Chapter02 - Applied Science Department (ASD) Centre for...

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PCE 0015 Chemistry for Engineers Foundation in Engineering ONLINE NOTES Chapter 2 QUANTUM THEORY AND ATOMIC STRUCTURE FOSEE , MULTIMEDIA UNIVERSITY (436821-T) MELAKA CAMPUS, JALAN AYER KEROH LAMA, 75450 MELAKA, MALAYSIA. Tel 606 252 3594 Fax 606 231 8799 URL: http://fosee.mmu.edu.my/~asd/ Applied Science Department (ASD) Centre for Foundation Studies and Extension Education (FOSEE)
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PCE0015 Chemistry For Engineers Chapter 2 2.1 Quantum Theory 2.2 Bohr’s Theory of the Hydrogen Atom 2.3 Duality of Electron 2.4 Quantum Number 2.5 Atomic Orbitals 2.6 Electron Configuration 2.7 Aufbau’s Principle Upon completion of this chapter, you should be able to : 1. Understand the quantum theory, which states that radiant energy is emitted by atoms and molecules in small discrete amounts, rather than over a continuous range. This behavior is governed by the relationship E = h ν . 2. Explain Bohr’s theory of the hydrogen atom 3. Express the dual nature of the electron. De Broglie extended Einstein’s wave-particle description of light to all matter in motion. The wavelength of a moving particles of mass m and velocity u is given by the De Broglie equation λ = h / mu. 4. Understand four quantum number, namely: the principal quantum number n identifies the main energy level, or shell, of the orbital; the angular momentum quantum number l indicates the shape of the orbital; the magnetic quantum number m l specifies the orientation of the orbital in space; and the electron spin quantum number m s , indicates the direction of the electron’s spin on its own axis. 5. Differentiate atomic orbitals, which consists of orbital s , p , d , …and so on. 6. Express electron configuration by using Pauli exclusion principle and Hund’s rules. 7. Understand the Aufbau principle, which provides the guideline for building up the element. __________________________________________________________________________________ HST/MAM 2/ 11
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PCE0015 Chemistry For Engineers Chapter 2 2.1 QUANTUM THEORY Classical physics assumed that atoms and molecules could emit (or absorb) any arbitrary amount of radiant energy. However, according to Planck, atoms and molecules could emit (or absorb) energy only in discrete quantities. Planck gave the name quantum to the smallest quantity of energy that can be emitted (or absorbed) in the form of electromagnetic radiation . The energy E of a single quantum of energy is given by E = h ν Where h is called Planck’s constant and ν is the frequency of radiation. The value of Planck’s constant is 6.63 X 10 -34 J s. According to the quantum theory, energy is always emitted in multiples of h ν , 2 h ν , 3 h ν , ….., but never, for example, 1.67 h ν or 4.98 h ν . At the time Planck presented his theory, he could not explain why energies should be foxed or quantized in this manner. Starting with his hypothesis, however, he had no trouble correlating the experimental data for emission by solids over the entire range of wavelength, they all supported the quantum
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Chapter02 - Applied Science Department (ASD) Centre for...

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