Periodic Table: COMPLETE Chapter in 1 Video | Quick Revision | Class 11 Arjuna JEE
Arjuna JEE・2 minutes read
The periodic table is essential in atomic chemistry for classifying known elements, with Mendeleev's table being a favorite JE Mains question, focusing on isoelectronic species and shield constant for periodic trends. Classification of elements in groups by similar properties is key, with Mendeleev's table initially predicting new elements accurately but having limitations like hydrogen's placement and grouping similar elements separately.
Insights
- The periodic table is designed to classify known elements based on their properties, with Mendeleev's table being a significant historical breakthrough in predicting new elements accurately.
- The modern periodic table organizes elements based on atomic number, with elements in the same group sharing similar properties due to their outermost shell's electronic configuration.
- Electronegativity values, which measure an atom's tendency to attract electrons, vary based on factors like atomic radius and effective nuclear charge, impacting bond energy and the nature of chemical bonds.
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Recent questions
What is the significance of the periodic table?
The periodic table classifies known elements based on properties.
How did Mendeleev contribute to the periodic table?
Mendeleev's table predicted new elements accurately.
What is the shielding effect in chemistry?
Shielding effect involves inner electrons repelling outer electrons.
How does electronegativity influence chemical bonds?
Electronegativity affects bond energy and charge distribution.
How do oxides vary across periods and groups?
Oxides' nature changes based on metallic or non-metallic character.
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Summary
00:00
"Periodic Table: Classification, Trends, and History"
- The periodic table is crucial in aortic chemistry, designed to classify known elements.
- Mendeleev's periodic table initially included 63 known elements and has been a favorite JE Mains question.
- Isoelectronic species and shield constant are key concepts in periodic trends.
- Z effective increases left to right in a period, affecting size and effectiveness.
- The session aims to revise the chapter on elements and periodic properties thoroughly.
- Classification of elements in groups with similar properties is the essence of the periodic table.
- Historical attempts at classification include Lavoisier, Proust, and Newland's Law of Octaves.
- Mendeleev's periodic table was a significant breakthrough, predicting new elements accurately.
- Mendeleev's periodic law states that elements' properties depend on their atomic mass.
- Mendeleev's table had some limitations, like the placement of hydrogen and isotopes, and grouping similar elements separately.
13:22
Evolution of Periodic Table and Elements
- The atomic mass of an element was less, leading to the identification of four groups: Cobalt, nickel, argon, and potassium.
- Mendeleev placed elements with different properties in the same group, like copper and sodium, due to similar reactivity.
- Mendeleev's periodic table was discarded in favor of the periodic law, which led to the modern periodic table design.
- Mosle's study revealed that atomic number, not atomic mass, determined an element's properties.
- The modern periodic table organizes elements based on atomic number, with 18 groups and seven periods.
- Elements in the same group share similar properties due to their outermost shell's electronic configuration.
- Periods in the periodic table indicate the number of electron shells in an element, with specific configurations for each period.
- S block elements have an electronic configuration of ns1 or ns2, while p block elements range from ns2 np1 to np6.
- D block elements have n-1d 1 to 10 ns0 to 2 configurations and are crucial in the periodic table.
- F block elements, like actinides and lanthanides, follow specific electronic configurations and are essential in classification.
27:02
Valence electrons determine group number in elements
- Valence electrons in p block elements determine group number
- Group number in p block element = 10 + number of valence electrons
- Group number for 1, 3, 4, and 5 valence electrons in p block elements are 13, 3, 14, and 15 respectively
- For d block elements, group number is determined by adding electrons in (n-1)d and n s orbitals
- F block elements have undefined group numbers
- Shielding effect explained as inner electrons repelling outer electrons
- Screening effect described as inner electrons acting as a shield for outer electrons
- Shielding effect order: s > p > d > f
- Calculation of shielding constant involves electronic configuration and group classification
- Ionic radius is smaller for cations and larger for anions due to effective nuclear charge and electron repulsion
41:23
Element Size Trends in the Periodic Table
- The radius of elements is highest during their respective periods, with neon having the maximum radius.
- Neon and argon are abundant in the second period, not the third.
- The size of elements in the third period decreases from left to right, except for chlorine.
- Potassium is larger than chromium, and rubidium is larger than noon.
- Poor shielding of D&F electrons starts after the third period, affecting size.
- Group number one elements increase in size down the group.
- Group 13 elements follow a specific order in size: aluminum, gallium, indium, and thallium.
- Gallium is smaller than aluminum due to poor shielding of 3D electrons.
- Lanthanide contraction affects the size of 4D and 5D series elements.
- Ionization energy increases with successive removal of electrons, influenced by factors like atomic radius, effective nuclear charge, penetration effect, and stability of half and fully filled orbitals.
54:00
Ionization Energy Trends in Periodic Table
- Electrons are the same as gallium and aluminium.
- Lafra was the first size ordered, encountering the same ionization energy issue.
- Group 14 elements include carbon, silicon, germanium, tin, and lead.
- Ionization energy exceeds that of tin due to poor shielding of electrons on lead's outer electrons.
- Nucleus attraction on lead's outer electrons is more than expected, making electron removal difficult.
- Group 2 generally has higher ion energy than Group 13, with beryllium's ion energy exceeding boron's.
- The trend of Group 2 having higher ion energy than Group 13 is not applicable in the second period but in the third and fourth periods.
- Poor shielding of D&F electrons causes Munna to have more helium than indium stones.
- Group 15 has higher ionization energy than Group 16, with the trend applicable in the second, third, and fourth periods.
- The order of 3d, 4d, and 5d elements varies, with exceptions like chromium, vanadium, and nickel in Group 5, 6, and 10.
01:07:00
Understanding Electronegativity and Periodic Trends
- Electronegativity is the tendency of an atom to pull electrons towards itself, creating partial positive and negative charges in a bond.
- Electronegativity values are not fixed and depend on the atoms an element is bonded with.
- Different scientists developed electronegativity scales like the Pauling scale and the Mulliken scale to measure this tendency.
- The difference in electronegativity between two atoms affects the bond energy and can be calculated using specific formulas.
- Trends in electronegativity show that values increase from left to right and decrease down a group in the periodic table.
- Factors like atomic radius, effective nuclear charge, and hybridization influence electronegativity values.
- Applications of electronegativity include determining the ionic character of a bond, bond length, and the nature of hydroxides.
- The valency of an element depends on the number of valence electrons and varies across periods and groups in the periodic table.
- Diagonal relationships in the periodic table show similarities in properties between elements in the same group and their diagonally aligned counterparts.
- The nature of oxides, whether acidic, basic, or neutral, is influenced by the metallic or non-metallic character of the element and changes across periods and groups.
01:19:58
Amphoteric Oxides and D Block Elements
- Amphoteric oxides include Zn, Be, Aluminum, Gallium, Tin, and Lead. To remember them, associate them with the phrase "Sir, dear saint, remember you in this way. Sir, you can keep a nice cow, it is a saint, isn't it?" Arsenic and antimony with a +3 oxidation state are also amphoteric, while those with a +5 charge are acidic.
- D block elements with oxidation states of +1 or +2 are basic, except for Cr2O3, which is amphoteric. V2O5 is amphoteric in nature. The session covered the nature of oxides extensively, with exceptions like Cr2O3 and V2O5.
