gaseous ion is described as electron affinity. The formation of anion is explained by electron affinity.
X(g) + e- ⎯⎯→ X- (g) + electron affinity
Watch -> Electron affinity video lecture https://youtu.be/vpU-v18BDhY
Explanation
The energy change that occurs when an electron is added to a gaseous atom is known as its electron affinity. For example, the associated energy change is -328 kJ/mol when a fluorine atom gains an electron to create F – (g) in the gaseous state.
Variation of electron affinity in periodic table
1) Variation of electron affinity in a period
The electron affinity generally increases across a period as you move from left to right. This is because the nuclear charge increases and the atomic radius decreases, which attaches the extra electron more tightly to the nucleus. However, the shielding effect is constant throughout the period. As a result, in each period, halogens have the highest electron affinities and alkali metals the lowest.
2) Variation of electron affinity in a group
In a group, the electron affinity decreases from top to bottom. This is because the shielding effect has increased. Because of the increased shielding effect, the additional electron binds less tightly to the nucleus. As a result, less energy is released.
Electronegativity
An atom’s ability to attract electrons to itself in a chemical bond is known as electronegativity.
OR
Electronegativity is a chemical property that describes an atom’s ability to attract electrons to itself.
An atom’s electronegativity is influenced by both its atomic number and the separation of its valence electrons from its charged nucleus.
Watch -> Electronegativity lecture https://youtu.be/eqPg1o5ueO4
Variation of electronegativity in periodic table
1) Variation in electronegativity in a group
The electronegativity of an atom decreases when we move from top to bottom in a group this is because the atomic size increases while moving down the group and as a result shielding effect increases. So, the tendency of an atom to attract electrons decreases. E . g Electronegativity of Oxygen(3.5) is greater than that of Sulphur (2.5)
2) Variation in electronegativity in a period
The electronegativity of an atom increases when we move from left to right in a period because the nuclear charges increase from left to while the electrons are added in the same shell and these electrons can’t shield each other. So, the electrons are attracted toward the nucleus more effectively. E.g the electronegativity of Oxygen(3.5) is less than that of fluorine(4.0).
