Spectroscopy - CH 221 - Experiment MEASUREMENT AND...

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CH 221 - Experiment MEASUREMENT AND ASSIGNMENT OF ATOMIC SPECTRA Purpose: The purpose of this laboratory is to become familiar with the emission spectrum of Hydrogen, the basis for modern quantum mechanics, and atomic orbital theory for the rest of the elements on the periodic chart. We will test the idea that the positions of the lines in the emission spectrum of Hydrogen can be calculated from a derived equation, by both predicting from calculation and then measure directly the emission lines of Hydrogen. Students will also compare the spectra of other elements, and use the characteristic emissions of salts to identify an unknown. While doing this exercise, we will take into account the limits of precision of the measurements. Introduction The emission from a collection of excited hydrogen atoms is characteristic of Hydrogen, and the emission spectrum for every other element is characteristic of those elements. These emission lines, characterized carefully on earth, are used to identify elements both in the laboratory and also in distant stars. The light that makes up these spectra originates in the atoms themselves, and the explanation of their origins was a hotly debated issue in the early twentieth century. The first successful theory for the structure of the hydrogen atom was formulated by Niels Bohr in 1913. Basically, Bohr assumed that a hydrogen atom consisted of a central proton around which an electron moved in a circular orbit, much as the earth moves around the sun. By imposing quantum theory on his model, Bohr then found that the electron could only be in orbits having certain definite radii. In addition, each allowed orbit had associated with it a definite energy. Thus, in the Bohr model, the electron orbital radii and the energies of the hydrogen atom are quantized or only allowed at certain energies. In subsequent, even better, models of the hydrogen atom, it was found that the orbits of fixed radii for the electron had to be replaced by fuzzy orbitals having fixed average distances from the proton. However, the associated fixed energies calculated by Bohr seem to be correct to this day. Bohr's equation for the allowed energies of the hydrogen atom is B E = - 2 n (equation 1) where n, the principle quantum number, can have the integer values 1, 2, 3. .., and B is a constant having the value 2.179 x 10 -18 joules. The normal, or ground state, energy of the hydrogen atom corresponds to the energy calculated when n = 1, and is the state in which the electron is, on the average, closest to the proton. The energy of the atom can be increased to any energy allowed by the Bohr equation (1) by pumping energy into the atom in some fashion (for example, by exposing it to electric current). Then the atom is said to be in an excited state. For example, if the energy of the atom corresponds to n = 2, the atom is said to be in the first excited state. As n increases, the energy of the atom increases (and vice versa) and the electron moves to new orbitals that, on the
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Spectroscopy - CH 221 - Experiment MEASUREMENT AND...

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