TLDR;
This video provides a comprehensive overview of the periodic table and its properties, focusing on electronic configuration, atomic size, ionisation energy, and electron affinity. It explains key principles like the Aufbau principle and effective nuclear charge, and also covers exceptions and trends within the periodic table.
- Electronic configuration and the Aufbau principle are explained.
- Trends in atomic size, ionisation energy, and electron affinity are discussed.
- Specific anomalies and exceptions, such as lanthanide contraction and the behaviour of certain elements within the Boron family, are highlighted.
Introduction [0:00]
The JEE Wallah channel welcomes students to the "Manzil" batch, aimed at achieving a good percentile in JEE Mains and securing an NIT. The goal is to prepare students thoroughly for the January attempt, with the batch designed to provide conceptual clarity and a concrete understanding of the material. The instructor assures that the session will change the way students view inorganic chemistry, regardless of their prior knowledge.
Topics to be covered [6:55]
The session will cover topics including atoms, shells, subshells, orbitals, the Aufbau principle, electronic configuration, the modern periodic table, IUPAC naming, atomic size, ionisation energy, electron affinity, electronegativity, Moseley's equation, and the nature of oxides. The aim is to solve PYQs.
Atom, Shell, Subshell [9:46]
An atom is defined as the smallest particle, comprising a nucleus with protons and neutrons, around which electrons revolve in shells. The number of protons in the nucleus is the atomic number (Z), which equals the number of electrons in a neutral atom. Shells are numbered (1, 2, 3, 4), corresponding to principal quantum numbers. The outermost shell is the ultimate shell, while the one inside it is the penultimate shell, and the one before that is the anti-penultimate shell. Each shell contains subshells: the first shell has only 's', the second has 's' and 'p', the third has 's', 'p', and 'd', and the fourth has 's', 'p', 'd', and 'f'. The number of orbitals and maximum electrons in each subshell are: s (1 orbital, 2 electrons), p (3 orbitals, 6 electrons), d (5 orbitals, 10 electrons), and f (7 orbitals, 14 electrons).
Aufbau Principle [18:40]
The Aufbau principle dictates that electrons first fill the subshells with the lowest energy. The energy sequence is: s, s p, s p, s d p, s d p, s f d p, s f d p. When writing electronic configurations, the principal quantum number for 'd' is one less than that of 's' and 'p', and for 'f' it is two less. The filling order is 1s, 2s, 2p, 3s, 3p, 4s, 3d, 4p, 5s, 4d, 5p, 6s, 4f, 5d, 6p, 7s, 5f, 6d, 7p. The maximum principal quantum number indicates the period number, and the maximum number of electrons in the energy sequence indicates the number of elements in that period.
Atomic Number Distribution [49:48]
The periodic table is divided into four blocks: s, p, d, and f. The s-block has two vertical columns, the p-block has six, the d-block has ten, and the f-block appears to have 14. The number of vertical columns corresponds to the maximum number of electrons each subshell can hold. The first period fills 1s, the second fills 2s and 2p, and so on. The 3d series elements are those where the 3d subshell is being filled. The 4f and 5f series are placed separately at the bottom of the table. The IUPAC numbering of groups is 1 to 18. The s-block consists of groups 1 and 2, the p-block spans from 13 to 18, and the d-block ranges from 3 to 12. The f-block elements are located within group 3.
Lanthanide Series [1:03:15]
The lanthanide series includes elements from Cerium (58) to Lutetium (71), belonging to the sixth period and group 3. The actinide series includes elements from Thorium (90) to Lawrencium (103), belonging to the seventh period and group 3. Elements 101-103 are part of the 5f series. Elements 104 onwards follow the group number corresponding to the last digit of their atomic number (e.g., element 104 belongs to group 4). Main group elements are s and p block elements, excluding inert gases. Alkali metals are group 1 elements, and alkaline earth metals are group 2 elements. The nitrogen family is group 15, chalcogens are group 16, and halogens are group 17. Transition elements are d-block elements, excluding group 12. Inner transition elements are f-block elements.
Important [1:23:20]
Elements 119 and 120, if discovered, would belong to the eighth period and the s-block. Element 119 would be in group 1 (alkali metal), and element 120 would be in group 2 (alkaline earth metal).
IUPAC name of 101 to 118 [1:25:36]
IUPAC nomenclature uses numerical roots for atomic numbers: 0 (nil), 1 (un), 2 (bi), 3 (tri), 4 (quad), 5 (pent), 6 (hex), 7 (sept), 8 (oct), 9 (enn). The name is formed by combining these roots and adding "-ium" at the end. For example, element 101 is Unnilunium (Unu).
Name of Elements from 101 to 118 [1:32:30]
Mendelevium (Md, 101), Nobelium (No, 102), Lawrencium (Lr, 103), Rutherfordium (Rf, 104), Dubnium (Db, 105), Seaborgium (Sg, 106), Bohrium (Bh, 107), Hassium (Hs, 108), Meitnerium (Mt, 109), Darmstadtium (Ds, 110), Roentgenium (Rg, 111), Copernicium (Cn, 112), Nihonium (Nh, 113), Flerovium (Fl, 114), Moscovium (Mc, 115), Livermorium (Lv, 116), Tennessine (Ts, 117), and Oganesson (Og, 118).
Electronic configuration [1:42:10]
Electronic configurations can be derived using the Aufbau principle sequence (s, s p, s p, s d p, s d p, s f d p, s f d p). Elements in the second period have configurations starting with Helium (He), third-period elements start with Neon (Ne), and so on. S-block elements have configurations ending in 's', p-block elements end in 'p', and d-block elements end in 'd'.
No of Electrons in Valence Shell [1:55:02]
Group 1 elements have one valence electron, group 2 has two, group 13 has three, group 14 has four, group 15 has five, group 16 has six, group 17 has seven, and group 18 typically has eight (except Helium, which has two). The first transition series (3d) includes elements from Scandium (21) to Zinc (30), the second (4d) from Yttrium (39) to Cadmium (48), the third (5d) from Lanthanum (57) followed by Hafnium (72) to Mercury (80), and the fourth (6d) from Actinium (89) followed by Rutherfordium (104) to Copernicium (112).
Questions [2:19:52]
To determine the group, period, block, and family of an element with a given atomic number, first, identify the period by comparing the atomic number to the range of atomic numbers for each period. Then, determine the group and block based on the element's position in the periodic table.
Transition Metals [2:26:16]
Transition metals are d-block elements that have partially filled d orbitals in their atomic or ionic states (excluding group 12 elements like Zn, Cd, and Hg).
Metals, Non-Metals & Metalloids [2:29:54]
Metals tend to lose electrons, non-metals gain electrons, and metalloids exhibit properties of both. Metalloids include elements such as Arsenic (As), Antimony (Sb), Tellurium (Te), and Germanium (Ge). Metals comprise the s-block, d-block, and f-block, along with some p-block elements.
Effective Nuclear Charge [2:36:58]
Effective nuclear charge (Zeff) is the net positive charge experienced by an electron in a multi-electron atom. It is calculated as Zeff = Z - σ, where Z is the atomic number and σ is the shielding constant. The shielding constant depends on the number of electrons in inner shells.
Penetration Effect [2:44:30]
The penetration effect refers to the ability of an electron to penetrate the inner electron shells and get closer to the nucleus. The order of penetration is ns > np > nd > nf. Subshells closer to the nucleus shield the outer electrons more effectively.
Atomic Radius [2:53:32]
Atomic radius is the distance between the centre of the nucleus and the outermost electron shell. It cannot be precisely determined for an individual atom due to the lack of a sharp boundary. Instead, it is estimated by measuring the distance between atoms in a combined state. Covalent radius is half the distance between two nuclei in a covalently bonded molecule. Metallic radius is half the distance between two adjacent metal ions in a metallic lattice. Van der Waals radius is half the distance between two non-bonded atoms in adjacent molecules.
Variation of Atomic size in the Periodic Table [3:02:05]
In a period, atomic size generally decreases from left to right due to increasing nuclear charge. In a group, atomic size increases from top to bottom due to the addition of new electron shells.
Ionic Radius [3:12:40]
Cations are smaller than their parent atoms because the removal of electrons increases the effective nuclear charge. Anions are larger than their parent atoms because the addition of electrons decreases the effective nuclear charge.
Isoelectronic species [3:19:16]
Isoelectronic species have the same number of electrons. In an isoelectronic series, the species with the greater nuclear charge will have a smaller radius.
Questions [3:26:24]
Atomic and ionic radii can be compared by considering the balance between nuclear charge and the number of electrons.
Boron Family [3:31:12]
In the Boron family, the expected trend of increasing atomic size down the group is disrupted between aluminium (Al) and gallium (Ga). Gallium is smaller than aluminium due to the presence of the 3d subshell, which has poor shielding effect, leading to a greater effective nuclear charge.
Lanthanide & Actinide Contraction [3:38:20]
Lanthanide contraction is the steady decrease in the size of the lanthanide ions (La to Lu) with increasing atomic number. This is due to the poor shielding of the 4f electrons, which causes the effective nuclear charge to increase, pulling the electron cloud inward.
Effect of Lanthanide Contraction on D Block [3:44:06]
Lanthanide contraction affects the size of the elements in the d-block. The sizes of the 4d and 5d series become similar due to lanthanide contraction.
3d Series [3:51:28]
In the 3d series, atomic size generally decreases from Scandium (Sc) to Chromium (Cr) due to increasing nuclear charge. From Manganese (Mn) to Nickel (Ni), the size remains nearly constant due to a balance between increasing nuclear charge and inter-electronic repulsion. From Copper (Cu) to Zinc (Zn), the size slightly increases due to increasing inter-electronic repulsion.
Ionisation Energy [4:02:45]
Ionisation energy (IE) is the minimum energy required to remove an electron from an isolated gaseous atom in its ground state. Successive ionisation energies increase (IE1 < IE2 < IE3) because it becomes more difficult to remove an electron from an increasingly positive ion.
Factors Affecting I.E [4:12:36]
Factors affecting ionisation energy include effective nuclear charge (Zeff), atomic size, penetration effect, and stable electronic configurations. Higher effective nuclear charge increases ionisation energy, while larger atomic size decreases it. The penetration effect makes it more difficult to remove electrons from s orbitals compared to p orbitals. Half-filled and fully filled electronic configurations are more stable and require more energy to remove an electron.
Questions [4:43:52]
The largest difference in successive ionisation energies indicates the number of valence electrons.
Boron Family (2) [4:50:12]
In the Boron family, the ionisation energy trend is B > Tl > Ga > Al > In. This trend is influenced by the presence of the 3d subshell in Gallium and the 4f subshell in Thallium, which have poor shielding effects.
Electron Gain Enthalpy [5:10:20]
Electron gain enthalpy (ΔHeg) is the energy change when an electron is added to an isolated gaseous atom. If energy is released, the process is exothermic (ΔHeg is negative). If energy is required, the process is endothermic (ΔHeg is positive).
Electron Affinity [5:21:20]
Electron affinity is the tendency of an atom to attract an additional electron. It is related to electron gain enthalpy. More negative electron gain enthalpy indicates a higher electron affinity.
Electronic Configuration (2) [5:30:22]
Stable electronic configurations, such as half-filled and fully filled subshells, resist the addition of an electron.
Penetration Effect [5:35:58]
The penetration effect influences electron gain enthalpy. Electrons are more readily added to orbitals closer to the nucleus.
UQT Points [5:38:48]
All cations have negative electron gain enthalpies because they readily accept electrons. Generally, neutral atoms have negative electron gain enthalpies, but there are exceptions.
Electron Affinity of F & Cl [5:49:12]
Chlorine has a more negative electron gain enthalpy than fluorine due to its larger size and lower electron density.
Important [6:04:30]
Second electron gain enthalpies are always positive because adding an electron to an already negative ion requires energy.
Relation between IE & EGE [6:08:00]
The ionisation energy of an anion is equal in magnitude but opposite in sign to the electron gain enthalpy of the corresponding neutral atom.
Electronegativity [6:16:45]
Electronegativity is the ability of an atom in a chemical compound to attract shared electrons to itself.
Pauling Scale [6:23:17]
The Pauling scale is a measure of electronegativity, with fluorine having the highest value (4.0). Electronegativity generally increases across a period and decreases down a group.
Diagonal Relationship [6:31:56]
Diagonal relationships occur between certain elements in adjacent groups and periods (e.g., Lithium (Li) and Magnesium (Mg)). These elements exhibit similar properties due to the balance of increasing charge and decreasing size.
Oxidation state in Oxides [6:35:05]
Oxides are compounds containing oxygen. The oxidation state of oxygen in oxides is typically -2.
Acidic Nature of Oxides [6:45:51]
Metallic oxides are generally basic, while non-metallic oxides are acidic. Amphoteric oxides can react with both acids and bases.
Modern Periodic Table [6:49:45]
Henry Moseley's experiments with X-rays showed that the atomic number is the fundamental property of an element, leading to the modern periodic table. Moseley's equation relates the frequency of X-rays emitted by an element to its atomic number.
Questions [6:52:54]
The periodic table's structure and properties can be used to predict the behaviour of elements and their compounds.
Thank you [6:56:10]
The instructor concludes the session, encouraging students to attend the live test and utilise the resources provided by Physics Wallah.
