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Understanding Atomic Structure: Quantum Numbers and Electronic Configurations

Explore the principles of atomic structure, quantum mechanics, and electronic configurations. Learn about elements, isotopes, atomic orbitals, and quantum numbers.

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Understanding Atomic Structure: Quantum Numbers and Electronic Configurations

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  1. LECTURE 2

  2. Lecture 2 Atomic Structure

  3. Plan 1. Quantum numbers 2. The principles of building of electronic structures 3. The periodic system of the elements

  4. Atom • An atomic nucleus consist of protons and neutrons • The number of protons in all atoms of the same element remain constant and is called atomic number which equals the number of electrons in an atom, while the number of neutrons differ in different isotopes. • Any atom is electroneutral

  5. Properties of Elemental Particles ParticleStateChargeMass (у.е.) Proton (p) core +1 1,00728 Neutron (n) core 0 1,00867 Positron (е) core+1 0,00055 Electron(е) shell -1 0,00055

  6. ATOM They are specified the chemical properties They are specified the mass of an atom NUCLEUS ELECTRONS PROTONS & NEUTRONS NUCLONS

  7. A = Z + N A – an atomic mass Z – A charge of a nucleus (the number of protons – atomic number) N – the number of neutrons Э А Z

  8. Isotopes of Chlorine 35 17 37 17 35 75,43 + 37  24,57 100 Cl(75,43%) Cl(24,57%) Ar = = 35,491

  9. Electronic Structure of an Atom

  10. The Quantum Mechanics Base • The moving of an electron hasambivalentnature: it is a particle as well as a wave at the same time • Therefore three principles of quantum mechanic have appeared

  11. QuantizationPrinciple (М. Plank, 1900) • Atoms emit the energy by doses which are multiple to the smallest magnitude named quantum • photon - h h= 6,626•10-34(J•s)– Plank constant =c Е = h

  12. Duality Principle • Under free moving the electrons have wavelike character • If the electrons interact with a matter they show the properties of particle • An electron has both properties at the same time

  13. The Uncertainly Principle(W. Heisenberg, 1927) • It is impossible to know simultaneously both the position and momentum of a small particle, such as an electron • A position and a momentum of an electron can be found only with some probability

  14. The state of an electron is determined by the probability its finding in definite region of space • The part of the atomic space where the probability of finding an electron is maximum (equal to 90%) is called an atomic orbital (AO)

  15. The stochastic model

  16. Orbitaldz2

  17. Schrodinger’s equation -the equation ofthree-dimensionalwave • It bonds energy, coordinates and wave function (). • The square of the wave function -2 determines the probability of finding the electron at a particular point

  18. The Quantum Numbers • Schrodinger’s equationis three-dimensional • Therefore there are three sets of quantum numbers to solve this equation • The size (energy) of an electronic cloud, its form and orientation change by the certain doses (quantums)

  19. Principle Quantum Number (n) • n- 1,2,3,….It characterizes energy (E) of a level, namely, the total energy of an electron on the orbital and the average distance from the nucleus • The energetic levelis the state of electrons in an atom with certain value of the principle quantum number • The main state of an atom has minimum energy • The excited state of an atom has higher values of electrons energy

  20. The state of Carbon Atom The main state of an atom The exited state of an atom

  21. Each principle level includes one or more sublevels, that is, the orbitals (electronic clouds) with different geometric forms • Combination of orbitals having the same value of principle quantum number is an energy level

  22. The Orbital Quantum Number (l) characterizes the form of an electronic cloud (an orbital) l= 0, 1, 2, 3…….n-1 Sublevel:s, p, d, f, g, h That is, the energetic level (n) is the collection of energetic sublevels differing in energy

  23. The Types of Atomic Orbitals S Px,Py,Pz dxz,dxy,dz2 dx2-y2,dyz

  24. The magnetic quantum number(ml)characterizes the direction of the electronic cloud (orbital) in space surrounding the nucleus • mlchange from–lto+l,and overall =2l + 1values • For example: l= 0(s); ml = 0 l = 1(p); ml = 0, +1, -1

  25. The spin number(ms)characterizes own magnetic moment of an electron that corresponds to those of a charged particle spinning on its axis. • Either of two spin is possible, clockwise or counterclockwise. • This quantum number is not related to n, l ml. • mshas values: +1/2or-1/2

  26. An Atomic Orbital (АО) • This is the state of an electron in an atom which is described by wave function with the set of three quantum numbers n, l, ml For example: 1s (n = 1, l = 0, ml = 0), 2p (n = 2, l = 1, ml = -1, 0, +1)

  27. Building the Electron Configuration of an Atom • The lowest energy principle: an electron fills АО with minimum energy • Pauli exclusion principle: no two electrons in an atom which can have the same set of four quantum numbers • Hund’s rule: the summary spin of electrons into one sublevel is maximum

  28. Klechkovsy’s rule • It defines the distribution of orbitals according to energy • An orbital having less sum of n+l has lower energy. • If two orbitals have equal sums of n+l then the less energy will be assigned to the orbital where n is less.

  29. Order of the Аtomical Orbital’s Filling 1s2s2p3s3p4s3d4p5s4d5p 6s4f5d6p7s5f6d7p

  30. Maximum filling of a sublevel by the electrons l - 2(2l+1)e s – 2e p – 6e d – 10e f – 14e

  31. Manner of Description of Electronic Structures • Electronic formula • Graphic structure • Energetic diagram

  32. Examples of Electronic Structures The full electronic formula Se - 1s22s22p63s23p64s23d104p4 The short formulaof Se - 4s24p4 Electronic-graphic structure S p d 4 3 2 1 Ti

  33. Energetic Diagram of Vanadium f p d • Е S 5 4 3 2 1

  34. The Promotion of an Electron For example: z = 24; Cr We are waiting electonic formula:1s22s22p63s23p64s23d4 realy:1s22s22p63s23p64s13d5

  35. Higher Energetic Stability • p6 d10f14 • p3 d5f7

  36. The Periodical Law & The Periodical System

  37. The Periodic Law • The properties of elements are a periodic function of their atomic numbers

  38. There are 4 types of the elements in the Periodical System : s, p, d, f – elements • It depends on the energy sublevel which fills last • On this sublevel the electrons are called valence electrons and they take part to form chemical bond

  39. s – elements: ns1 & ns2 p – elements: ns2np16 d – elements: ns2(n-1)d1  10 f – elements: ns2(n-2)f1  14

  40. A period is a horizontal row of elements. The atoms of every row have equal number of energetic levels which partially or fully filling by electrons

  41. Short periods • 1 period (n=1): 2 elements (1s1&1s2) • 2 period (n=2): 8 elements (2s12pо - 2s22p6) • 3 period (n=3): 8 elements (3s13pо - 3s23p6)

  42. Long Periods • 4 period(n=4): 18 elements(4s13dо4pо - 4s23d104p6) • 5 period(n=5): 18 elements(5s14d05pо - 5s24d105p6) • 6 period(n=6): 32 elements(6s14fо5d106p6 - 6s24f145d106p6) • 7 period(n=7): 32 elements7s25fо6d107p6………..

  43. A group is a vertical column of elements. • They have the same electronic configuration of atoms • They have equal number of outmost electrons • They have the same max valance and similar chemical properties

  44. Periodicity of the Properties of Elements • Atomic and ionic radii • Ionization energy • Electroaffinity • Electronegativity • Valence of the elements

  45. Atomic and ionic Radii of Elements • The radius of an atom (or ion) is the distance from nucleus to maximum of the electronic density of the outmost orbital of this atom • In periods the radii of atoms decrease • In groups the radii of atoms increase

  46. Radii of Cations & Anions • The transformation of an atom into cation gives the sharp decreasing of orbital radius • The transformation of an atom into anion almost does not change orbital radius Rкат<Rат < Rан Cl+ < Cl < Cl–

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