16. Sulfur

Sulfur has the same electron configuration as oxygen, though as a larger atom and one with a d-orbital into which it can hybridize, sulfur has different properties and can make different molecular geometries than oxygen. (The wireframe simply indicates the boundary of the n=3 shell, since there are no electrons defining the boundary of its sphere.)

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As the structure below indicates, it is common for sulfur to make two bonds. In its natural crystalline form, sulfur forms S8 rings where each sulfur atoms is bonded to two adjacent sulfur atoms. Like oxygen, sulfur can also form the 2– ion by gaining two electrons in an ionic interaction.

But if sulfur is approached by highly electronegative atoms like oxygen or fluorine, they can induce sulfur’s paired valence electrons to uncouple, hybridizing with the d-orbital, and allowing sulfur to have up to 6 unpaired electrons. Sulfur can therefore make up to 6 covalent bonds, as we see in the SF6 molecule.

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15. Phosphorus

Phosphorus has the same tetrahedral sp3 electron configuration as nitrogen, though, as a larger atom, the chemistry of phosphorus has important differences from that of nitrogen. With an available d-orbital into which to hybridize, phosphorus can achieve a greater variety of molecular geometries than can nitrogen. (The wireframe simply indicates the boundary of the n=3 shell, since there are no electrons defining the boundary of its sphere.)

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14. Silicon

Silicon has the same structure as carbon, a hybridized, tetrahedral sp3 configuration. Since it is a larger atom, the four degenerate n=3 electrons form a less coherent quantum system than they do in carbon. The larger size, along with the presence of a d-orbital in the 3rd shell into which to expand the electron octet, gives silicon different (more semi-metallic) properties than carbon. (The wireframe simply indicates the boundary of the n=3 shell, since there are no electrons defining the boundary of its sphere.)

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13. Aluminium

Aluminium has a similar electron configuration to boron, and may therefore form an sp2 trigonal planar configuration. Since it is larger than boron, its valence electrons fill a larger volume and are therefore less coherent as a single quantum system. This gives aluminium different properties, making it more metallic in nature than the semi-metallic boron. (The wireframe simply indicates the boundary of the n=3 shell, since there are no electrons defining the boundary of its sphere.)

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Aluminium can be convinced to lose its three valence electrons in an ionic interaction, and it will lose all three in order to reach the stability of the 2s22p6 noble gas configuration, which is a multi-di-electron state with two concentric full shells. That is why aluminium forms a 3+ ionic state.

Neutral aluminium (Al) atom (L) compared to the much smaller Al3+ ion (R)

When we compare aluminium to boron, we see a significant size difference (see below). We also see a significant difference between the magnetic susceptibility of aluminium versus the other Group III elements since aluminium is paramagnetic (χm=16.5) while the others are diamagnetic. Aluminium’s paramagnetic strength is also very similar to that of sodium (χm=16). This might imply that, in its crystalline metallic state, aluminium might delocalize only one of its three valence electrons, perhaps allowing the remaining two 3s-electrons to retain some of their preferred di-electron character, stabilizing the 3rd shell with spherical electron density.

Size comparison between boron (left) and aluminium (right)

Boron’s diamagmetism (χm=–6.7), on the other hand, would appear to emerge from the fact that it usually makes covalent bonds, in which all of its electrons are paired.

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12. Magnesium

Magnesium has a di-electron making up its 3rd shell, with two full core shells within that have the identical configuration to neon. Given its larger size, the 3s2 di-electron is not as well-bound as the 2s2 di-electron on beryllium. Magnesium will therefore donate its valence electrons more readily than beryllium, but it will not be as reactive in doing so as the alkali earth metals below it in Group II such as calcium, strontium, or barium.

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Magnesium is more willing to lose its two valence electrons in an ionic interaction than is beryllium, and it will lose both at the same time in order to reach the stability of the 2s22p6 noble gas configuration, which is a multi-di-electron state with two concentric full shells. That is why magnesium forms a 2+ ionic state.

Neutral magnesium (Mg) atom (L) compared to the much smaller Mg2+ ion (R)

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OTHER GROUP II ELEMENTS: Beryllium, Magnesium, Calcium, Strontium, Barium

11. Sodium

Sodium has a single valence electron in its 3rd shell with two full core shells within that have the identical configuration to neon. This makes sodium keen to donate its single valence electron in order to regain the electron symmetry of neon, resulting in its 1+ ionic character when interacting with other non-metal atoms. Being larger than lithium or hydrogen, the lower electrostatic force from the nucleus and the greater core electron shielding cause sodium’s valence electron to be more weakly bound, and this makes sodium more reactive than lithium. Pure sodium metal reacts violently, sometimes explosively, when placed in water as it donates its valence electron to oxygen. The heat of this reaction ignites the hydrogen gas that is also produced, burning with a yellow flame.

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Sodium will give up its valence electron readily in an ionic interaction in order to reach the stability of the 2s22p6 noble gas configuration, which is a multi-di-electron state with two concentric full shells. That is why sodium forms a 1+ ionic state.

Neutral sodium (Na) atom (L) compared to the much smaller Na+ ion (R)

Salt

When sodium and chlorine interact, sodium gives the electron it wants to lose to chlorine, which is keen to gain it. This forms both atoms into their ions and allows both to achieve full shell configurations. The ions can then stick to each other because of their opposite charges, forming sodium chloride (NaCl) crystals. This process is called ionic bonding.

Na + Cl prefer to become Na+ + Cl, which can then form NaCl.

Sodium chloride (NaCl) dissolves in water because the polar H2O molecules and the ions in the crystal attract each other. The water molecules can therefore tug ions off the crystal and still satisfy the ion’s desire to attract their opposite polarity.

Polar water (H2O) molecules dissolving salt (NaCl).

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OTHER GROUP I ELEMENTS: Lithium, Sodium, Potassium, Rubidium, Cesium

10. Neon

Neon has two full electron shells — an inner core 1s shell, and a 2s shell containing a full p-orbital resonance. For atoms with fewer than 6 p-electrons, tetrahedral symmetry is the most stable. With 6 electrons in a full 2p orbital, neon can achieve stability with octahedral symmetry, and its 2s electrons can unhybridize and return to their preferred spherical di-electron state.

This high degree of symmetry, with all electrons paired, renders neon, like most other noble gases, unreactive.

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Once again, the spheres in these diagrams are only meant to indicate the directions of the orbital lobes. These lobes are not spheres. Only s-orbitals are spherical. As we mentioned on the key shown on the periodic table graphic, a representation that would be closer to reality would need to show the lobes spreading out within their allowed shell as much as their mutual repulsion will allow. These lobes represent phase-locked, resonant, coherent, harmonic, stationary waves:

Neon (left) and only the p-orbitals (right)
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An alternative way to consider the electron structure of neon and of the full n=2 shell would be to continue the tetrahedral trend of the preceding atoms, and render neon as a tetrahedral arrangement of four di-electrons. Intuitively, that seems more stable as it involves a greater degree of field cancellation than a full p-orbital with a single electron occupying each lobe.

Tetrahedral view of neon’s orbital symmetry
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Note, again, that these small spheres in the second shell represent the center of each orbital’s focus and region of highest electron density, but the entire shell will be filled with electron density, save for the “double-layer” nodal regions that divide the four equal volumes of the shell, where electron density will be lowest. We might therefore approximate the maximum volume that each sp3 orbital di-electron resonance can occupy as one fourth of the volume of shell 2 (excluding the volume of shell 1).

Spherical tetrahedral platonic solid compared to sp3 hybrid orbital lobes.

What is intriguing about the nesting of these spherical tetrahedra is that they align in such a way as to minimize repulsion between layers. The place of lowest electron density — where the nodal vertices intersect — on one shell is set against the highest electron density at the center of a face on the adjacent concentric shell.

Nested full-shell tetrahedra (left), with tetrahedral argon (right) as an example..
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SEE OTHER NOBLE GASES: Helium, Neon, Argon, Krypton, Xenon

9. Fluorine

Fluorine has five electrons in its p-orbital. This is not a sphere-shaped harmonic, and so five p-electrons cannot achieve a stable electron symmetry around a spherical core. Fluorine therefore cannot simply add its five p-electrons on top of the same (2s2) configuration that beryllium has, as shown here, because it would not be a stable configuration:

Each p-orbital lobe holds 1 electron. An electron pair occupies two opposite lobes.

The asymmetry therefore causes fluorine to hybridize its 2s and 2p electrons in order to achieve tetrahedral symmetry. Its sp3 hybrid orbitals feature three di-electrons and one unpaired electron, rendering it very reactive in search of that final electron-pairing. One more electron will give it a full shell, like neon, and that is a very attractive state for the atom. In addition, a high effective nuclear charge gives fluorine the highest electronegativity in its row, but because it is the smallest of the Group VII elements, its electronegativity is also the highest on the periodic table. (The wireframe simply indicates the boundary of the n=2 shell, since there are no electrons defining the boundary of its sphere.)

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Fluorine is so eager to obtain an extra electron to fill its second shell that it can bond with just about any atom on the periodic table, even several of the usually-unreactive noble gases, forcing them to donate electrons into that bond. Fluorine can therefore make a single covalent bond or gain an electron in ionic interaction in order to reach the stability of the 2s22p6 noble gas configuration of neon — a multi-di-electron state with two concentric full shells. That is why fluorine forms a 1– ionic state. The negative ion is larger than its neutral version because electrons now outnumber protons. This results in a lower effective nuclear charge — a lower average attraction by the nucleus on each electron.

Neutral fluorine (F) atom (left) compared to the larger fluoride (F) ion (right)

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OTHER GROUP VII HALOGENS: Fluorine, Chlorine, Bromine, Iodine

8. Oxygen

Oxygen has four electrons in its p-orbital. This is not a sphere-shaped harmonic, and so four p-electrons cannot achieve a stable electron symmetry around a spherical core. Oxygen therefore cannot simply add its four p-electrons on top of the same (2s2) configuration that beryllium has, as shown here, because it would not be a stable configuration:

Each p-orbital lobe holds 1 electron. An electron pair occupies two opposite lobes.

The asymmetry therefore causes oxygen to hybridize its 2s and 2p electrons in order to achieve tetrahedral symmetry. Its sp3 hybrid orbitals feature two di-electrons and 2 unpaired, degenerate electrons. This is why oxygen typically makes 2 bonds, and it is this tetrahedral geometry that gives the water (H2O) molecule its characteristic bent shape (see below). (The wireframe simply indicates the boundary of the n=2 shell, since there are no electrons defining the boundary of its sphere.)

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Oxygen is keen to obtain two extra electrons to fill its second shell. It is the second most electronegative element after fluorine, and can therefore be quite reactive. Oxygen can make one or more covalent bonds or it can gain two electrons in ionic interaction in order to reach the stability of the 2s22p6 noble gas configuration of neon — a multi-di-electron state with two concentric full shells. That is why oxygen forms a 2– ionic state. The negative ion is much larger than its neutral version because electrons now outnumber protons by two. This results in a much lower effective nuclear charge — a lower average attraction by the nucleus on each electron.

Neutral oxygen (O) atom (left) compared to the larger oxide (O2–) ion (right)

Water

When oxygen gains two electrons via covalent bonding with two hydrogen atoms, the water (H2O) molecule is formed. Its asymmetrical structure gives it very important properties. The most significant is that it creates an electron imbalance which makes the side of the molecule where the hydrogen atoms attach slightly positive (𝛿+) and the opposite (oxygen) side slightly negative (𝛿–). This polarity causes water molecules to stick to each other and to certain substances rather effectively. It makes it harder to change water’s temperature, causing oceans and lakes to moderate local climates. It makes ice less dense than water, causing ice to float, ensuring that aquatic species survive the winter. It allows a great variety of substances (like NaCl) to dissolve in water, and it allows others to resist water and form cellular structures within a water environment. It is truly a remarkable molecule.

Formation of the water (H2O) molecule
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Dioxygen (O2)

The dioxygen (O2) molecule has a surprisingly strong paramagnetism. This means it is strongly attracted into a magnetic field.

Liquid dioxygen (O2) poured between the poles of an electromagnet are attracted to it. (Source: Harvard Natural Sciences Lecture Demonstrations, Youtube)

Gaseous dioxygen (O2(g)) has a magnetic susceptibility of χm=3,449, which is stronger than the rare earth metal cerium (χm=2,450). Liquid dioxygen (O2(l)) has a χm=7,699, which is stronger than the rare earth metal neodymium (χm=5,628). Furthermore, dioxygen gas has more than six times the paramagnetism of manganese, which has five unpaired electrons.

The Molecular Orbital Theory attributes the paramagnetism of dioxygen to the presence of two unpaired electrons in anti-bonding molecular orbitals. It is proposed, in the Quantum TORCH, that two oxygen atoms form a unique quantum symmetry when they bond, and that this unique electron arrangement gives rise to the molecule’s strong paramagnetism. (See the Quantum TORCH for more detail.)


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7. Nitrogen

Nitrogen has three electrons in its p-orbital. This is not a sphere-shaped harmonic, and so three p-electrons cannot achieve a stable electron symmetry around a spherical core. Nitrogen therefore cannot simply add its three p-electrons on top the same (2s2) configuration that beryllium has, as shown here, because it would not be a stable configuration:

Each p-orbital lobe holds 1 electron. An electron pair occupies two opposite lobes.

The asymmetry therefore causes the three p-orbital electrons to combine (hybridize) with the two s-orbital electrons in order to achieve 4-directional, symmetry. Nitrogen’s tetrahedral (sp3) symmetry features one di-electron (lone pair) and 3 unpaired, degenerate electrons. This is why nitrogen typically makes 3 bonds. The quantum states of the three degenerate electrons become linked, which stabilizes them and thereby increases nitrogen’s ionization energy. (The wireframe simply indicates the boundary of the n=2 shell, since there are no electrons defining the boundary of its sphere.)

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Nitrogen is keen to obtain three extra electrons to fill its second shell. It can make one or more covalent bonds or it can gain three electrons in ionic interaction in order to reach the stability of the 2s22p6 noble gas configuration of neon, which is a multi-di-electron state with two concentric full shells. That is why nitrogen forms a 3– ionic state. The negative ion is much larger than its neutral version because electrons now outnumber protons by three. This results in a much lower effective nuclear charge — a lower average attraction by the nucleus on each electron.

Neutral nitrogen (N) atom (left) compared to the larger nitride (N3–) ion (right)

Ammonia (NH3)

When nitrogen bonds covalently with three hydrogen atoms, the biologically important ammonia (NH3) molecule is formed. Like the water molecule, its asymmetrical structure gives it important properties. The most significant is that it creates an electron imbalance which makes the side of the molecule where the hydrogen atoms attach slightly positive (𝛿+) and the opposite (oxygen) side slightly negative (𝛿–). The presence of the di-electron on the nitrogen atom also makes this molecule alkaline (basic).

Formation of the ammonia (NH3) molecule
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