Class 11 -Structure of Atom / L-1/Learn basic to advanced Electronic configuration (easiest way) скачать в хорошем качестве

Class 11 -Structure of Atom / L-1/Learn basic to advanced Electronic configuration (easiest way)

how to write electronic configuration of element explain the electronic configuration of chromium atom explain the electronic configuration of copper atom what is ground state and excited state how to write electronic configuration of Uranium atom electronic configuration all important questions covered take your chemistry from 0 to Hero level easiest way to understand and do numericals on electronic configuration Electron Configuration he electron configuration of an atomic species (neutral or ionic) allows us to understand the shape and energy of its electrons. Many general rules are taken into consideration when assigning the "location" of the electron to its prospective energy state, however these assignments are arbitrary and it is always uncertain as to which electron is being described. Knowing the electron configuration of a species gives us a better understanding of its bonding ability, magnetism and other chemical properties. The electron configuration is the standard notation used to describe the electronic structure of an atom. Under the orbital approximation, we let each electron occupy an orbital, which can be solved by a single wavefunction. In doing so, we obtain three quantum numbers (n,l,ml), which are the same as the ones obtained from solving the Schrodinger's equation for Bohr's hydrogen atom. Hence, many of the rules that we use to describe the electron's address in the hydrogen atom can also be used in systems involving multiple electrons. When assigning electrons to orbitals, we must follow a set of three rules: the Aufbau Principle, the Pauli-Exclusion Principle, and Hund's Rule. The wavefunction is the solution to the Schrödinger equation. By solving the Schrödinger equation for the hydrogen atom, we obtain three quantum numbers, namely the principal quantum number (n), the orbital angular momentum quantum number (l), and the magnetic quantum number (ml). There is a fourth quantum number, called the spin magnetic quantum number (ms), which is not obtained from solving the Schrödinger equation. Together, these four quantum numbers can be used to describe the location of an electron in Bohr's hydrogen atom. These numbers can be thought of as an electron's "address" in the atom. Aufbau Principle The word 'Aufbau' is German for 'building up'. The Aufbau Principle, also called the building-up principle, states that electron's occupy orbitals in order of increasing energy. The order of occupation is as follows. This order of occupation roughly represents the increasing energy level of the orbitals. Hence, electrons occupy the orbitals in such a way that the energy is kept at a minimum. That is, the 7s, 5f, 6d, 7p subshells will not be filled with electrons unless the lower energy orbitals, 1s to 6p, are already fully occupied. Also, it is important to note that although the energy of the 3d orbital has been mathematically shown to be lower than that of the 4s orbital, electrons occupy the 4s orbital first before the 3d orbital. This observation can be ascribed to the fact that 3d electrons are more likely to be found closer to the nucleus; hence, they repel each other more strongly. Nonetheless, remembering the order of orbital energies, and hence assigning electrons to orbitals, can become rather easy when related to the periodic table. Hund's Rule Hund's Rule states that when electrons occupy degenerate orbitals (i.e. same n and l quantum numbers), they must first occupy the empty orbitals before double occupying them. Furthermore, the most stable configuration results when the spins are parallel (i.e. all alpha electrons or all beta electrons). Nitrogen, for example, has 3 electrons occupying the 2p orbital. According to Hund's Rule, they must first occupy each of the three degenerate p orbitals, namely the 2px orbital, 2py orbital, and the 2pz orbital, and with parallel spins Pauli-Exclusion Principle Wolfgang Pauli postulated that each electron can be described with a unique set of four quantum numbers. Therefore, if two electrons occupy the same orbital, such as the 3s orbital, their spins must be paired. Although they have the same principal quantum number (n=3), the same orbital angular momentum quantum number (l=0), and the same magnetic quantum number (ml=0), they have different spin magnetic quantum numbers (ms=+1/2 and ms=-1/2). List of anomalous electronic configurations are: Chromium :- [Ar] 3d5 4s1 Copper :- [Ar] 3d10 4s1 Niobium :- [Kr] 4d4 5s1 Molybdenum :— [Kr] 4d5 5s1 Ruthenium :— [Kr] 4d7 5s1 Rhodium :— [Kr] 4d8 5s1 Palladium :- [Kr] 4d10 5s0 Silver :— [Kr] 4d10 5s1 Lanthanum :— [Xe] 5d1 6s2 Cerium :— [Xe] 4f1 5d1 6s2 Gadolinium :— [Xe] 4f7 5d1 6s2 Platinum :— [Xe] 4f14 5d9 6s1 Gold :— [Xe] 4f14 5d10 6s1 Actinium :— [Rn] 6d1 7s2 Thorium:— [Rn] 6d2 7s2 Protactinium :— [Rn] 5f2 6d1 7s2 Uranium :— [Rn] 5f3 6d1 7s2 Neptunium :— [Rn] 5f4 6d1 7s2 Curium :— [Rn] 5f7 6d1 7s2