Atomic hydrogen
Hydrogen1  

 
General  
Name, symbol  protium, ^{1}H 
Neutrons  0 
Protons  1 
Nuclide data  
Natural abundance  99.985% 
Halflife  Stable 
Isotope mass  1.007825 u 
Spin  ½+ 
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 mass of the universe.^{[1]} (Most of the universe's mass is not in the form of chemical elements—that is, "baryonic" matter—but is made up of dark matter and dark energy.)
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).
Contents
 1 Production and reactivity
 2 Isotopes
 3 Quantum theoretical analysis
 4 See also
 5 References
 6 Books
 7 External links
Production and reactivity
The H–H bond is one of the toughest bonds in chemistry, with a bond dissociation enthalpy of 435.88 kJ/mol at 298 K (25 °C; 77 °F). As a consequence of this strong bond, H_{2} dissociates to only a minor extent until higher temperatures. At 3,000 K (2,730 °C; 4,940 °F), the degree of dissociation is just 7.85%:^{[2]}
 H_{2} ⇌ 2 H
Hydrogen atoms are so reactive that they combine with almost all elements.
Isotopes
The most abundant isotope, hydrogen1, protium, or light hydrogen, contains no neutrons; other isotopes of hydrogen, such as deuterium or tritium, contain one or more neutrons. 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.
Quantum 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.
In 1913, Niels Bohr obtained the spectral frequencies of the hydrogen atom after making a number of simplifying assumptions. These assumptions, the cornerstones of the Bohr model, were not fully correct but did yield fairly correct energy answers (with a relative error in the ground state ionization energy of around α^{2}/4 or around 10^{5}). Bohr's results for the frequencies and underlying energy values were duplicated by the solution to the Schrödinger equation in 1925–1926. The solution to the Schrödinger equation for hydrogen is analytical, giving a simple expressions for the hydrogen energy levels and thus the frequencies of the hydrogen spectral lines. The solution of the Schrödinger equation goes much further than the Bohr model, because it also yields 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.
The Schrödinger equation is not fully accurate. The next improvement was the Dirac equation (see below).
Solution of Schrödinger equation: Overview of results
The solution of the Schrödinger equation (wave equations) 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 a (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^{[3]} 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.^{[4]}
In 1979 the (non relativistic) hydrogen atom was solved for the first time within Feynman's path integral formulation of quantum mechanics.^{[5]}^{[6]} 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, are given by the Sommerfeld expression:
 $\backslash begin\{array\}\{rl\}\; E\_\{j\backslash ,n\}\; \&\; =\; m\_\backslash text\{e\}c^2\backslash left[\backslash left(1+\backslash left[\backslash dfrac\{\backslash alpha\}\{nj\backslash frac\{1\}\{2\}+\backslash sqrt\{\backslash left(j+\backslash frac\{1\}\{2\}\backslash right)^2\backslash alpha^2\}\}\backslash right]^2\backslash right)^\{1/2\}1\backslash right]\; \backslash \backslash \; \&\; \backslash approx\; \backslash dfrac\{m\_\backslash text\{e\}c^2\backslash alpha^2\}\{2n^2\}\; \backslash left[1\; +\; \backslash dfrac\{\backslash alpha^2\}\{n^2\}\backslash left(\backslash dfrac\{n\}\{j+\backslash frac\{1\}\{2\}\}\; \; \backslash dfrac\{3\}\{4\}\; \backslash right)\; \backslash right]\; ,\; \backslash end\{array\}$
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. 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).
The value
 $\backslash frac\{m\_\{\backslash text\{e\}\}\; c^2\backslash alpha^2\}\{2\}\; =\; \backslash frac\{0.51\backslash ,\backslash text\{MeV\}\}\{2\; \backslash cdot\; 137^2\}\; =\; 13.6\; \backslash ,\backslash text\{eV\}$
is called the Rydberg constant and was first found from the Bohr model as given by
 $13.6\; \backslash ,\backslash text\{eV\}\; =\; \backslash frac\{m\_\{\backslash text\{e\}\}\; e^4\}\{8\; h^2\; \backslash varepsilon\_0^2\},$
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:
 $1\; \backslash ,\backslash text\{Ry\}\; \backslash equiv\; h\; c\; R\_\backslash infty\; =\; 13.605\backslash ;692\backslash ;53(30)\; \backslash ,\backslash text\{eV\}.$^{[7]}
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: $R\_M\; =\; \backslash frac\{R\_\backslash infty\}\{1+m\_\{\backslash text\{e\}\}/M\},$ where M is the mass of the atomic nucleus. For hydrogen1, the quantity $m\_\{\backslash text\{e\}\}/M,$ 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:
 $\backslash psi\_\{n\backslash ell\; m\}(r,\backslash vartheta,\backslash varphi)\; =\; \backslash sqrt$,
 $a\_0$ is the Bohr radius,
 $L\_\{n\backslash ell1\}^\{2\backslash ell+1\}(\backslash rho)$ is a generalized Laguerre polynomial of degree n − ℓ − 1, and
 $Y\_\{\backslash ell\}^\{m\}(\backslash vartheta,\; \backslash varphi\; )\; \backslash ,$ 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,^{[8]} and Mathematica.^{[9]} In other places, the Laguerre polynomial includes a factor of $(n+\backslash ell)!$,^{[10]} or the generalized Laguerre polynomial appearing in the hydrogen wave function is $L\_\{n+\backslash ell\}^\{2\backslash ell+1\}(\backslash rho)$ instead. ^{[11]}
The quantum numbers can take the following values:
 $n=1,2,3,\backslash ldots$
 $\backslash ell=0,1,2,\backslash ldots,n1$
 $m=\backslash ell,\backslash ldots,\backslash ell.$
Additionally, these wavefunctions are normalized (i.e., the integral of their modulus square equals 1) and orthogonal:
 $\backslash int\_0^\{\backslash infty\}\; r^2\; dr\backslash int\_0^\{\backslash pi\}\; \backslash sin\; \backslash vartheta\; d\backslash vartheta\; \backslash int\_0^\{2\; \backslash pi\}\; d\backslash varphi\backslash ;\; \backslash psi^*\_\{n\backslash ell\; m\}(r,\backslash vartheta,\backslash varphi)\backslash psi\_\{n\text{'}\backslash ell\text{'}m\text{'}\}(r,\backslash vartheta,\backslash varphi)=\backslash langle\; n,\backslash ell,\; m\; \; n\text{'},\; \backslash ell\text{'},\; m\text{'}\; \backslash rangle\; =\; \backslash delta\_\{nn\text{'}\}\; \backslash delta\_\{\backslash ell\backslash ell\text{'}\}\; \backslash delta\_\{mm\text{'}\},$
where $\; n,\; \backslash ell,\; m\; \backslash rangle$ is the representation of the wavefunction $\backslash psi\_\{n\backslash ell\; m\}$ in Dirac notation, and $\backslash delta$ is the Kronecker delta function.^{[12]}
Angular momentum
The eigenvalues for Angular momentum operator:
 $L^2\backslash ,\; \; n,\; \backslash ell,\; m\backslash rangle\; =\; \{\backslash hbar\}^2\; \backslash ell(\backslash ell+1)\backslash ,\; \; n,\; \backslash ell,\; m\; \backslash rang$
 $L\_z\backslash ,\; \; n,\; \backslash ell,\; m\; \backslash rang\; =\; \backslash hbar\; m\; \backslash ,\; n,\; \backslash ell,\; m\; \backslash rang.$
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.^{[13]} There are:
 $n1$ total nodes,
 $l$ of which are angular nodes:
 $m$ angular nodes go around the $\backslash phi$ axis (in the xy plane). (The figure above does not show these nodes since it plots crosssections through the xzplane.)
 $lm$ (the remaining angular nodes) occur on the $\backslash theta$ (vertical) axis.
 $n\; \; l\; \; 1$ (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.
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.
See also

References
Books
 Section 4.2 deals with the hydrogen atom specifically, but all of Chapter 4 is relevant.
 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 
pl:Wodór atomowy