Lab_2_manual - UOIT Chem. 1800, W10: Exp. 2-7.9 2. ATOMIC...

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UOIT Chem. 1800, W10: Exp. 2-7.9 52 2. ATOMIC EMISSION SPECTRA Objectives 1) To observe the atomic emission spectra of two elements: mercury and hydrogen; 2) to calculate the Rydberg constant from the spectrum of the hydrogen atom; 3) to observe the spectra for several cations and note that cations give characteristic colours to flames; 4) to understand spectra are a valuable tool for identifying chemical species. Introduction In the late 19 th century Johann Balmer observed that when electric current was passed through hydrogen gas, the gas emitted light. When the gas was viewed through a spectroscope (see below), he observed discrete, bright lines against a dark background in the visible region of the electromagnetic spectrum. The discrete lines indicate that the hydrogen atom must exist in discrete energy levels. Otherwise, the emitted light would resemble a rainbow. He found he could relate the wavelengths ( ± ) of the lines to a constant, k, using the relationship: (2.1) 11 2 1 22 λ =− ± ² ³ ´ µ k n where ‘n’ was an integer > 2. The series of lines that he observed was called the Balmer series. Subsequently, series of lines were found in the ultraviolet (the Lyman series) and in the infrared (the Paschen, Brackett and Pfund series). These extra lines allowed the generalization of equation (2.1): (2.2) 1 1 2 2 2 ± ² ³ ´ µ R nn H where n 2 > n 1 . R H is known as the Rydberg constant and in this equation has units of reciprocal
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UOIT Chem. 1800, W10: Exp. 2-7.9 53 wavelength (commonly expressed in cm -1 ). The Rydberg constant was a purely empirical value and its literature value is 1.0973 x10 5 cm -1 . It had no theoretical basis until Bohr’s model of the hydrogen atom. In Bohr’s model of the atom the electron is restricted to specific “orbits” circling the nucleus. When it is in the orbit closest to the nucleus the electron has its lowest energy. The lowest energy state (electron in the first orbit) is called the “ground state”. The electron in the lowest orbit can be raised (excited) to a higher state by the addition of energy. However, the energy must exactly match the energy difference between the initial and final states; the electron can only absorb energy in discrete (“quantized”) amounts. When the electron returns to the ground state (or any lower energy state) it will emit energy exactly equal to the energy difference between the upper and lower states. This is why the spectrum of the hydrogen gas appears as discrete lines. Bohr postulated that electrons existed with discrete (or quantized) amounts of angular momentum as they circled the nucleus. Based on this assumption Bohr derived a theoretical expression for the energy levels of the electron in the hydrogen atom: (2.3) E Zem hn n e = 24 0 2 2 2 8 1 ε where: m e = mass of an electron e = charge on an electron h = Planck’s constant ± 0 = permittivity of free space Z = atomic number n = integer The first four of these are constants and Z = 1 for hydrogen. The (1 / n 2 ) factor means the electron is limited to discrete energy levels. If the constants are evaluated, then for hydrogen the
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UOIT Chem. 1800, W10: Exp. 2-7.9
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This note was uploaded on 11/11/2010 for the course CHEMISTRY CHEM1800 taught by Professor Krista during the Winter '10 term at UOIT.

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Lab_2_manual - UOIT Chem. 1800, W10: Exp. 2-7.9 2. ATOMIC...

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