Hydrogen line-emission spectrum

Explain the reason for the hydrogen line-emission spectrum. 

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. 

A. Schuster 8 found that the spectrum of ammonia in the discharge tube shows a broad, greenish-yellow band between 5688 and 5627. G. Magnanini observed the spectrum of the flame of ammonia burning in oxygen exhibits a large number of hydrogen lines. This spectrum was also observed by J. M. Eder, who measured 240 lines between A=5000 and 2262 for the extreme ultra-violet. The emission spectrum has seven characteristic bands—one between the red and ultra-violet. 

Bohr developed an equation to calculate all of the possible energies of the electron in a hydrogen atom. His values agreed with those calculated from the wavelengths observed in hydrogen s line-emission spectrum. In fact, his values matched with the experimental values so well that his atomic model that is described earlier was quickly accepted. 

Figure 5.8 shows an illustration of the characteristic purple-pink glow produced by excited hydrogen atoms and the visible portion of hydrogen s emission spectrum responsible for producing the glow. Note how the line nature of hydrogens atomic emission spectrum differs from that of a continuous spectrum. 

The dual wave-particle model of light accounted for several previously unexplainable phenomena, but scientists still did not understand the relationships among atomic structure, electrons, and atomic emission spectra. Recall that hydrogens atomic emission spectrum is discontinuous that is, it is made up of only certain frequencies of light. Why are the atomic emission spectra of elements discontinuous rather than continuous Niels Bohr, a Danish physicist working in Rutherford s laboratory in 1913, proposed a quantum model for the hydrogen atom that seemed to answer this question. Bohr s model also correctly predicted the frequencies of the lines in hydrogens atomic emission spectrum. 

What electron transitions account for the Balmer series Hydrogen s emission spectrum comprises three series of lines. Some wavelengths are ultraviolet (Lyman series) and infrared (Paschen series). Visible wavelengths comprise the Balmer series. The Bohr atomic model attributes these spectral lines to transitions from higher-energy states with electron orbits in which n = n, to lower-energy states with smaller electron orbits in which n = nf. 

Hydrogen One Hne in hydrogen s emission spectrum has a wavelength of 486 nm. Examine Figure 5.22 to determine the line s color. What is the line s frequency ... 

The simplest type of electronic spectra result from transitions within the simplest atom, hydrogen. The emission spectrum of the hydrogen atom at ultraviolet wavelengths consists of a series of emission peaks (or in a photographic emulsion, dark lines) called the Lyman series. 

As shown with hydrogen, the emission spectrum of an element consists of several series of lines arising from transitions from some higher energy state to a common lower state. In hydrogen these series are quite well separated, with only minimal overlap. Reference to Figure 2-2 will make this clear. 

Figure 1.12 The line emission spectrum of hydrogen atoms in the visible region.

Each canponent color is focused at a definite position, according to its wavelength, and forms a colored image of the slit on the photographic plate. The colored irtuxges are called spectral lines, (b) The line emission spectrum of hydrogen atoms. 

One of the lines in the Balmer series of the hydrogen atom emission spectrum is at 397 nm. It results from a transition from an upper energy level ton = 2. What is the principal quantum number of the upper level ... 

Quantum theory was developed to explain observations such as the photoelectric effect and the line-emission spectrum of hydrogen. 

Figure 23.2 The Visible Portion of the Hydrogen Atom Emission Spectrum (Simulated). Each wavelength represented produces an image of the slit of the spectrograph. If only discrete wavelengths are present, as in this case, the spectrum is called a line spectrum.

Continnous and line emission spectra. From the top down The continuous visible spectrum the line emission spectra for sodium (Na). hydrogen (H). and mercury (Hg). 

Historically, the visible emission lines shown in Figure 15-3 were the first atomic hydrogen lines discovered. They were found in the spectrum of the sun by W. H. Wollaston in 1802. In 1862, A. J. Angstrom announced that there must be hydrogen in the solar atmosphere. These lines were detected first because of the lesser experimental difficulties in the visible spectral region. They are called the "Balmer series because J. J. Balmer was able to formulate a simple mathematical relation among the frequencies (in It S). The ultraviolet series shown in Figure 15-3 was... 

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 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. 

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.)...

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