DOCSLIB.ORG

EXPERIMENT Hydrogen Emission Spectra

Total Page：16

File Type：pdf, Size：1020Kb

Download full-text PDF Read full-text
EXPERIMENT Hydrogen Emission Spectra EXPERIMENT Hydrogen Emission Spectra PURPOSE In this experiment you will use a simple spectroscope to observe the line spectrum of hydrogen, identify the wavelength of each transition, and determine the corresponding energy levels. You will also show that these wavelengths fit a pattern of energy states described by a simple formula containing a single constant and integer quantum numbers. BACKGROUND When an electric current is passed through a sample of gas in a sealed tube the energy excites the electrons of the atoms, causing them to jump to higher energy levels, n (high). Then, as the electrons fall back from the higher to lower energy levels, n (low), that energy is emitted as light. When this light is passed through a prism we do not observe a continuous spectrum, but a line spectrum. In other words, only certain wavelengths of light are observed. This observation strongly suggests that the energy absorbed/emitted by an electron is quantized, or restricted to certain discreet values. Each line in the spectrum corresponds to a particular electronic transition between discreet energy levels. For the emission Figures 1- A simple spectroscope on the right and on the left “Electron transitions for the Hydrogen atom” is shown along with corresponding wavelengths for the visible region. Note that transition between any pair of states such that ni > nf produces a photon; however, only those transition with nf = 2 and ni= 3, 4, 5, or 6, happen to produce photons in visible range wavelengths. Energy (ΔE) for these transitions can be calculated: (ΔE = − [ Ef − Ei]). Page 1 In this lab, by using a simple spectroscope, at first you will construct a calibration graph using four emissions in atomic spectrum of mercury. In the second part you will measure the position of four lines in the atomic spectrum of hydrogen. The only lines you will be able to observe are those of the Balmer series, visible region of spectrum between 400 and 700 nm. The lines of the Balmer series are the lines for which n (lower), is equal to 2 (i.e. n (low) = 2). Other transitions show up in other regions of the electromagnetic spectrum. For example, all transitions with n (low) = 1 fall in the UV region of the spectrum, while all transitions with n (low) = 3 fall in the infrared. With different detection equipment we could observe those transitions as well. Once you have identified the transition of each line according to the observed color using figure 1, you will be able determine the corresponding wavelength, λ using calibration graph. Then you could calculate the Rydberg’s constant using the Rydberg equation (below). In other words, given that n (low or final) = 2 (true for all lines in the visible region), you can determine the n (high or initial), and subsequently the energy level from which the "excited" electron came down from. It is important to note that Rydberg constant is in SI unit “reciprocal meter (m-1)” and represent a 1 limiting value of the highest wavenumber ( ) of a photon that can be emitted from the hydrogen 휆 atom or, alternatively, the wavenumber of the lowest –energy photon capable of ionizing the hydrogen atom from its ground state. Rydberg’s equation & wavelength Rydberg’s equation & energy 1 1 1 ℎ푐 1 1 = 푅 ( − ) ∆퐸 = ℎ푣 = = −푅퐻ℎ푐 ( 2 − 2) 퐻 2 2 휆 푛푓 푛푖 휆 푛푓 푛 ℎ푐 1 1 1 7 −1 1 1 −18 = 1.096776 × 10 푚 ( − ) ∆퐸 = ℎ푣 = = −2.179 × 10 퐽 ( 2 − 2) 2 2 휆 푛푓 푛푖 휆 푛푓 푛 Note: Rydberg's Constant is 1.096776 x 10 7 m-1 which is a distance. Some internet sites say that Rydberg's constant is equal to 2.17869 x 10 -18 Joules but this is not correct. They are using the constant that is derived and is equal to (RH) times (h) times (c). Page 2 Name ____________________________ LAB Atomic Emission Spectrum REPORT Section _________ Date _____________ Instructor _________________________ Part I: Calibration of the spectroscope (a) Measure the position of the four observed lines corresponding to the known wavelengths by examining the mercury spectrum and enter the result in the table below. Observed line Positions Color Known Wavelengths (nm) recorded to the nearest 0.1 cm Violet 404.7 Blue 435.8 Green 546.1 Yellow 579.0 (b) Describe any other lines that do not correspond to the known wavelengths. (c) Make a calibration graph by plotting the known wavelengths versus the recorded line positions on the graph below Page 3 Name ____________________________ LAB Atomic Emission Spectrum REPORT Section _________ Date _____________ Instructor _________________________ Part II: Observation and measurement of the hydrogen spectrum: (a) Exam the hydrogen spectrum and enter the observed line positions in the table below and then determine the corresponding wavelengths using the mercury calibration graph. Observed line Wavelength form the Known Percentage Color position (cm)* calibration graph wavelength (nm) error (nm) Red 656.43 Greenish Blue 486.26 Bluish Violet 434.16 Purple 410.28 *recorded to the nearest 0.1 cm (b) Calculate the corresponding energy per phone (kJ/photon) for each calibrated wavelength and then calculate the corresponding energy (kJ/mol) per mole by using Avogadro’s number. (c) Using Figure1 determines the values of quantum numbers n1 and n2 for the initial and final states of the transitions that give rise to each line. Wavelength from the Photon energy Value of ni Value of nf calibration graph (nm) (kJ/mol) (initial state) (final state) (d) Using the equations provided, calculate the value of the Rydberg constant for each wavelength value in part (c). Enter these values in the table in part e) below. Give one example to show your method of calculation. (e) Look up the reported value of RH_____________________. Compute your percentage error for determination. Wavelength from the Experimental value of RH Percentage error calibration graph (nm) Page 4 Name ____________________________ LAB Atomic Emission Spectrum REPORT Section _________ Date _____________ Questions Instructor _________________________ 1. Using the Bohr equation, show that the energy of the transition from the lowest (ground) state (n1 = 1) to the highest energy state (n2 = ∞) is given by ∆E (kJ/mol ) = -RH (hc), the derivative of Rydberg constant in terms of Planck’s constant and speed of light. Therefore, the magnitude of this constant is equal to the ionization energy of one mole of hydrogen atoms, in units of kJ/mol. 2. How many emission are possible when hydrogen atom spectra in exited state return to lower energy state and ultimately ground state for emission spectral lines when a hydrogen atom’s electron in the n = 4 quantum level of a Bohr atom drops to the ground state (i.e. n = 1)? Use a diagram to illustrate each of the possible transitions responsible for the spectral lines. 3. Use the Rydberg equation to calculate the wavelength in nm of the light emitted when the electron in a hydrogen atom undergoes a transition from n = 6 to n = 2. Consult figure1 to determine the color of light emitted. 4. Calculate the energy in units of Joules for each photon of light in question above. Page 5 .
Load more

Recommended publications
  • Quantum Theory of the Hydrogen Atom
    Quantum Theory of the Hydrogen Atom Chemistry 35 Fall 2000 Balmer and the Hydrogen Spectrum n 1885: Johann Balmer, a Swiss schoolteacher, empirically deduced a formula which predicted the wavelengths of emission for Hydrogen: l (in Å) = 3645.6 x n2 for n = 3, 4, 5, 6 n2 -4 •Predicts the wavelengths of the 4 visible emission lines from Hydrogen (which are called the Balmer Series) •Implies that there is some underlying order in the atom that results in this deceptively simple equation. 2 1 The Bohr Atom n 1913: Niels Bohr uses quantum theory to explain the origin of the line spectrum of hydrogen 1. The electron in a hydrogen atom can exist only in discrete orbits 2. The orbits are circular paths about the nucleus at varying radii 3. Each orbit corresponds to a particular energy 4. Orbit energies increase with increasing radii 5. The lowest energy orbit is called the ground state 6. After absorbing energy, the e- jumps to a higher energy orbit (an excited state) 7. When the e- drops down to a lower energy orbit, the energy lost can be given off as a quantum of light 8. The energy of the photon emitted is equal to the difference in energies of the two orbits involved 3 Mohr Bohr n Mathematically, Bohr equated the two forces acting on the orbiting electron: coulombic attraction = centrifugal accelleration 2 2 2 -(Z/4peo)(e /r ) = m(v /r) n Rearranging and making the wild assumption: mvr = n(h/2p) n e- angular momentum can only have certain quantified values in whole multiples of h/2p 4 2 Hydrogen Energy Levels n Based on this model, Bohr arrived at a simple equation to calculate the electron energy levels in hydrogen: 2 En = -RH(1/n ) for n = 1, 2, 3, 4, .
    [Show full text]
  • Rydberg Constant and Emission Spectra of Gases
    Page 1 of 10 Rydberg constant and emission spectra of gases ONE WEIGHT RECOMMENDED READINGS 1. R. Harris. Modern Physics, 2nd Ed. (2008). Sections 4.6, 7.3, 8.9. 2. Atomic Spectra line database https://physics.nist.gov/PhysRefData/ASD/lines_form.html OBJECTIVE - Calibrating a prism spectrometer to convert the scale readings in wavelengths of the emission spectral lines. - Identifying an "unknown" gas by measuring its spectral lines wavelengths. - Calculating the Rydberg constant RH. - Finding a separation of spectral lines in the yellow doublet of the sodium lamp spectrum. INSTRUCTOR’S EXPECTATIONS In the lab report it is expected to find the following parts: - Brief overview of the Bohr’s theory of hydrogen atom and main restrictions on its application. - Description of the setup including its main parts and their functions. - Description of the experiment procedure. - Table with readings of the vernier scale of the spectrometer and corresponding wavelengths of spectral lines of hydrogen and helium. - Calibration line for the function “wavelength vs reading” with explanation of the fitting procedure and values of the parameters of the fit with their uncertainties. - Calculated Rydberg constant with its uncertainty. - Description of the procedure of identification of the unknown gas and statement about the gas. - Calculating resolution of the spectrometer with the yellow doublet of sodium spectrum. INTRODUCTION In this experiment, linear emission spectra of discharge tubes are studied. The discharge tube is an evacuated glass tube filled with a gas or a vapor. There are two conductors – anode and cathode - soldered in the ends of the tube and connected to a high-voltage power source outside the tube.
    [Show full text]
  • Improving the Accuracy of the Numerical Values of the Estimates Some Fundamental Physical Constants
    Improving the accuracy of the numerical values of the estimates some fundamental physical constants. Valery Timkov, Serg Timkov, Vladimir Zhukov, Konstantin Afanasiev To cite this version: Valery Timkov, Serg Timkov, Vladimir Zhukov, Konstantin Afanasiev. Improving the accuracy of the numerical values of the estimates some fundamental physical constants.. Digital Technologies, Odessa National Academy of Telecommunications, 2019, 25, pp.23 - 39. hal-02117148 HAL Id: hal-02117148 https://hal.archives-ouvertes.fr/hal-02117148 Submitted on 2 May 2019 HAL is a multi-disciplinary open access L’archive ouverte pluridisciplinaire HAL, est archive for the deposit and dissemination of sci- destinée au dépôt et à la diffusion de documents entific research documents, whether they are pub- scientifiques de niveau recherche, publiés ou non, lished or not. The documents may come from émanant des établissements d’enseignement et de teaching and research institutions in France or recherche français ou étrangers, des laboratoires abroad, or from public or private research centers. publics ou privés. Improving the accuracy of the numerical values of the estimates some fundamental physical constants. Valery F. Timkov1*, Serg V. Timkov2, Vladimir A. Zhukov2, Konstantin E. Afanasiev2 1Institute of Telecommunications and Global Geoinformation Space of the National Academy of Sciences of Ukraine, Senior Researcher, Ukraine. 2Research and Production Enterprise «TZHK», Researcher, Ukraine. *Email: [email protected] The list of designations in the text: l
    [Show full text]
  • Chapter 09584
    Author's personal copy Excitons in Magnetic Fields Kankan Cong, G Timothy Noe II, and Junichiro Kono, Rice University, Houston, TX, United States r 2018 Elsevier Ltd. All rights reserved. Introduction When a photon of energy greater than the band gap is absorbed by a semiconductor, a negatively charged electron is excited from the valence band into the conduction band, leaving behind a positively charged hole. The electron can be attracted to the hole via the Coulomb interaction, lowering the energy of the electron-hole (e-h) pair by a characteristic binding energy, Eb. The bound e-h pair is referred to as an exciton, and it is analogous to the hydrogen atom, but with a larger Bohr radius and a smaller binding energy, ranging from 1 to 100 meV, due to the small reduced mass of the exciton and screening of the Coulomb interaction by the dielectric environment. Like the hydrogen atom, there exists a series of excitonic bound states, which modify the near-band-edge optical response of semiconductors, especially when the binding energy is greater than the thermal energy and any relevant scattering rates. When the e-h pair has an energy greater than the binding energy, the electron and hole are no longer bound to one another (ionized), although they are still correlated. The nature of the optical transitions for both excitons and unbound e-h pairs depends on the dimensionality of the e-h system. Furthermore, an exciton is a composite boson having integer spin that obeys Bose-Einstein statistics rather than fermions that obey Fermi-Dirac statistics as in the case of either the electrons or holes by themselves.
    [Show full text]
  • 1. Physical Constants 1 1
    1. Physical constants 1 1. PHYSICAL CONSTANTS Table 1.1. Reviewed 1998 by B.N. Taylor (NIST). Based mainly on the “1986 Adjustment of the Fundamental Physical Constants” by E.R. Cohen and B.N. Taylor, Rev. Mod. Phys. 59, 1121 (1987). The last group of constants (beginning with the Fermi coupling constant) comes from the Particle Data Group. The figures in parentheses after the values give the 1-standard- deviation uncertainties in the last digits; the corresponding uncertainties in parts per million (ppm) are given in the last column. This set of constants (aside from the last group) is recommended for international use by CODATA (the Committee on Data for Science and Technology). Since the 1986 adjustment, new experiments have yielded improved values for a number of constants, including the Rydberg constant R∞, the Planck constant h, the fine- structure constant α, and the molar gas constant R,and hence also for constants directly derived from these, such as the Boltzmann constant k and Stefan-Boltzmann constant σ. The new results and their impact on the 1986 recommended values are discussed extensively in “Recommended Values of the Fundamental Physical Constants: A Status Report,” B.N. Taylor and E.R. Cohen, J. Res. Natl. Inst. Stand. Technol. 95, 497 (1990); see also E.R. Cohen and B.N. Taylor, “The Fundamental Physical Constants,” Phys. Today, August 1997 Part 2, BG7. In general, the new results give uncertainties for the affected constants that are 5 to 7 times smaller than the 1986 uncertainties, but the changes in the values themselves are smaller than twice the 1986 uncertainties.
    [Show full text]
  • Compton Wavelength, Bohr Radius, Balmer's Formula and G-Factors
    Compton wavelength, Bohr radius, Balmer's formula and g-factors Raji Heyrovska J. Heyrovský Institute of Physical Chemistry, Academy of Sciences of the Czech Republic, Dolejškova 3, 182 23 Prague 8, Czech Republic. [email protected] Abstract. The Balmer formula for the spectrum of atomic hydrogen is shown to be analogous to that in Compton effect and is written in terms of the difference between the absorbed and emitted wavelengths. The g-factors come into play when the atom is subjected to disturbances (like changes in the magnetic and electric fields), and the electron and proton get displaced from their fixed positions giving rise to Zeeman effect, Stark effect, etc. The Bohr radius (aB) of the ground state of a hydrogen atom, the ionization energy (EH) and the Compton wavelengths, λC,e (= h/mec = 2πre) and λC,p (= h/mpc = 2πrp) of the electron and proton respectively, (see [1] for an introduction and literature), are related by the following equations, 2 EH = (1/2)(hc/λH) = (1/2)(e /κ)/aB (1) 2 2 (λC,e + λC,p) = α2πaB = α λH = α /2RH (2) (λC,e + λC,p) = (λout - λin)C,e + (λout - λin)C,p (3) α = vω/c = (re + rp)/aB = 2πaB/λH (4) aB = (αλH/2π) = c(re + rp)/vω = c(τe + τp) = cτB (5) where λH is the wavelength of the ionizing radiation, κ = 4πεo, εo is 2 the electrical permittivity of vacuum, h (= 2πħ = e /2εoαc) is the 2 Planck constant, ħ (= e /κvω) is the angular momentum of spin, α (= vω/c) is the fine structure constant, vω is the velocity of spin [2], re ( = ħ/mec) and rp ( = ħ/mpc) are the radii of the electron and
    [Show full text]
  • The Fine Structure Constant, the Rydberg Constant and the Planck
    Perspective iMedPub Journals Polymer Sciences 2017 www.imedpub.com ISSN 2471-9935 Vol.3 No.2:13 DOI: 10.4172/2471-9935.100028 The Fine Structure Constant, the Rydberg YinYue Sha* Constant and the Planck Constant Dongling Engineering Center, Ningbo Institute of Technology, Zhejiang University, PR China Abstract *Corresponding author: YinYue Sha So let's say that the electron is orbiting the proton in the innermost orbital, and then, according to the equilibrium relationship of forces, there's the following [email protected] formula: F=K × Qp × Qe/(Rb × Rb)=Me × Ve × Ve/Rb (1) Dongling Engineering Center, Ningbo The K is the electromagnetic constant; Qp is the charge of proton; Qe is the charge Institute of Technology, Zhejiang University, of electron; Rb is the Bohr atom radius; Me is the mass of the electron; Ve is the PR China. speed of electrons. Tel: +86 574 8822 9048 Keywords: Protons; Electronic; Mass; Charge; Bohr atom radius; Fine structure constant; The Rydberg constant; Planck constant Citation: Sha YY (2017) The Fine Structure Received: July 12, 2017; Accepted: September 20, 2017; Published: September 30, Constant, the Rydberg Constant and the 2017 Planck Constant. Polym Sci Vol.3 No.2:13 One, the fine structure constant and the Bohr atomic radius. Three, The kinetic energy of the ground state electron and the Planck constant From the formula (1), we can obtain the formula for calculating the basic radius of the hydrogen atom: The kinetic energy of the ground state electron: Rb=K × Qp × Qe/(Me × Ve2) (2) Eb=1/2 × Me × Ve2=2.1798719696853
    [Show full text]
  • Giant Rydberg Excitons in Cu2o Probed by Photoluminescence
    Giant Rydberg excitons in Cu2O probed by photoluminescence excitation spectroscopy Marijn A. M. Versteegh,1, ∗ Stephan Steinhauer,1 Josip Bajo,1 Thomas Lettner,1 Ariadna Soro,1, † Alena Romanova,1, 2 Samuel Gyger,1 Lucas Schweickert,1 Andr´eMysyrowicz,3 and Val Zwiller1 1Department of Applied Physics, KTH Royal Institute of Technology, Stockholm, Sweden 2Johannes Gutenberg University Mainz, Germany 3Laboratoire d’Optique Appliqu´ee, ENSTA, CNRS, Ecole´ Polytechnique, Palaiseau, France (Dated: May 18, 2021) Rydberg excitons are, with their ultrastrong mutual interactions, giant optical nonlinearities, and very high sensitivity to external fields, promising for applications in quantum sensing and nonlin- ear optics at the single-photon level. To design quantum applications it is necessary to know how Rydberg excitons and other excited states relax to lower-lying exciton states. Here, we present pho- toluminescence excitation spectroscopy as a method to probe transition probabilities from various excitonic states in cuprous oxide, and we show giant Rydberg excitons at T = 38 mK with principal quantum numbers up to n = 30, corresponding to a calculated diameter of 3 µm. Rydberg excitons are hydrogen-atom-like bound Rydberg excitons in Cu2O were identified and studied electron-hole pairs in a semiconductor with principal in electric [17, 18] and magnetic fields [19–23]. Studies quantum number n ≥ 2. Rydberg excitons share com- of the absorption spectrum revealed quantum coherences mon features with Rydberg atoms, which are widely stud- [24] and brought an understanding of the interactions of ied as building blocks for emerging quantum technologies, Rydberg excitons with phonons and photons [25–27], di- such as quantum computing and simulation, quantum lute electron-hole plasma [13, 14], and charged impurities sensing, and quantum photonics [1].
    [Show full text]
  • Atomic Structure
    1 Atomic Structure 1 2 Atomic structure 1-1 Source 1-1 The experimental arrangement for Ruther- ford's measurement of the scattering of α particles by very thin gold foils. The source of the α particles was radioactive radium, a rays encased in a lead block that protects the surroundings from radiation and confines the α particles to a beam. The gold foil used was about 6 × 10-5 cm thick. Most of the α particles passed through the gold leaf with little or no deflection, a. A few were deflected at wide angles, b, and occasionally a particle rebounded from the foil, c, and was Gold detected by a screen or counter placed on leaf Screen the same side of the foil as the source. The concept that molecules consist of atoms bonded together in definite patterns was well established by 1860. shortly thereafter, the recognition that the bonding properties of the elements are periodic led to widespread speculation concerning the internal structure of atoms themselves. The first real break- through in the formulation of atomic structural models followed the discovery, about 1900, that atoms contain electrically charged particles—negative electrons and positive protons. From charge-to-mass ratio measurements J. J. Thomson realized that most of the mass of the atom could be accounted for by the positive (proton) portion. He proposed a jellylike atom with the small, negative electrons imbedded in the relatively large proton mass. But in 1906-1909, a series of experiments directed by Ernest Rutherford, a New Zealand physicist working in Manchester, England, provided an entirely different picture of the atom.
    [Show full text]
  • Hydrogen-Deuterium Mass Ratio
    Hydrogen-Deuterium Mass Ratio 1 Background 1.1 Balmer and Rydberg Formulas By the middle of the 19th century it was well established that atoms emitted light at discrete wavelengths. This is in contrast to a heated solid which emits light over a continuous range of wavelengths. The difference between continuous and discrete spectra can be readily observed by viewing both incandescent and fluorescent lamps through a diffraction grating; hand held gratings are available for you to do this in the laboratory. Around 1860 Kirchoff and Bunsen discovered that each element has its own characteristic spectrum. The next several decades saw the accumulation of a wealth of spectroscopic data from many elements. What was lacking though, was a mathematical relation between the various spectral lines from a given element. The visible spectrum of hydrogen, being rela- tively simple compared to the spectra of other elements, was a particular focus of attempts to find an empirical relation between the wavelengths of its spectral lines. In 1885 Balmer discovered that the wavelengths λn of the then nine known lines in the hydrogen spectrum were described to better than one part in a thousand by the formula 2 2 λn = 3646 × n /(n − 4) A˚ (angstrom), (1) where n = 3, 4, 5,... for the various lines in this series, now known as the Balmer series. (1 angstrom = 10−10 m.) The more commonly used unit of length on this scale is the nm (10−9 m), but since the monochromator readout is in angstroms, we will use the latter unit throughout this write-up.
    [Show full text]
  • The Rydberg Constant
    The Rydberg Constant The Rydberg Constant The light you see when you plug in a hydrogen gas discharge tube is a shade of lavender, with some pinkish tint at a higher current. If you observe the light through a spectroscope, you can identify four distinct lines of color in the visible light range. The history of the study of these lines th dates back to the late 19 century, where we meet a high school mathematics teacher from Basel, Switzerland, named Johann Balmer. Balmer created an equation describing the wavelengths of the visible hydrogen emission lines. However, he did not support his equation with a physical explanation. In a paper written in 1885, Balmer proposed that his equation could be used to predict the entire spectrum of hydrogen, including the ultra-violet and the infrared spectral lines. The Balmer equation is shown below. -8 where m and n were integers, and h = 3654.6 10 cm. When one solves the equation using n =2 and m = 3, 4, 5, or 6, the calculated wavelengths are very close to the four emission lines in the visible light range for a hydrogen gas discharge tube. Balmer apparently derived his equation by trail and error. Sadly, he would not live to see Niels Bohr and Johannes Rydberg prove the validity of his equation. Johannes Rydberg was a mathematics teacher like Balmer (he also taught a bit of physics). In 1890, Rydberg’s research of spectroscopy (inspired, it is said, by the work of Dmitri Mendeleev) led to his discovery that Balmer’s equation was a specific case of a more general principle.
    [Show full text]
  • 28-1 Line Spectra and the Hydrogen Atom Figure 28.1 Gives Some Examples of the Line Spectra Emitted by Atoms of Gas
    28-1 Line Spectra and the Hydrogen Atom Figure 28.1 gives some examples of the line spectra emitted by atoms of gas. The atoms are typically excited by applying a high voltage across a glass tube that contains a particular gas. By observing the light through a diffraction grating, the light is separated into a set of wavelengths that characterizes the element. The spectrum of light emitted by excited hydrogen atoms is shown Figure 28.1: Line spectra from hydrogen (top) and helium in Figure 28.1(a). The decoding of the (bottom). A line spectrum is like the fingerprint of an element. hydrogen spectrum represents one of Astronomers, for instance, can determine what a star is made of by the great scientific mystery stories. carefully examining the spectrum of light emitted by the star. First on the scene was the Swiss mathematician and schoolteacher, Johann Jakob Balmer (1825 – 1898), who published an equation in 1885 giving the wavelengths in the visible spectrum emitted by hydrogen. The Swedish physicist Johannes Rydberg (1854 – 1919) followed up on Balmer’s work in 1888 with a more general equation that predicted all the wavelengths of light emitted by hydrogen: , (Eq. 28.1: The Rydberg equation for the hydrogen spectrum) 7 –1 where R = 1.097 ! 10 m is the Rydberg constant, and the two n’s are integers, with n2 greater than n1. Neither Balmer nor Rydberg had a physical explanation to justify their equations, however, so the search was on for such a physical explanation. The breakthrough was made by the Danish physicist Niels Bohr (1885 – 1962), who showed that if the angular momentum of the electron in a hydrogen atom was quantized in a particular way (related to Planck’s constant, in fact), that the energy levels for an electron within the hydrogen atom were also quantized, with the energies of the electrons being given by: , (Eq.
    [Show full text]