Hydrogen atom

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Hydrogen atom
Hydrogen 1.svg

Hydrogen atom

Name, symbol protium,1H
Neutrons 0
Protons 1
Nuclide data
Natural abundance 99.985%
Isotope mass 1.007825 u
Spin 12+
Excess energy 7288.969± 0.001 keV
Binding energy 0.000± 0.0000 keV
Depiction of a hydrogen atom showing the diameter as about twice the Bohr model radius. (Image not to scale)

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, H2. "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.


The most abundant isotope, hydrogen-1, 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 half-life 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, H3O+, is meant, not a literal ionized single hydrogen atom. In that case, the acid transfers the proton to H2O to form H3O+.

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 two-body problem physical system which has yielded many simple analytical solutions in closed-form.

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]

t_{fall} \approx \frac{ a_0^3}{4 r_0^2 c} \approx 1.6 *10^{-11} s

Where a_0 is the Bohr radius and r_0 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:

  1. Electrons can only be in certain, discrete orbitals, thereby having a discrete radius and energy.
  2. Electrons do not emit radiation while in one of these stationary states.
  3. 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:

L = n \hbar where n = 1,2,3,...

and \hbar is Planck constant over 2 \pi. 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]

E_n = - \frac{  m_e e^4}{2 ( 4 \pi \epsilon_0)^2 \hbar^2 } \frac{1}{n^2}   ,

where m_e   is the electron mass, e   is the electron charge, \epsilon_0   is the electric permeability, and n   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 \alpha^2 \approx 10^{-5}   , where \alpha   is the fine-structure constant. The Bohr model assumed circular orbits where, as developed by Sommerfeld, elliptial orbits adds other quantum numbers besides n   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 quantum-mechanical 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 non-relativistic 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, and m (both are integers). The angular momentum quantum number = 0, 1, 2, ... determines the magnitude of the angular momentum. The magnetic quantum number m = −, ..., + determines the projection of the angular momentum on the (arbitrarily chosen) z-axis.

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. = 0, 1, ..., n − 1.

Due to angular momentum conservation, states of the same 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 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 z-axis, 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 z-axis for the directional quantization of the angular momentum vector is immaterial: an orbital of given 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 4-component "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]

\begin{array}{rl} E_{j\,n} & = -m_\text{e}c^2\left[1-\left(1+\left[\dfrac{\alpha}{n-j-\frac{1}{2}+\sqrt{\left(j+\frac{1}{2}\right)^2-\alpha^2}}\right]^2\right)^{-1/2}\right] \\ & \approx -\dfrac{m_\text{e}c^2\alpha^2}{2n^2} \left[1 + \dfrac{\alpha^2}{n^2}\left(\dfrac{n}{j+\frac{1}{2}} - \dfrac{3}{4} \right) \right] , \end{array}

where α is the fine-structure constant and j is the "total angular momentum" quantum number, which is equal to | ± 12| 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

\frac{m_{\text{e}} c^2\alpha^2}{2}  = \frac{0.51\,\text{MeV}}{2 \cdot 137^2} = 13.6 \,\text{eV}

is called the Rydberg constant and was first found from the Bohr model as given by

-13.6 \,\text{eV} = -\frac{m_{\text{e}} e^4}{8 h^2 \varepsilon_0^2},

where me is the electron mass, e is the elementary charge, h is the Planck constant, and ε0 is the vacuum permittivity.

This constant is often used in atomic physics in the form of the Rydberg unit of energy:

1 \,\text{Ry} \equiv h c R_\infty = 13.605\;692\;53(30) \,\text{eV}.[9]

The exact value of the Rydberg constant above assumes that the nucleus is infinitely massive with respect to the electron. For hydrogen-1, hydrogen-2 (deuterium), and hydrogen-3 (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 RM for a hydrogen atom (one electron), R is given by: R_M = \frac{R_\infty}{1+m_{\text{e}}/M}, where M is the mass of the atomic nucleus. For hydrogen-1, the quantity m_{\text{e}}/M, is about 1/1836 (i.e. the electron-to-proton 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.


The normalized position wavefunctions, given in spherical coordinates are:

 \psi_{n\ell m}(r,\vartheta,\varphi) = \sqrt {{\left (  \frac{2}{n a_0} \right )}^3 \frac{(n-\ell-1)!}{2n[(n+\ell)!]^3}} e^{- \rho / 2} \rho^{\ell} L_{n-\ell-1}^{2\ell+1}(\rho) Y_{\ell}^{m}(\vartheta, \varphi )
3D illustration of the eigenstate \psi_{4,3,1}. Electrons in this state are 45% likely to be found within the solid body shown.


 \rho = {2r \over {na_0}} ,
 a_0 is the Bohr radius,
 L_{n-\ell-1}^{2\ell+1}(\rho) is a generalized Laguerre polynomial of degree n − 1, and
 Y_{\ell}^{m}(\vartheta, \varphi ) \, is a spherical harmonic function of degree 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 (n+\ell)!,[12] or the generalized Laguerre polynomial appearing in the hydrogen wave function is  L_{n+\ell}^{2\ell+1}(\rho) 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:

\int_0^{\infty} r^2 dr\int_0^{\pi} \sin \vartheta d\vartheta \int_0^{2 \pi} d\varphi\; \psi^*_{n\ell m}(r,\vartheta,\varphi)\psi_{n'\ell'm'}(r,\vartheta,\varphi)=\langle n,\ell, m | n', \ell', m' \rangle = \delta_{nn'} \delta_{\ell\ell'} \delta_{mm'},

where | n, \ell, m \rangle is the state represented by the wavefunction  \psi_{n\ell m} in Dirac notation, and  \delta is the Kronecker delta function.[14]

The wavefunctions in momentum space are related to the wavefunctions in position space through a Fourier transform

 \phi(p, \vartheta_p, \varphi_p) = (2\pi\hbar)^{-3/2} \int e^{-i \vec{p} \cdot \vec{r} / \hbar} \psi(r,\vartheta,\varphi) dV,

which, for the bound states, results in [15]

 \phi(p, \vartheta_p, \varphi_p) = \sqrt{\frac{2}{\pi} \frac{(n-l-1)!}{(n+l)!}} n^2 2^{2l+2} l! \frac{n^l p^l}{(n^2 p^2 + 1)^{l+2}} C_{n-l-1}^{l+1}\left(\frac{n^2 p^2 - 1}{n^2 p^2 + 1}\right) Y_l^m({\vartheta_p, \varphi_p}),

where  C_N^\alpha(x) denotes a Gegenbauer polynomial and  p is in units of  \hbar/a_0 .

Angular momentum

The eigenvalues for Angular momentum operator:

 L^2\, | n, \ell, m\rangle = {\hbar}^2 \ell(\ell+1)\, | n, \ell, m \rangle
 L_z\, | n, \ell, m \rangle = \hbar m \,| n, \ell, m \rangle.

Visualizing the hydrogen electron orbitals

Probability densities through the xz-plane for the electron at different quantum numbers (, across top; n, down side; m = 0)

The image to the right shows the first few hydrogen atom orbitals (energy eigenfunctions). These are cross-sections of the probability density that are color-coded (black represents zero density and white represents the highest density). The angular momentum (orbital) quantum number is denoted in each column, using the usual spectroscopic letter code (s means  = 0, p means  = 1, d means  = 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 cross-sectional plane is the xz-plane (z is the vertical axis). The probability density in three-dimensional space is obtained by rotating the one shown here around the z-axis.

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, = 0).

An image with more orbitals is also available (up to higher numbers n and ).

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:

  • n-1 total nodes,
  • l of which are angular nodes:
    • m angular nodes go around the \phi axis (in the xy plane). (The figure above does not show these nodes since it plots cross-sections through the xz-plane.)
    • l-m (the remaining angular nodes) occur on the \theta (vertical) axis.
  • n - l - 1 (the remaining non-angular 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 so-called spin-orbit 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 two-body 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(12) and 2P(12) levels of Hydrogen to have exactly the same energy, which is in a contradiction with observations (Lamb-Retherford experiment).

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


  1. Palmer, D. (13 September 1997). "Hydrogen in the Universe". NASA. Retrieved 5 February 2008.<templatestyles src="Module:Citation/CS1/styles.css"></templatestyles>
  2. 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: |accessdate= (help)<templatestyles src="Module:Citation/CS1/styles.css"></templatestyles>
  3. "Derivation of Bohr's Equations for the One-electron Atom" (PDF). University of Massachusetts Boston. Retrieved 12/10/2015. Check date values in: |accessdate= (help)<templatestyles src="Module:Citation/CS1/styles.css"></templatestyles>
  4. Lua error in Module:Citation/CS1/Identifiers at line 47: attempt to index field 'wikibase' (a nil value).
  5. 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)<templatestyles src="Module:Citation/CS1/styles.css"></templatestyles>
  6. Lua error in Module:Citation/CS1/Identifiers at line 47: attempt to index field 'wikibase' (a nil value).
  7. Lua error in Module:Citation/CS1/Identifiers at line 47: attempt to index field 'wikibase' (a nil value).
  8. Sommerfeld, Arnold (1919). Atombau und Spektrallinien'. Braunschweig: Friedrich Vieweg und Sohn. ISBN 3-87144-484-7.<templatestyles src="Module:Citation/CS1/styles.css"></templatestyles> German English
  9. 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
  10. Messiah, Albert (1999). Quantum Mechanics. New York: Dover. p. 1136. ISBN 0-486-40924-4.<templatestyles src="Module:Citation/CS1/styles.css"></templatestyles>
  11. LaguerreL. Wolfram Mathematica page
  12. Griffiths, David (1995). Introduction to Quantum Mechanics. New Jersey: Pearson Education, Inc. p. 152. ISBN 0-13-111892-7.<templatestyles src="Module:Citation/CS1/styles.css"></templatestyles>
  13. Condon and Shortley (1963). The Theory of Atomic Spectra. London: Cambridge. p. 441.<templatestyles src="Module:Citation/CS1/styles.css"></templatestyles>
  14. Introduction to Quantum Mechanics, Griffiths 4.89
  15. Physics of atoms and molecules, B. H. Bransden and C. H. Joachain. Appendix 5
  16. Summary of atomic quantum numbers. Lecture notes. 28 July 2006


  • Griffiths, David J. (1995). Introduction to Quantum Mechanics. Prentice Hall. ISBN 0-13-111892-7.<templatestyles src="Module:Citation/CS1/styles.css"></templatestyles> 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 0-582-44401-2.<templatestyles src="Module:Citation/CS1/styles.css"></templatestyles>
  • 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.fu-berlin.de)

External links

(none, lightest possible)
Hydrogen atom is an
isotope of hydrogen
Decay product of:
Decay chain
of hydrogen atom
Decays to:

pl:Wodór atomowy