Emission spectra hydrogen

Whereas the emission spectrum of the hydrogen atom shows only one series, the Balmer series (see Figure 1.1), in the visible region the alkali metals show at least three. The spectra can be excited in a discharge lamp containing a sample of the appropriate metal. One series was called the principal series because it could also be observed in absorption through a column of the vapour. The other two were called sharp and diffuse because of their general appearance. A part of a fourth series, called the fundamental series, can sometimes be observed. 

Continuous sources The sources of choice for measurements in the ultraviolet spectral region are hydrogen or deuterium lamps [1]. When the gas pressure is 30 to 60 X10 Pa they yield a continuous emission spectrum. The maxima of their radiation emission occur at different wavelengths (Hi A = 280 nm Di 2 = 220 nm). This means that the deuterium lamp is superior for measurements in the lower UV region (Fig. 15). 

The Humphreys series is set of spectral lines in the emission spectrum of atomic hydrogen that ends in the fifth excited state. 

Schematic representation of an apparatus that measures the emission spectrum of a gaseous element. Emission lines appear bright against a dark background. The spectmm shown is the emission spectrum for hydrogen atoms.

The atomic spectra of most elements are complex and show little regularity. However, the emission spectrum of the hydrogen atom is sufficiently simple to be described by a single formula ... 

Bohr s idea of restricted energy levels was revolutionary, because scientists at that time thought that the electron in a hydrogen atom could have any energy, not just the ones described by Equation. hi contrast, Bohr interpreted the hydrogen emission spectrum to mean that electrons bound to atoms can have only certain specific energy values. 

The Bohr atom went a long way toward explaining the nature of atoms, but there were problems. Although scientists could calculate the emission spectrum of hydrogen using the Bohr model, the model could not account for the spectra of heavier atoms. The biggest problem with the Bohr atom, however, lay in its lack of a... 

By the early twentieth century, scientists had analyzed the spectra of most elements. They knew that each element produced a characteristic emission spectrum. Because hydrogen is the simplest atom, much of the research to understand the nature of atomic spectra centered on it. 

The discovery of two other series of emission lines of hydrogen came later. They are named for their discoverers the Lyman series in the ultraviolet range and Paschen series in the infrared region. Although formulas were devised to calculate the spectral lines, the physics behind the math was not understood until Niels Bohr proposed his quantized atom. Suddenly, the emission spectrum of hydrogen made sense. Each line represented the energy released when an excited electron went from a higher quantum state to a lower one. 

Draw a picture of the electron jump corresponding to the first line in the visible emission spectrum of hydrogen according to the Bohr theory. 

Eventually, other series of lines were found in other regions of the electromagnetic spectrum. The Lyman series was observed in the ultraviolet region, whereas the Paschen, Brackett, and Pfund series were observed in the infrared region of the spectrum. All of these lines were observed as they were emitted from excited atoms, so together they constitute the emission spectrum or line spectrum of hydrogen atoms. 

The first application of quantum theory to a problem in chemistry was to account for the emission spectrum of hydrogen and at the same time explain the stability of the nuclear atom, which seemed to require accelerated electrons in orbital motion. This planetary model is rendered unstable by continuous radiation of energy. The Bohr postulate that electronic angular momentum should be quantized in order to stabilize unique orbits solved both problems in principle. The Bohr condition requires that... 

Figure 2.1 Electronic orbitals and the resulting emission spectrum in the hydrogen atom, (a) Bohr orbitals of the hydrogen atom and the resulting spectral series, (b) emission spectrum of atomic hydrogen. The spectrum in (b) is calibrated in terms of wavenumber (P), which is reciprocal wavelength. The Balmer series, which consists of those transitions terminating on the second orbital, give rise to emission lines in the visible region of the spectrum. ( 1990 John Wiley Sons, Inc. Reprinted from Brady, 1990, by permission of the publisher.)...

Figure 12.6 The emission spectrum of hydrogen in the UY, visible and near infrared, showing the families of lines labeled Lyman (n = 1), Balmer (n — 2), and Paschen (n = 3).

Table 12.1. The wavelengths of the major spectral lines (nm) in the emission spectrum of sodium, arranged in series which are analogous to those in hydrogen.

Thinking Critically How can the single electron in a hydrogen atom produce all of the lines found in its emission spectrum ... 

With the lights dimmed, examine the hydrogen emission spectrum with the... 

How do you think the lines you observed in the hydrogen emission spectrum relate to the energy of the electrons in a hydrogen atom Include a diagram in your answer. 

Compare the hydrogen emission spectrum to the other spectra you observed. Why do you think emission spectra are different for different elements ... 

Why do you think that scientists initially focused their attention on the hydrogen emission spectrum rather than the emission spectrum of another element ... 

How the Bohr model explains the coloured lines in hydrogen s emission spectrum. When an excited electron falls from a higher energy level to a lower energy level (shown by the downward-pointing arrows), it emits a photon with a specific wavelength that corresponds to one of the coloured lines in the spectrum. 

Bohr s realization that the atom s energy is quantized—that electrons are restricted to specific energy levels (orbits)— was an astounding achievement. As you have seen, this model successfully predicted the coloured lines in the visible-light portion of hydrogen s emission spectrum. It also successfully predicted other lines, shown in Figure 3.11, that earlier chemists had discovered in the ultraviolet and infrared portions of hydrogen s emission spectrum. 

The emission spectrum for mercury shows that it has more spectral lines than the emission spectrum for hydrogen. 

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