Question: Does a quantized electron bond make an atom more stable?
About the article "Quantum", section "Abstract".

Candidate passages:
The electrons in an atom are attracted to the protons in the nucleus by the electromagnetic force.
This force binds the electrons inside an electrostatic potential well surrounding the smaller nucleus, which means that an external source of energy is needed for the electron to escape.
The closer an electron is to the nucleus, the greater the attractive force.
Hence electrons bound near the center of the potential well require more energy to escape than those at greater separations.
Electrons, like other particles, have properties of both a particle and a wave.
The electron cloud is a region inside the potential well where each electron forms a type of three-dimensional standing wave—a wave form that does not move relative to the nucleus.
This behavior is defined by an atomic orbital, a mathematical function that characterises the probability that an electron appears to be at a particular location when its position is measured.
Only a discrete (or quantized) set of these orbitals exist around the nucleus, as other possible wave patterns rapidly decay into a more stable form.
Orbitals can have one or more ring or node structures, and they differ from each other in size, shape and orientation.
Each atomic orbital corresponds to a particular energy level of the electron.
The electron can change its state to a higher energy level by absorbing a photon with sufficient energy to boost it into the new quantum state.
Likewise, through spontaneous emission, an electron in a higher energy state can drop to a lower energy state while radiating the excess energy as a photon.
These characteristic energy values, defined by the differences in the energies of the quantum states, are responsible for atomic spectral lines.
The amount of energy needed to remove or add an electron—the electron binding energy—is far less than the binding energy of nucleons.
For example, it requires only 13.6 eV to strip a ground-state electron from a hydrogen atom, compared to 2.23 "million" eV for splitting a deuterium nucleus.
Atoms are electrically neutral if they have an equal number of protons and electrons.
Atoms that have either a deficit or a surplus of electrons are called ions.
Electrons that are farthest from the nucleus may be transferred to other nearby atoms or shared between atoms.
By this mechanism, atoms are able to bond into molecules and other types of chemical compounds like ionic and covalent network crystals. 
In chemistry, an unpaired electron is an electron that occupies an orbital of an atom singly, rather than as part of an electron pair.
As the formation of electron pairs is often energetically favourable, either in the form of a chemical bond or as a lone pair, unpaired electrons are relatively uncommon in chemistry, because an entity that carries an unpaired electron is usually rather reactive.
In organic chemistry they typically only occur briefly during a reaction on an entity called a radical; however, they play an important role in explaining reaction pathways.
Radicals are uncommon in s- and p-block chemistry, since the unpaired electron occupies a valence p orbital or an sp, sp2 or sp3 hybrid orbital.
These orbitals are strongly directional and therefore overlap to form strong covalent bonds, favouring dimerisation of radicals.
Radicals can be stable if dimerisation would result in a weak bond or the unpaired electrons are stabilised by delocalisation.
In contrast, radicals in d- and f-block chemistry are very common.
The less directional, more diffuse d and f orbitals, in which unpaired electrons reside, overlap less effectively, form weaker bonds and thus dimerisation is generally disfavoured.
These d and f orbitals also have comparatively smaller radial extension, disfavouring overlap to form dimers.
The more stable entities with unpaired electrons do exist, e.g.
the nitric oxide molecule has one.
According to Hund's rule, the spins of unpaired electrons are aligned parallel and this gives these molecules paramagnetic properties.
The most stable examples of unpaired electrons are found on the atoms and ions of lanthanides.
The incomplete f-shell of these entities does not interact very strongly with the environment they are in and this prevents them from being paired.
The ion with the largest number of unpaired electrons is Gd3+ with seven unpaired electrons. 
The electron affinity of an atom can be described either as the energy gained by an atom when an electron is added to it, or conversely as the energy required to detach an electron from a singly charged anion.
The sign of the electron affinity can be quite confusing, as atoms that become more stable with the addition of an electron (and so are considered to have a higher electron affinity) show a decrease in potential energy; i.e.
the energy gained by the atom appears to be negative.
For atoms that become less stable upon gaining an electron, potential energy increases, which implies that the atom gains energy.
In such a case, the atom's electron affinity value is positive.
Consequently, atoms with a more negative electron affinity value are considered to have a lower electron affinity (they are more receptive to gaining electrons), and vice versa.
However in the reverse scenario where electron affinity is defined as the energy required to detach an electron from an anion, the energy value obtained will be of the same magnitude but have the opposite sign.
This is because those atoms with a high electron affinity are less inclined to give up an electron, and so take more energy to remove the electron from the atom.
In this case, the atom with the more positive energy value has the higher electron affinity.
As one progresses from left to right across a period, the electron affinity will increase.
Although it may seem that Fluorine should have the greatest electron affinity, the small size of fluorine generates enough repulsion that Chlorine has the greatest electron affinity. 
The electron affinity of an atom is the amount of energy released when an electron is added to a neutral atom to form a negative ion.
Although electron affinity varies greatly, some patterns emerge.
Generally, nonmetals have more positive electron affinity values than metals.
Chlorine most strongly attracts an extra electron.
The electron affinities of the noble gases have not been measured conclusively, so they may or may not have slightly negative values.
Electron affinity generally increases across a period.
This is caused by the filling of the valence shell of the atom; a group 17 atom releases more energy than a group 1 atom on gaining an electron because it obtains a filled valence shell and is therefore more stable.
A trend of decreasing electron affinity going down groups would be expected.
The additional electron will be entering an orbital farther away from the nucleus.
As such this electron would be less attracted to the nucleus and would release less energy when added.
However, in going down a group, around one-third of elements are anomalous, with heavier elements having higher electron affinities than their next lighter congenors.
Largely, this is due to the poor shielding by d and f electrons.
A uniform decrease in electron affinity only applies to group 1 atoms. 
Bonds with one or three electrons can be found in radical species, which have an odd number of electrons.
The simplest example of a 1-electron bond is found in the dihydrogen cation, H2+.
One-electron bonds often have about half the bond energy of a 2-electron bond, and are therefore called "half bonds".
However, there are exceptions: in the case of dilithium, the bond is actually stronger for the 1-electron Li2+ than for the 2-electron Li2.
This exception can be explained in terms of hybridization and inner-shell effects.
The simplest example of three-electron bonding can be found in the helium dimer cation, He2+.
It is considered a "half bond" because it consists of only one shared electron (rather than two) in addition to one lone electron on each atom; in molecular orbital terms, the third electron is in an anti-bonding orbital which cancels out half of the bond formed by the other two electrons.
Another example of a molecule containing a 3-electron bond, in addition to two 2-electron bonds, is nitric oxide, NO.
The oxygen molecule, O2 can also be regarded as having two 3-electron bonds and one 2-electron bond, which accounts for its paramagnetism and its formal bond order of 2.
Chlorine dioxide and its heavier analogues bromine dioxide and iodine dioxide also contain three-electron bonds.
Molecules with odd-electron bonds are usually highly reactive.
These types of bond are only stable between atoms with similar electronegativities.