Atomic Structure and Bonding in Materials

1. ATOMIC STRUCTURE & INTERATOMIC BONDING PART-1 DR. RABIUL HUSSAIN SCHOOL OF MATERIAL SCIENCE AND ENGINEERING JIMMA INSTITUTE OF TECHNOLOGY JIMMA UNIVERSITY, ETHIOPIA

2. Why study atomic structure and interatomic bonding?

3. Example:

5. ATOMIC STRUCTURE: Fundamental Concepts • Proton and electron, charged=1.60 x10-19 C • Mass of electron = 9.11x10-31 kg • Mass of protons and neutrons=1.67 x 10-27 kg • Atomic number(Z)=the number of protons • Atomic mass =mass of protons+neutrons • Isotope: Atoms of same element with different atomic masses (due to different numbers of neutrons) • Atomic mass unit(amu): 1amu=1/12TH of atomic mass of C • One mole = 6.023x1023 atoms or molecules(Avogadro’s Number)

6. Structure of an Atom

7. ATOMIC MODELS: How different sub-atomic particles fit in an ATOM

8. 400 BC-Greek Philosopher Democritus • Democritus coined the term “atomos” which means “uncuttable” or “the smallest indivisible particle of matter.” • Democritus postulated that atoms were completely solid, hard and small particles having no internal structure and has an infinite variety of shapes and sizes. • This theory was ignored for more than 2000 years.

9. Atomic Theory of Matter

10. Dalton’s Atomic Theory(1803): • John Dalton was a English instructor and natural philosopher. • “Each element is consist of a single unique type of atom and can join to form chemical compounds.” • He put forward the first theory of atom regarded as the BEGINNING OF MODERN ATOMIC THEORY.

11. Postulates of Daltons Atomic Theory

12. Postulates of Daltons Atomic Theory

13. Mendelev’s Classification of Elements(1869):

14. Mendelev’s Periodic Table

15. Today: Modern Periodic Table of Elements

16. Discovery of Electron by J.J.Thomson (1898-1903)

17. Thomson’s Atomic Model: PLUM PUDDING MODEL(1904) • J. J. Thomson postulated that an atom consisted of a diffuse cloud of positive charge with the negative electrons embedded randomly in it. This model, shown in Fig. is often called the plum pudding model because the electrons are like raisins dispersed in a pudding (the positive charge cloud), as in plum pudding, a favorite English dessert.

18. E. Rutherford’s Nuclear Atomic Model (1909-1911) • Tested and disproved the Plum Pudding model. • Carried out famous alpha (α) ray scattering experiment on Gold foil results of which led to Nuclear Model of Atom.

19. Rutherford’s α-ray Scattering Experiment Results: • Majority of particles transmitted (pass through) or deflected through small angles. • Tiny fraction deflected by large angles or bounced back through the original path.

20. Conclusion and Rutherford’s Nuclear Atomic Model • Disproved the plum pudding model. • Most of the space in an atom is empty. • Large amount of the atoms charge and mass is concentrated into a small region which he called nucleus. Fig. Rutherford’s Nuclear model of atom

21. How tiny is the Nucleus?

26. Bohr Atomic Model:

29. Schematic Representation of Bohr Atom • Bohr atomic model: • Electrons revolve around the atomic nucleus in discrete orbital and the position of any particular electron is more or less defined in terms of its orbitals. • The energies of electrons are quantized.

30. Limitations of Bohr Model: • It was unable to explain several phenomena involving electrons in solid. • So we need some other model to explain the behavior of solids. This leads us to QUANTUM MECHANICS and the model proposed is known as WAVE-MECHANICAL MODEL OF ATOM.

31. WAVE MECHANICAL MODEL To overcome the deficiencies in Bohr’s model • Electron is both wave-like and particle like. • Electron is no longer be considered as a particle in discrete orbital. • Electron can be at various locations in a discrete orbital (uncertainty of location).

34. Heisenberg’s Uncertainty Principle:

35. Schrodinger Wave Equation:

40. Quantum Mechanics & Atomic Orbitals • Electron density: Probability of finding electron at a point. • Higher density of dots in a region: Larger the values of ψ2.

42. QUANTUM NUMBERS • Every electron in an atom is characterized by 4 parameters called quantum numbers. • According to quantum mechanics, each electron in an atom is characterized by four different quantum numbers three of which (n, l and ml ) specify the wave function that gives the probability of finding the electron at various points of space around the nucleus. • These four quantum numbers are: • Principal quantum number n: distance of an electron from the nucleus • Second quantum number l: shape of the electron subshell • Third quantum number ml: number of energy states for each subshell • Forth quantum number ms: spin moment of an electron

43. Quantum numbers • Principal quantum number (n): n=1,2,3,4…[ shells=K,L,M,N….]. As n increases : • The orbital becomes larger ; • The electron spends more time farther from the nucleus • The electron has higher energy and therefore less tightly bound to the nucleus. • Azimuthal quantum number(l): l=0,1,2,3…(n-1) [subshells: s,p,d,f] • Defines the shape of the orbitals; • The value of l for a particular orbital is designated by s, p, d and f corresponding l values of 0,1,2,3 respectively. • Magnetic quantum number(ml): • Have integral values from l to –l including zero. • Describes the orientation of orbital in space. • Spin quantum number (ms ): • Spin moment of an electron either -1/2 0r +1/2.

44. Quantum numbers

45. Quantum Numbers

46. Relative energies of various shells and subshells

47. Sequence of filling electrons in orbitals: Aufbau Principle Governed by three principles: • Pauli’s Exclusion principle: No two electrons in an atom can have all four quantum numbers same. • Electrons have discrete energy states and an electron occupy the orbital from lower energy to higher energy. • Hund’s Rule: the lowest energy is attained by maximising the number of electrons with the same electron spin. [ For degenerate energy levels, each orbital is filled with one electron before electrons are paired up]

48. Electronic Configuration

49. THANK YOU FOR YOUR ATTENTION QUESTIONS?????