Classification of Elements & Periodicity in Properties Class 11 Chemistry One Shot| NCERT Cha 3 CBSE
LearnoHub - Class 11, 12・82 minutes read
Understanding the periodic table involves the arrangement of elements based on properties and trends, with classification into s, p, d, and f blocks. Chemical reactivity is influenced by elements' physical properties and electron behavior, showcasing extremes in reactivity from alkali metals to halogens.
Insights
- Elements are grouped based on similar properties in the periodic table, aiding in recognizing trends rather than individual properties.
- Historical advancements in the periodic table, from Lavoisier's classification to Mendeleev's atomic mass arrangement to Moseley's modern atomic number-based organization, have shaped our understanding of elements.
- The periodic table provides a systematic way to classify elements into blocks (s, p, d, f) based on electronic configuration, guiding the arrangement of elements by atomic number and unveiling trends in physical properties like atomic radius and ionization enthalpy.
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Recent questions
What is the periodic table?
A chart organizing elements by properties.
How are elements classified?
Elements are grouped by similar properties.
What are magic numbers in chemistry?
Numbers indicating repeating properties.
How are electron configurations determined?
By filling orbitals based on energy levels.
What is the significance of electronegativity?
Reflects atom's ability to attract electrons.
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Summary
00:00
Chemistry: Elements, Properties, and Periodic Table
- Understanding chemistry involves comprehending the elements and their properties.
- Elements are categorized based on their properties, not just in a periodic table.
- Similar elements with comparable properties are grouped together.
- The periodic table aids in recognizing trends in properties rather than individual element properties.
- The history of the periodic table dates back to 1789 with Lavoisier's classification of elements.
- Mendeleev's periodic law arranged elements by increasing atomic mass.
- Henry Moseley's discovery led to the modern periodic law based on atomic number.
- The modern periodic table organizes elements by atomic number, showcasing periodic properties.
- Magic numbers like 2, 8, 18, 32 signify intervals where elements with similar properties repeat.
- By adding magic numbers to an element's atomic number, one can find the element below with similar properties.
16:36
Naming Elements: Discovery, Classification, and Nomenclature
- Discovery of element 104 by American and Soviet scientists
- Naming dilemma for the newly discovered element
- Establishment of IUPAC for systematic nomenclature
- Provision of temporary names until full confirmation of discovery
- IUPAC's code for assigning temporary names based on atomic numbers
- Examples of temporary names based on the code
- Official naming of elements once confirmed, such as Mendelvium, Nobelium, etc.
- Use of the periodic table to determine the position of elements
- Determining the period of an element based on the principal quantum number
- Classification of elements into blocks (s, p, d, f) based on electronic configuration
- Explanation of A block elements, exceptions like hydrogen and helium
- Explanation of P block elements and their electronic configurations
- Pattern of electronic configurations in P block elements
- Stability of noble gases due to fully filled orbitals
- Introduction to D block elements and the concept of inner d orbitals
- Example of titanium's electronic configuration and the filling of orbitals based on energy levels.
32:46
Electronic Configuration and Periodic Trends Explained
- Electrons are filled in the 4s orbital before the 3d orbital in electronic configuration.
- The last electron goes into the d orbital, making it a d block element with an outermost shell of n = 4.
- D block elements act as a bridge between s and p block elements, known as transition elements.
- F block elements are placed separately due to their unique electronic configuration, with the last electron going into the f orbital.
- The group number for s block elements is equal to the number of valence electrons.
- P block elements have a group number equal to 10 plus the number of electrons in the valence shell.
- D block elements' group number is the sum of electrons in the n-1d subshell and the valence shell.
- Metals are predominant in the periodic table, with over 78% of elements being metals.
- Metals are usually solid at room temperature, while nonmetals are typically gases or solids.
- Metallic character increases from top to bottom in a group and decreases from left to right in a period due to nuclear charge and electron acceptance tendencies.
48:37
Trends in Atomic and Ionic Radii
- Physical properties like electronegativity, electron gain enthalpy, and atomic radii show trends.
- Calculating atomic radius is complex due to the electron cloud nature and lack of fixed boundaries.
- Atomic radius decreases across a period due to increased nuclear charge and electron addition in the same shell.
- Moving down a group increases atomic radius as the number of shells and principal quantum number (n) rises.
- Ionic radii can be estimated by measuring distances between ions in ionic crystals.
- Ionic radii follow trends similar to atomic radii, with ions being smaller than atoms and larger than their atoms.
- Isoelectronic species have the same number of electrons but different atomic numbers, affecting their sizes.
- Ionization enthalpy is the energy required to remove an electron from an atom, with trends across periods and groups.
- Ionization enthalpy increases across a period due to rising nuclear charge and difficulty in removing electrons.
- Ionization enthalpy decreases down a group as electrons move further from the nucleus, making removal easier.
01:04:38
Nuclear Charges Influence Ionization Enthalpy and Electron Gain
- Boron has a nuclear charge of +5, while beryllium has a nuclear charge of +4.
- The ionization enthalpy is higher for the element with a higher nuclear charge.
- Beryllium's outer electron is in the S orbital, which attracts electrons more towards the nucleus compared to P electrons.
- Beryllium has a lower ionization enthalpy than Boron due to the spherical shape of S block electrons.
- Nitrogen's ionization enthalpy is lower than oxygen's due to electron-electron repulsion in oxygen's P orbital.
- Electron gain enthalpy refers to how easily an atom can gain an electron.
- Noble gases in Group 18 have positive electron gain enthalpy values due to their stability.
- Electron gain enthalpy becomes more negative as we move left to right across a period.
- Halogens like chlorine and fluorine have very high negative electron gain enthalpy values due to their electron-hungry nature.
- Electron gain enthalpy becomes less negative as we move down a group, except for oxygen and fluorine, which have less negative values than sulfur and chlorine due to electron-electron repulsion in smaller orbitals.
01:20:38
"Electronegativity Trends and Oxidation States Explained"
- Electronegativity of an element is not constant and varies depending on the atom it combines with.
- Electronegativity changes element to element, influenced by the atom it interacts with.
- An analogy is used to explain how electronegativity shifts based on the relative power of atoms.
- Electronegativity trends are discussed, showing an increase from left to right in a period due to nuclear charge.
- Electronegativity decreases when moving from top to bottom in a group due to increased distance between electrons and the nucleus.
- Trends in physical properties are explained, emphasizing the importance of observing left to right for periods and top to bottom for groups.
- Oxidation states are defined as the charge an atom acquires when forming a molecule with another atom.
- The method to determine oxidation states is detailed, based on the number of outermost electrons.
- Oxidation states for elements in different groups are calculated, with specific examples like lithium hydride and calcium hydride.
- The unique behavior of second period elements is attributed to their small size, high charge to radius ratio, and limited valence orbitals available for bonding.
01:35:38
Chemical Reactivity Trends in Periodic Table
- Periodic table reveals trends in chemical reactivity, with extremes on both ends being highly reactive, while the middle elements are less reactive due to electron behavior.
- Alkali metals on the left side are eager to donate electrons, making them highly reactive, while halogens on the right side are electron-hungry, also leading to high reactivity.
- Reactivity is evident in reactions with oxygen, where metals like sodium form basic oxides, halogens like chlorine form acidic oxides, and elements like aluminum produce amphoteric oxides, showcasing varied chemical behaviors.
