Hydrogen atom
Hydrogen atom  

Hydrogen atom 

General  
Name, symbol  protium,^{1}H 
Neutrons  0 
Protons  1 
Nuclide data  
Natural abundance  99.985% 
Isotope mass  1.007825 u 
Spin  ^{1}⁄_{2}+ 
Excess energy  7288.969± 0.001 keV 
Binding energy  0.000± 0.0000 keV 
A hydrogen atom is an atom of the chemical element hydrogen. The electrically neutral atom contains a single positively charged proton and a single negatively charged electron bound to the nucleus by the Coulomb force. Atomic hydrogen constitutes about 75% of the elemental (baryonic) mass of the universe.^{[1]}
In everyday life on Earth, isolated hydrogen atoms (usually called "atomic hydrogen" or, more precisely, "monatomic hydrogen") are extremely rare. Instead, hydrogen tends to combine with other atoms in compounds, or with itself to form ordinary (diatomic) hydrogen gas, H_{2}. "Atomic hydrogen" and "hydrogen atom" in ordinary English use have overlapping, yet distinct, meanings. For example, a water molecule contains two hydrogen atoms, but does not contain atomic hydrogen (which would refer to isolated hydrogen atoms).
Attempts to develop a theoretical understanding of the hydrogen atom have been important to the history of quantum mechanics.
Contents
 1 Isotopes
 2 Hydrogen ion
 3 Theoretical analysis
 4 See also
 5 References
 6 Books
 7 External links
Isotopes
The most abundant isotope, hydrogen1, protium, or light hydrogen, contains no neutrons and is just a proton and an electron. Protium is stable and makes up 99.9885% of naturally occurring hydrogen by absolute number (not mass).
Deuterium contains one neutron and one proton. Deuterium is stable and makes up 0.0115% of naturally occurring hydrogen and is used in industrial processes like nuclear reactors and Nuclear Magnetic Resonance.
Tritium contains two neutrons and one proton and is not stable, decaying with a halflife of 12.32 years. Because of the short half life, Tritium does not exist in nature except in trace amounts.
Higher isotopes of hydrogen are only created in artificial accelerators and reactors and have half lives around the order of 10^{−22} seconds.
The formulas below are valid for all three isotopes of hydrogen, but slightly different values of the Rydberg constant (correction formula given below) must be used for each hydrogen isotope.
Hydrogen ion
Hydrogen is not found without its electron in ordinary chemistry (room temperatures and pressures), as ionized hydrogen is highly chemically reactive. When ionized hydrogen is written as "H^{+}" as in the solvation of classical acids such as hydrochloric acid, the hydronium ion, H_{3}O^{+}, is meant, not a literal ionized single hydrogen atom. In that case, the acid transfers the proton to H_{2}O to form H_{3}O^{+}.
Ionized hydrogen without its electron, or free protons, are common in the interstellar medium, and solar wind.
Theoretical analysis
The hydrogen atom has special significance in quantum mechanics and quantum field theory as a simple twobody problem physical system which has yielded many simple analytical solutions in closedform.
Failed classical description
Experiments by Rutherford in 1909 showed the structure of the atom be a dense, positive nucleus with a light, negative charge orbiting around it. This immediately caused problems on how such a system could be stable. Classical electromagnetism had shown that any accelerating charge radiates energy described through the Larmor formula. If the electron is assumed to orbit in a perfect circle and radiates energy adiabatically, the electron would spiral into the nucleus with a fall time of:^{[2]}
Where is the Bohr radius and is the classical electron radius. If this were true, all atoms would instantly collapse, however atoms seem to be stable. Furthermore, the spiral inward would release a smear of electromagnetic frequencies as the orbit got smaller. Instead, atoms were observed to only emit discrete frequencies of light. The resolution would lie in the development of quantum mechanics.
Bohr Model
In 1913, Niels Bohr obtained the energy levels and spectral frequencies of the hydrogen atom after making a number of simplifying assumptions in order to account for the failed Classical model. The assumptions included:
 Electrons can only be in certain, discrete orbitals, thereby having a discrete radius and energy.
 Electrons do not emit radiation while in one of these stationary states.
 An electron can gain or lose energy by jumping from one discrete orbital to another.
The assumption that angular momentum was quantized can be expressed as:
where
and is Planck constant over . Using this, the force relation between the Centripetal force and Coulomb's force, and energy conservation, Bohr derived the energy of each orbital of the hydrogen atom to be:^{[3]}

 ,
where is the electron mass, is the electron charge, is the electric permeability, and is the quantum number (now known as the principal quantum number). Bohr's predict matched experiments measuring the hydrogen spectral series to the first order, giving more confidence to a theory that used quantized values.
There were still problems with Bohr's model, it failed to predict other spectral lines such as fine structure and hyperfine structure, it could only predict energy levels for hydrogen like (single electron) atoms with any accuracy and the predicted values were only correct to , where is the finestructure constant. The Bohr model assumed circular orbits where, as developed by Sommerfeld, elliptial orbits adds other quantum numbers besides and changes energy values. Furthermore it didn't explain many other observed phenomena such as the Zeeman effect, Stark effect and violates the uncertainty principle.
These issues were resolved with the full development of quantum mechanics and the Schrödinger equation in 1925–1926. The solutions to the Schrödinger equation for hydrogen are analytical, giving a simple expression for the hydrogen energy levels and thus the frequencies of the hydrogen spectral lines and fully reproduced the Bohr model and went beyond it. It also yields two other quantum numbers and the shape of the electron's wave function ("orbital") for the various possible quantummechanical states, thus explaining the anisotropic character of atomic bonds.
The Schrödinger equation also applies to more complicated atoms and molecules. When there is more than one electron or nucleus the solution is not analytical and either computer calculations are necessary or simplifying assumptions must be made.
Since the Schrödinger equation is only valid for nonrelativistic quantum mechanics, the solutions it yields for the hydrogen atom are not entirely correct. The Dirac equation of relativistic quantum theory improves these solutions (see below).
Solution of Schrödinger equation
The solution of the Schrödinger equation (wave equation) for the hydrogen atom uses the fact that the Coulomb potential produced by the nucleus is isotropic (it is radially symmetric in space and only depends on the distance to the nucleus). Although the resulting energy eigenfunctions (the orbitals) are not necessarily isotropic themselves, their dependence on the angular coordinates follows completely generally from this isotropy of the underlying potential: the eigenstates of the Hamiltonian (that is, the energy eigenstates) can be chosen as simultaneous eigenstates of the angular momentum operator. This corresponds to the fact that angular momentum is conserved in the orbital motion of the electron around the nucleus. Therefore, the energy eigenstates may be classified by two angular momentum quantum numbers, ℓ and m (both are integers). The angular momentum quantum number ℓ = 0, 1, 2, ... determines the magnitude of the angular momentum. The magnetic quantum number m = −ℓ, ..., +ℓ determines the projection of the angular momentum on the (arbitrarily chosen) zaxis.
In addition to mathematical expressions for total angular momentum and angular momentum projection of wavefunctions, an expression for the radial dependence of the wave functions must be found. It is only here that the details of the 1/r Coulomb potential enter (leading to Laguerre polynomials in r). This leads to a third quantum number, the principal quantum number n = 1, 2, 3, .... The principal quantum number in hydrogen is related to the atom's total energy.
Note that the maximum value of the angular momentum quantum number is limited by the principal quantum number: it can run only up to n − 1, i.e. ℓ = 0, 1, ..., n − 1.
Due to angular momentum conservation, states of the same ℓ but different m have the same energy (this holds for all problems with rotational symmetry). In addition, for the hydrogen atom, states of the same n but different ℓ are also degenerate (i.e. they have the same energy). However, this is a specific property of hydrogen and is no longer true for more complicated atoms which have an (effective) potential differing from the form 1/r (due to the presence of the inner electrons shielding the nucleus potential).
Taking into account the spin of the electron adds a last quantum number, the projection of the electron's spin angular momentum along the zaxis, which can take on two values. Therefore, any eigenstate of the electron in the hydrogen atom is described fully by four quantum numbers. According to the usual rules of quantum mechanics, the actual state of the electron may be any superposition of these states. This explains also why the choice of zaxis for the directional quantization of the angular momentum vector is immaterial: an orbital of given ℓ and m′ obtained for another preferred axis z′ can always be represented as a suitable superposition of the various states of different m (but same l) that have been obtained for z.
Alternatives to the Schrödinger theory
In the language of Heisenberg's matrix mechanics, the hydrogen atom was first solved by Wolfgang Pauli^{[4]} using a rotational symmetry in four dimension [O(4)symmetry] generated by the angular momentum and the Laplace–Runge–Lenz vector. By extending the symmetry group O(4) to the dynamical group O(4,2), the entire spectrum and all transitions were embedded in a single irreducible group representation.^{[5]}
In 1979 the (non relativistic) hydrogen atom was solved for the first time within Feynman's path integral formulation of quantum mechanics.^{[6]}^{[7]} This work greatly extended the range of applicability of Feynman's method.
Mathematical summary of eigenstates of hydrogen atom
In 1928, Paul Dirac found an equation that was fully compatible with Special Relativity, and (as a consequence) made the wave function a 4component "Dirac spinor" including "up" and "down" spin components, with both positive and "negative" energy (or matter and antimatter). The solution to this equation gave the following results, more accurate than the Schrödinger solution.
Energy levels
The energy levels of hydrogen, including fine structure (excluding Lamb shift and hyperfine structure), are given by the Sommerfeld fine structure expression:^{[8]}
where α is the finestructure constant and j is the "total angular momentum" quantum number, which is equal to ℓ ± ^{1}⁄_{2} depending on the direction of the electron spin. This formula represents a small correction to the energy obtained by Bohr and Schrödinger as given above. The factor in square brackets in the last expression is nearly one; the extra term arises from relativistic effects (for details, see #Features going beyond the Schrödinger solution). It is worth noting that this expression was first obtained by A. Sommerfeld in 1916 based on the relativistic version of the old Bohr theory. Sommerfeld has however used different notation for the quantum numbers.
The value
is called the Rydberg constant and was first found from the Bohr model as given by
where m_{e} is the electron mass, e is the elementary charge, h is the Planck constant, and ε_{0} is the vacuum permittivity.
This constant is often used in atomic physics in the form of the Rydberg unit of energy:
 ^{[9]}
The exact value of the Rydberg constant above assumes that the nucleus is infinitely massive with respect to the electron. For hydrogen1, hydrogen2 (deuterium), and hydrogen3 (tritium) the constant must be slightly modified to use the reduced mass of the system, rather than simply the mass of the electron. However, since the nucleus is much heavier than the electron, the values are nearly the same. The Rydberg constant R_{M} for a hydrogen atom (one electron), R is given by: where M is the mass of the atomic nucleus. For hydrogen1, the quantity is about 1/1836 (i.e. the electrontoproton mass ratio). For deuterium and tritium, the ratios are about 1/3670 and 1/5497 respectively. These figures, when added to 1 in the denominator, represent very small corrections in the value of R, and thus only small corrections to all energy levels in corresponding hydrogen isotopes.
Wavefunction
The normalized position wavefunctions, given in spherical coordinates are:
where:

 ,
 is the Bohr radius,
 is a generalized Laguerre polynomial of degree n − ℓ − 1, and
 is a spherical harmonic function of degree ℓ and order m. Note that the generalized Laguerre polynomials are defined differently by different authors. The usage here is consistent with the definitions used by Messiah,^{[10]} and Mathematica.^{[11]} In other places, the Laguerre polynomial includes a factor of ,^{[12]} or the generalized Laguerre polynomial appearing in the hydrogen wave function is instead.^{[13]}
The quantum numbers can take the following values:
Additionally, these wavefunctions are normalized (i.e., the integral of their modulus square equals 1) and orthogonal:
where is the state represented by the wavefunction in Dirac notation, and is the Kronecker delta function.^{[14]}
The wavefunctions in momentum space are related to the wavefunctions in position space through a Fourier transform
which, for the bound states, results in ^{[15]}
where denotes a Gegenbauer polynomial and is in units of .
Angular momentum
The eigenvalues for Angular momentum operator:
Visualizing the hydrogen electron orbitals
The image to the right shows the first few hydrogen atom orbitals (energy eigenfunctions). These are crosssections of the probability density that are colorcoded (black represents zero density and white represents the highest density). The angular momentum (orbital) quantum number ℓ is denoted in each column, using the usual spectroscopic letter code (s means ℓ = 0, p means ℓ = 1, d means ℓ = 2). The main (principal) quantum number n (= 1, 2, 3, ...) is marked to the right of each row. For all pictures the magnetic quantum number m has been set to 0, and the crosssectional plane is the xzplane (z is the vertical axis). The probability density in threedimensional space is obtained by rotating the one shown here around the zaxis.
The "ground state", i.e. the state of lowest energy, in which the electron is usually found, is the first one, the 1s state (principal quantum level n = 1, ℓ = 0).
An image with more orbitals is also available (up to higher numbers n and ℓ).
Black lines occur in each but the first orbital: these are the nodes of the wavefunction, i.e. where the probability density is zero. (More precisely, the nodes are spherical harmonics that appear as a result of solving Schrödinger's equation in polar coordinates.)
The quantum numbers determine the layout of these nodes.^{[16]} There are:
 total nodes,
 of which are angular nodes:
 angular nodes go around the axis (in the xy plane). (The figure above does not show these nodes since it plots crosssections through the xzplane.)
 (the remaining angular nodes) occur on the (vertical) axis.
 (the remaining nonangular nodes) are radial nodes.
Features going beyond the Schrödinger solution
There are several important effects that are neglected by the Schrödinger equation and which are responsible for certain small but measurable deviations of the real spectral lines from the predicted ones:
 Although the mean speed of the electron in hydrogen is only 1/137th of the speed of light, many modern experiments are sufficiently precise that a complete theoretical explanation requires a fully relativistic treatment of the problem. A relativistic treatment results in a momentum increase of about 1 part in 37,000 for the electron. Since the electron's wavelength is determined by its momentum, orbitals containing higher speed electrons show contraction due to smaller wavelengths.
 Even when there is no external magnetic field, in the inertial frame of the moving electron, the electromagnetic field of the nucleus has a magnetic component. The spin of the electron has an associated magnetic moment which interacts with this magnetic field. This effect is also explained by special relativity, and it leads to the socalled spinorbit coupling, i.e., an interaction between the electron's orbital motion around the nucleus, and its spin.
Both of these features (and more) are incorporated in the relativistic Dirac equation, with predictions that come still closer to experiment. Again the Dirac equation may be solved analytically in the special case of a twobody system, such as the hydrogen atom. The resulting solution quantum states now must be classified by the total angular momentum number j (arising through the coupling between electron spin and orbital angular momentum). States of the same j and the same n are still degenerate. Thus, direct analytical solution of Dirac equation predicts 2S(^{1}⁄_{2}) and 2P(^{1}⁄_{2}) levels of Hydrogen to have exactly the same energy, which is in a contradiction with observations (LambRetherford experiment).
 There are always vacuum fluctuations of the electromagnetic field, according to quantum mechanics. Due to such fluctuations degeneracy between states of the same j but different l is lifted, giving them slightly different energies. This has been demonstrated in the famous LambRetherford experiment and was the starting point for the development of the theory of Quantum electrodynamics (which is able to deal with these vacuum fluctuations and employs the famous Feynman diagrams for approximations using perturbation theory). This effect is now called Lamb shift.
For these developments, it was essential that the solution of the Dirac equation for the hydrogen atom could be worked out exactly, such that any experimentally observed deviation had to be taken seriously as a signal of failure of the theory.
Due to the high precision of the theory also very high precision for the experiments is needed, which utilize a frequency comb.
See also
References
 ↑ Palmer, D. (13 September 1997). "Hydrogen in the Universe". NASA. Retrieved 5 February 2008.
 ↑ Olsen, James; McDonald, Kirk (March 7, 2005). "Classical Lifetime of a Bohr Atom" (PDF). Joseph Henry Laboratories, Princeton University. Retrieved 12/10/2015. Check date values in:
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(help)  ↑ "Derivation of Bohr’s Equations for the Oneelectron Atom" (PDF). University of Massachusetts Boston. Retrieved 12/10/2015. Check date values in:
accessdate=
(help)  ↑ Pauli, W (1926). "Über das Wasserstoffspektrum vom Standpunkt der neuen Quantenmechanik". Zeitschrift für Physik. 36: 336–363. Bibcode:1926ZPhy...36..336P. doi:10.1007/BF01450175.
 ↑ Kleinert H. (1968). "Group Dynamics of the Hydrogen Atom" (PDF). Lectures in Theoretical Physics, edited by W.E. Brittin and A.O. Barut, Gordon and Breach, N.Y. 1968: 427–482. line feed character in
journal=
at position 24 (help)  ↑ Duru I.H., Kleinert H. (1979). "Solution of the path integral for the Hatom" (PDF). Physics Letters B. 84 (2): 185–188. Bibcode:1979PhLB...84..185D. doi:10.1016/03702693(79)902806.
 ↑ Duru I.H., Kleinert H. (1982). "Quantum Mechanics of HAtom from Path Integrals" (PDF). Fortschr. Phys. 30 (2): 401–435. Bibcode:1982ForPh..30..401D. doi:10.1002/prop.19820300802.
 ↑ Sommerfeld, Arnold (1919). Atombau und Spektrallinien'. Braunschweig: Friedrich Vieweg und Sohn. ISBN 3871444847. German English
 ↑ P.J. Mohr, B.N. Taylor, and D.B. Newell (2011), "The 2010 CODATA Recommended Values of the Fundamental Physical Constants" (Web Version 6.0). This database was developed by J. Baker, M. Douma, and S. Kotochigova. Available: http://physics.nist.gov/constants. National Institute of Standards and Technology, Gaithersburg, MD 20899. Link to R_{∞}, Link to hcR_{∞}
 ↑ Messiah, Albert (1999). Quantum Mechanics. New York: Dover. p. 1136. ISBN 0486409244.
 ↑ LaguerreL. Wolfram Mathematica page
 ↑ Griffiths, David (1995). Introduction to Quantum Mechanics. New Jersey: Pearson Education, Inc. p. 152. ISBN 0131118927.
 ↑ Condon and Shortley (1963). The Theory of Atomic Spectra. London: Cambridge. p. 441.
 ↑ Introduction to Quantum Mechanics, Griffiths 4.89
 ↑ Physics of atoms and molecules, B. H. Bransden and C. H. Joachain. Appendix 5
 ↑ Summary of atomic quantum numbers. Lecture notes. 28 July 2006
Books
 Griffiths, David J. (1995). Introduction to Quantum Mechanics. Prentice Hall. ISBN 0131118927. Section 4.2 deals with the hydrogen atom specifically, but all of Chapter 4 is relevant.
 Bransden, B.H.; C.J. Joachain (1983). Physics of Atoms and Molecules. Longman. ISBN 0582444012.
 Kleinert, H. (2009). Path Integrals in Quantum Mechanics, Statistics, Polymer Physics, and Financial Markets, 4th edition, Worldscibooks.com, World Scientific, Singapore (also available online physik.fuberlin.de)
External links
 Physics of hydrogen atom on Scienceworld
 Interactive graphical representation of orbitals
 Applet which allows viewing of all sorts of hydrogenic orbitals
 The Hydrogen Atom: Wave Functions, and Probability Density "pictures"
 Basic Quantum Mechanics of the Hydrogen Atom
 "Research team takes image of hydrogen atom" Kyodo News, Friday, 5 November 2010 – (includes image)
Lighter: (none, lightest possible) 
Hydrogen atom is an isotope of hydrogen 
Heavier: hydrogen2 
Decay product of: neutronium1 helium2 
Decay chain of hydrogen atom 
Decays to: Stable 