PERIODIC TABLE in 60 Minutes || Full Chapter Revision || Class 11th JEE

JEE Wallah・2 minutes read

The text discusses the properties of metalloids, metals, and non-metals, detailing their characteristics and classification in the periodic table. It also covers concepts such as effective nuclear charge, penetration power, atomic radius trends, ionization energy, metallic character, electron affinity, and electronegativity, providing a comprehensive overview of these fundamental principles in chemistry.

Insights

Metalloids exhibit properties of both metals and non-metals, positioned between these two categories in the periodic table.
The effective nuclear charge, influenced by the number of protons and shielding effect of inner electrons, impacts the size of atoms, with higher charges leading to smaller atomic radii.

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Recent questions

What are metalloids?
Elements with metal and non-metal characteristics.

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Summary

00:00

Element Classification and Atomic Structure Explained

Metalloids are elements that have characteristics of both metals and non-metals, lying on the border line between the two categories.
Alkali metals, alkaline earth metals, and other groups in the periodic table are classified as metals.
Non-metals are elements found in the P block of the periodic table, such as carbon, nitrogen, and sulfur.
Inert gases, like helium, are considered non-metals due to their properties.
Silicon is a metalloid, characterized by its green color.
Effective nuclear charge is the force of attraction on an electron by the nucleus, calculated as the difference between attractions and repulsions.
The ultimate cell in an atom is the outermost electron shell, while the penultimate cell is the second outermost shell.
The test electron in the ultimate cell experiences forces of attraction and repulsion, determining the effective nuclear charge.
The effective nuclear charge is influenced by the number of protons in the nucleus and the shielding effect of inner electrons.
Penetration power, also known as shielding effect, refers to the ability of an electron orbital to approach the nucleus, with S orbitals having higher penetration power than P orbitals.

13:51

Atomic Structure and Periodic Trends Explained

The nucleus attracts an electron, but if the orbital is outside the nucleus, the s orbital remains spherical.
The penetration power is crucial, as removing work from the nucleus can affect the force of attractions.
The shape of the p orbital is double with a dumbbell structure, representing a torn old curtain.
The force of attractions determines the filling of test electrons, with the p orbital having holes due to its torn structure.
The atomic radius decreases from left to right in the periodic table due to the increasing effective nuclear charge.
Moving from top to bottom in the periodic table results in a constant increase in atomic radius due to the addition of new shells.
The 4f series experiences a rapid decrease in atomic radius from left to right, known as the lanthanide contraction.
Lanthanide contraction occurs due to the poor shielding of the 4f orbital, leading to a significant decrease in atomic radius.
Ionic radius is influenced by the effective nuclear charge, with cations having smaller radii when the charge increases.
Removing an electron from an atom increases the effective nuclear charge, resulting in a decrease in ionic radius.

26:12

"Effective Nuclear Charge and Atomic Properties"

The text discusses the concept of effective nuclear charge and its impact on the size of atoms.
It explains that the size of an atom is determined by the effective nuclear charge, with higher charges resulting in smaller sizes.
The text delves into the concept of isoelectronic species, where atoms have the same number of electrons but different numbers of protons.
It highlights that the number of protons in an atom affects its effective nuclear charge and, consequently, its atomic radius.
The text emphasizes the relationship between ionization energy and effective nuclear charge, stating that as the charge increases, so does the energy required to remove an electron.
It details how ionization energy increases as one moves from left to right across the periodic table due to the rising effective nuclear charge.
The text explains that ionization energy decreases as one moves down a group in the periodic table because of the increasing atomic size.
It discusses the metallic character of elements, which refers to their ability to donate electrons easily.
The text explains that metallic character increases as one moves down a group in the periodic table due to decreasing ionization energy.
It concludes by discussing the concept of lanthanide contraction and its impact on ionization energy in the transition metals.

38:38

Electron Affinity: Key Factors and Trends

There will be 48 protests, with a specific focus on the electron's position in relation to protons.
The radius of electrons and protons is crucial in determining the force of attraction between them.
Diagonal relationships in elements like lithium and magnesium showcase similar characteristics due to their atomic radius.
The concept of electron affinity is explained as the energy released when an electron is added to an atom.
The electron affinity is influenced by factors like atomic size and effective nuclear charge.
Fluorine and chlorine's electron configurations impact their electron affinity and energy release.
The electron affinity is higher for chlorine compared to fluorine due to their atomic sizes and electron densities.
The P block elements exhibit varying electron affinities, with neon requiring significant energy to gain electrons.
Oxygen stands out in the halogen family for its high electron affinity, leading to energy release upon electron addition.
Iodine is highlighted as the weakest member of the halogen family, surpassing sulfur in electron affinity.

52:06

Understanding Electronegativity and Oxides

Electronegative elements have a strong nucleus that can pull electrons towards themselves, resulting in negative charges.
Electronegativity is measured on a scale where Fluorine has the highest value of 4.0, with values decreasing as you move left to right.
Remember the electronegativity scale: Fluorine (4.0), Oxygen (3.5), Nitrogen (3.0), Carbon (2.5), Hydrogen (2.1).
Oxides can be classified as amphoteric, acidic, or basic based on their reaction properties and oxidation states.
Basic oxides are typically found in metals, with oxidation states determining whether they are basic or amphoteric.
Non-metal oxides, like those of Boron, Carbon, Nitrogen, Phosphorus, and Chlorine, are always acidic, while metalloids' oxides are also acidic in nature.

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