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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As we saw in the case of neon, the 2nd shell orbitals are more like spherical tetrahedra, and the 3rd shell is a single electron in a spherical s-orbital. These orbitals represent phase-locked, resonant, coherent, harmonic, stationary waves.

Sodium will give up its valence electron readily in an ionic interaction in order to reach the stability of a full 2nd shell. This is the same electron configuration as the 2s22p6 noble gas configuration of neon — 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, and it occurs between a metal (from the left side of the periodic table) and a non-metal (from the right side). The term “salt” can also be used to apply to any ionic crystal.

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. As each ion leaves the crystal, it becomes hydrated — surrounded by water molecules.

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

When the water is allowed to evaporate from the salt solution, the ions become increasingly exposed to one another, and the solid crystals re-form due to electrostatic attraction.


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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. With 6 electrons in a 2p orbital, neon is believed to achieve stability with octahedral symmetry, and its 2s electrons can unhybridize and return to their preferred spherical di-electron state (see below). It is proposed in the theory of Sub-Quantum Chemistry, however, that neon is more stable with a tetrahedral valence shell of 4 di-electrons (shown here).

In either case, its high degree of symmetry, with all electrons paired, renders neon, like helium, unable to bond, and therefore chemically unreactive.

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Intuitively, it seems more stable to consider the electron structure of neon (and of the full n=2 shell) as a tetrahedral arrangement of four di-electrons since 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 (left) compared to the unhybridized version (center) and the traditional lobe view (right)

The small 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. The outer spheres indicate the center of each orbital’s focus and region of highest electron density. The orbitals themselves are more like spherical tetrahedra that can only occupy volume within their shell. The entire shell will be filled with electron density. It will be highest at the center of the face of each orbital (as in the traditional sp3 lobe shapes) and will decrease toward the nodal regions between orbitals, where electron density will be lowest (though not necessarily zero). These lobes represent phase-locked, resonant, coherent, harmonic, stationary waves.

Neon’s four 2sp3-orbitals surrounding a 1s2 core di-electron shell

In the case of argon (Ar), since there are two concentric tetrahedral shells, they will align to form an antiparallel nested tetrahedral geometry because this minimizes 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. This is therefore the lowest energy state that nested tetrahedral shells can achieve.

Argon’s nested full-shell tetrahedra (left), and with one outer orbital raised (right)
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Unlike fluorine, neon is one of the two most unreactive elements on the periodic table. One of its primary uses (along with other noble gases) is as a gas in lighting (glow discharge) tubes, and it gives neon signs their characteristic bright red color.


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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 extremely (and dangerously) reactive in search of that final electron-pairing. One more electron will give it a full 2nd 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, and because it is the smallest of the Group VII elements, its electronegativity is also the highest on the periodic table. (The wireframe indicates the boundary of the n=2 shell.)

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The small spheres above simply indicate the directions of maximum electron density. The orbitals themselves will be more like spherical tetrahedra that can only occupy volume within their shell. The entire shell will be filled with electron density. It will be highest at the center of the face of each orbital (as in the traditional sp3 lobe shapes) and will decrease toward the nodal regions between orbitals, where electron density will be lowest (though not necessarily zero). In the case of fluorine, the three orbitals containing di-electrons will each occupy slightly more volume than the one containing the unpaired electron.

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 the larger (and usually-unreactive) noble gases, forcing them to donate electrons into that bond. Fluorine can therefore make a single covalent bond, achieving the same electron configuration as the 2s22p6 noble gas configuration of neon — a multi-di-electron state with two concentric full shells.

Fluorine can also gain an electron in an ionic interaction in order to reach the stability of a full 2nd shell. That is why fluorine forms a 1 ionic state. The negative ion is larger than the neutral atom because electrons now outnumber protons by one. 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)

Fluorine gas (F2) is so reactive that if it is simply passed over carbon, the carbon will spontaneously combust in it. In contrast, if we pass oxygen gas over carbon, it will only combust if ignited.

Although hydrofluoric acid (HF) is a weak acid, it cannot be stored in a glass container because it will degrade the glass due to fluorine’s high reactivity.


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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 indicates the boundary of the n=2 shell.)

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The small spheres above simply indicate the directions of maximum electron density. The orbitals themselves will be more like spherical tetrahedra that can only occupy volume within their shell. The entire shell will be filled with electron density. It will be highest at the center of the face of each orbital (as in the traditional sp3 lobe shapes) and will decrease toward the nodal regions between orbitals, where electron density will be lowest (though not necessarily zero). In the case of oxygen, the orbitals containing the two ‘di-electron’ lone pairs (dark blue) will each occupy slightly more volume than the two containing the unpaired electrons.

Oxygen is keen to obtain two extra electrons to fill its 2nd shell and achieve maximum stability. That would be the same electron configuration as the 2s22p6 noble gas configuration of neon — a multi-di-electron state with two concentric full shells. One way oxygen can do this is by making two covalent bonds. (See examples below.) Oxygen is the second most electronegative element (after fluorine), and it can therefore be very reactive as it strongly attracts electrons to itself.

Oxygen can also fill its 2nd shell by gaining two electrons in an ionic interaction (with a metal atom). That is why oxygen forms a 2 ionic state. The negative ion is much larger than the neutral atom 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 (H2O)

When oxygen gains two electrons via covalent bonding with two hydrogen atoms, the water (H2O) molecule is formed. Its asymmetrical structure, and the fact that oxygen pulls electrons more strongly than hydrogen, 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. This 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 (like oils) 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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Combustion Reactions

Oxygen has a high electronegativity, which means that oxygen atoms pull electrons strongly toward themselves. This causes them to react and bond (share electrons) with other atoms in order to gain some of their electron density. It is difficult for an oxygen atom to do that when bonding to another oxygen atom, though, because they are both competing for the shared electrons equally strongly. But when oxygen is near other atoms, which almost all hold their electrons less tightly, the oxygen atoms will prefer to bond with them because they will be able to attract the other atom’s electron density more successfully.

When atoms combine with oxygen, we call the process combustion because it usually gives off so much heat that it manifests fire or an explosion. An example of this reaction occurs when we burn carbon or hydrocarbon compounds in air.

When natural gas (methane CH4) burns (as we see in the image above), oxygen atoms from the oxygen (O2) molecules in the air let go of one another and instead bond to all of the available carbon and hydrogen atoms in the methane because their electrons are easier to attract. Water vapor (H2O) and carbon dioxide (CO2) are formed as a result. The chemical equation for this reaction is:

CH4 + 2O2 —> 2H2O + CO2

Other common combustion reactions occur in the formation of water (H2O) from hydrogen (H2) and oxygen (O2) and the combustion of octane (C8H18) in (the soon-to-be-obsolete) gasoline car engines.


Dioxygen (O2)

The dioxygen (O2) molecule (usually just called “oxygen”) has a surprisingly strong paramagnetism, which means it is strongly attracted into a magnetic field (as shown in the video below). This effect implies that there must be unpaired electrons in the molecule. Unpaired electrons can align themselves with an external magnetic field, and this causes them to be attracted into it. However, the oxygen molecule seems, at first glance, to have no unpaired electrons, so its strong observed paramagnetism needs to be explained.

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

The strength of this magnetic effect is called magnetic susceptibility (χm). By way of perspective, oxygen gas (O2(g)) has a magnetic susceptibility of χm=3,449, which is stronger than the rare earth metal cerium (χm=2,450). Liquid oxygen (O2(l)) has χm=7,699, which is stronger than the rare earth metal neodymium (χm=5,628) — which has four unpaired f-electrons. Furthermore, oxygen gas has more than six times the paramagnetism of manganese (χm=529), which has five unpaired d-electrons.

The Molecular Orbital Theory attributes the paramagnetism of dioxygen (O2) to the presence of two unpaired electrons in anti-bonding molecular orbitals (see Triplet Oxygen). An alternate (and at this point speculative) explanation for this phenomenon may emerge from the new theory of Sub-Quantum Chemistry, in which a set of quantum electron interactions unique to the oxygen molecule gives rise to a stronger than expected magnetic susceptibility (χm) value.


Ozone (O3)

A different naturally occurring form (or allotrope) of oxygen is ozone (O3), which consists of three oxygen atoms bonded to each other in a bent shape.

Ozone is formed in the stratosphere when incoming ultraviolet light hits oxygen molecules, and ozone is also broken back down into oxygen when it is hit by ultraviolet light. The ozone layer therefore protects us from harmful radiation by acting as a UV shield in two mutually-reinforcing ways.

Ozone is a molecule that demonstrates an interesting electron state known as resonance. Resonance occurs when the electrons have more than one way to make the same structure, though the structure is not symmetrical overall.

The two equivalent Lewis Dot Structures for ozone (O3)

Both versions of this structure feature a single bond, a double bond, and 6 di-electrons. They are exact mirror images of one another, and the traditional understanding of this in chemistry is that the actual structure is some form of average between these two structures. Since that is difficult to depict, we show both alternatives and we imagine the resulting average.

In Sub-Quantum Chemistry, Lewis Dot Structure is extended in order to incorporate all of the quantum interactions between electrons, especially those of quantum spin. Spin Lewis Structure then allows us to depict a single symmetrical resonance state for ozone, and according to this new theory, it looks like this:

The Spin Lewis Structure (right) and space-filling structure (left) for singlet ozone (O3), according to Sub-Quantum Chemistry theory.

More detail about Spin Lewis Structure (and ozone’s resonance structure) will be released soon.


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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 indicates the boundary of the n=2 shell.)

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The small spheres above simply indicate the directions of maximum electron density. The orbitals themselves will be more like spherical tetrahedra that can only occupy volume within their shell. The entire shell will be filled with electron density. It will be highest at the center of the face of each orbital (as in the traditional sp3 lobe shapes) and will decrease toward the nodal regions between orbitals, where electron density will be lowest (though not necessarily zero). In the case of nitrogen, the orbital containing the ‘di-electron’ lone pair (dark blue) will occupy slightly more volume than the three containing an unpaired electron.

Nitrogen is keen to obtain three extra electrons to fill its 2nd shell and achieve maximum stability. That would be the same electron configuration as the 2s22p6 noble gas configuration of neon — a multi-di-electron state with two concentric full shells. One way nitrogen can do this is by making three covalent bonds. An important example of this is the triple bond that nitrogen atoms make with each other when forming the nitrogen molecule (N2).

A triple bond forms when the three unpaired electrons on one nitrogen atom each pair with one of the unpaired electrons on the other nitrogen atom. (The pairs of dots represent the remaining di-electrons on each atom.)

A triple bond is a strong bond, and this makes it more difficult for nitrogen molecules to be broken apart. This causes more of them to persist, and this is the reason the Earth’s atmosphere contains mostly nitrogen molecules (78%). Oxygen molecules are bound with a weaker double bond, and they are therefore more reactive, more easily separated, and therefore a smaller portion of them (21%) remains in the atmosphere.

Nitrogen can also fill its 2nd shell by gaining three electrons in ionic interaction. That is why nitrogen forms a 3 ionic state. The negative ion is much larger than the neutral atom 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, which expands the size of the electron shell as it is attracted inward with less force.

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 (nitrogen) side slightly negative (𝛿–). The presence of the di-electron (lone pair) on the nitrogen atom also makes this molecule alkaline (basic) in solution.

Formation of the ammonia (NH3) molecule
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The ammonia molecule is vitally important in the manufacture of fertilizer. The invention of a process to make it efficiently (by Fritz Haber) revolutionized humanity’s ability to provide enough agricultural produce to feed our growing global population.


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6. Carbon

Carbon has two electrons in its p-orbital. These are not sphere-shaped harmonics and, like boron, carbon’s p-electrons cannot achieve symmetry around a sphere on their own because they lie in orthogonal directions.

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

Carbon therefore cannot simply add its two p-electrons on top of the same (2s2) configuration that beryllium has, as shown above, because it would not be a stable configuration. The asymmetry therefore causes the two electrons in the 2s orbital to uncouple from their di-electron state and hybridize with the two p-electrons in order to achieve 4-directional, tetrahedral symmetry. This yields four electrons with equivalent (degenerate) energy. Carbon is therefore the first atom that can achieve tetrahedral electron geometry on its own, which is why carbon typically makes 4 bonds. (The wireframe indicates the boundary of the n=2 shell.)

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The small spheres above simply indicate the directions of maximum electron density. The orbitals themselves will be more like spherical tetrahedra that can only occupy volume within their shell. The entire shell will be filled with electron density. It will be highest at the center of the face of each orbital (as in the traditional sp3 lobe shapes) and will decrease toward the nodal regions between orbitals, where electron density will be lowest (though not necessarily zero).

Carbon’s traditional sp3-orbital lobe shapes (left), and space-filling views (center), and exploded view (right)

Carbons small size, being only the 6th element, as well as the fact that it can bond in up to four directions, makes it one of the most versatile elements, structurally, on the periodic table. It can facilitate linear, trigonal, or tetrahedral electron geometry and it can make single, double, triple, or resonance bonds in linear, bent, or ring configurations.

Carbon can form structures as hard and beautiful as diamond, as strong as carbon fiber, as fine and versatile as graphene and nanotubes, as well as being the spindle around which the chemistry of life itself turns. It is almost single-handedly responsible for organic chemistry, which makes genetic information storage and biochemistry possible.

Methane (CH4)

When carbon bonds covalently with 4 hydrogen atoms, the symmetrical and non-polar methane (natural gas) molecule is formed.

The formation of methane (CH4) gas
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Methane (CH4) can be “oxidized” in air. A little instigating energy will cause molecular bonds to be destabilized just enough to allow the oxygen in the air to attack both the hydrogen and the carbon in the methane molecule. Since carbon and hydrogen both have a lower electronegativity than oxygen, they provide an easier source of electrons for oxygen atoms than oxygen atoms provide for each other. The result is that oxygen atoms from the oxygen molecules (O2) in the air let go of one another and instead bond to all of the carbon and hydrogen atoms in the methane. Water vapor (H2O) and carbon dioxide (CO2) are formed as a result. This is an example of a combustion (burning) reaction. The chemical equation for this reaction is:

CH4 + 2O2 —> 2H2O + CO2

Interestingly, it is then these two products (H2O & CO2) that plants take in and convert into sugar (C6H12O6) and oxygen (O2) through photosynthesis. As such, the more trees and green plants on the planet, the more carbon dioxide (CO2) is removed from the atmosphere. The more phytoplankton and kelp forests in the ocean, the more carbon dioxide (CO2) is removed from the water, and thereby indirectly from the atmosphere.

As carbon sequestration technologies are investigated, we should note that photosynthesis is the most effective carbon sequestration technology at hand. It lowers atmospheric carbon levels and provides valuable nutrients to the ecosystem.

If we fight to limit carbon emission but we do not stop global deforestation, our gain is offset by our loss, and our climate problems will worsen. Not to mention the rate of species extinctions, each of which plays a valuable role in trapping carbon into the soil and life cycle.


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5. Boron

Boron is the 5th element on the periodic table because it has 5 protons in its nucleus. With five electrons, boron is the first atom to contain electrons that are in a p-orbital. This is not a sphere-shaped harmonic, and so a single electron in a p-orbital cannot find a symmetrical arrangement around a sphere by itself. Boron therefore cannot simply add its p-electron on top of the same (2s2) configuration that beryllium has, as shown here, because it would not be stable:

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

The asymmetry therefore causes the two electrons in the 2s orbital to uncouple from their di-electron state and form a tri-electron state with the single p-electron. This is called hybridization. Three equal-energy (degenerate) electrons can now achieve maximum stability by assuming a trigonal planar (sp2-hybridized) arrangement around the core electron shell, because this minimizes their mutual repulsion by having them as far from one another as they can get. (The wireframe indicates the boundary of the n=2 shell.)

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The small spheres above simply indicate the directions of maximum electron density. The orbitals themselves will be more like three equal longitudinal sections that can only occupy volume within their shell. The entire shell will be filled with electron density. It will be highest at the center of the face of each orbital (as in the traditional sp2 lobe shapes) and will decrease toward the nodal regions between orbitals, where electron density will be lowest (though not necessarily zero).

Boron’s three 2sp2-orbitals surrounding a 1s2 core di-electron shell (left), with top view (right).

Each of these three hybrid orbitals contains one electron. This arrangement is symmetrical in the equatorial plane (in 2 dimensions) but does not have equivalent symmetry in all directions. This makes boron keen to connect with more electrons, because that would give the atom a 4th direction and tetrahedral symmetry (like carbon). Four directions is more stable around a sphere than three directions. Boron will therefore either bond ionically by losing its 3 valence electrons and forming the spherically symmetrical B3+ ion, or it will seek to gain electrons through covalent bonding. This duality underscores boron’s metalloid (semi-metal) character.

Covalent bonding, where another atom donates both electrons for the bond, allows boron to achieve tetrahedral geometry, which is why it forms adducts — combination molecules like B2H6 — and molecular structures like BH3NH3. It is also why boron can participate in (and be doped into) both trigonal and tetrahedral carbon crystals.

When boron atoms replace carbon atoms in a carbon crystal, they create points of lower electron density (or relative positivity) in the crystal because boron atoms contain 5 electrons while the carbon atoms surrounding them each have 6 electrons. (If nitrogen were doped into an adjacent region of the crystal, it would have points of relative negativity because nitrogen atoms contain 7 electrons.) Such doping gives a crystal important properties, which are leveraged, for example, when making a pn junction diode, used in electronics and photovoltaic (solar) cells.

p-n junction diode (Raffamaiden, CC BY-SA 3.0)

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4. Beryllium

Beryllium is the 4th element on the periodic table because it has 4 protons in its nucleus. It therefore has 4 electrons. The first two occupy an inner core shell, which is the 1s2 di-electron, just like helium’s. This first shell is surrounded by two more electrons in a 2nd shell s-orbital, forming a 2s2 di-electron. This makes beryllium less willing to lose an electron than lithium, giving beryllium a higher ionization energy and making it less reactive than lithium.

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This image provides another view of the spherical beryllium atom. The 2nd shell introduces a significantly larger volume to the atom, which is why it can accommodate four times as many electrons as the 1st shell.

While beryllium can become an ion, it does so less readily than magnesium (which has the same electron configuration but is one shell larger). Beryllium can be convinced to lose its two valence electrons to more electronegative atoms in an ionic interaction, and it will lose both at the same time in order to reach the stability of the 1s2 di-electron state, like helium. That is why beryllium forms a 2+ ion. The positive ion is much smaller than its neutral version because protons now outnumber electrons by two. This results in a much higher effective nuclear charge — a higher average attraction by the nucleus on each electron, which shrinks the size of the electron shell as it is attracted inward with more force.

Neutral beryllium (Be) atom (left) compared to the much smaller beryllium (Be2+) ion (right)

Beryllium is also the only metal in Group II that is diamagnetic — meaning that it repels away from any type of magnetic field. This may be due to the fact that the 2s2 di-electron state is small and well-bound, rendering the valence electrons less susceptible to aligning with an external magnetic field. (See Sub-Quantum Chemistry for more details.)


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OTHER GROUP II ELEMENTS: (Helium), Beryllium, Magnesium, Calcium, Strontium, Barium

3. Lithium

Lithium is the 3rd element on the periodic table because it has 3 protons in its nucleus. Lithium therefore has three electrons. The first two occupy an inner core shell, which is the 1s2 di-electron, just like helium’s. This first shell is surrounded by a single electron in a 2nd shell s-orbital. This unpaired 2nd shell electron makes lithium somewhat reactive since it would rather all its electrons were paired. It will therefore seek to lose this unpaired 2s1 electron in an ionic interactions (see below).

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This image provides another view of the spherical lithium atom. The 2nd shell introduces a significantly larger volume to the atom, which is why it can accommodate four times as many electrons as the 1st shell.

When lithium loses its valence electron in an ionic interaction, it achieves the stability of the 1s2 di-electron state, the highly stable electron state we find in helium. That is why lithium forms a 1+ ion. The positive ion it becomes is smaller than the neutral atom because protons now outnumber electrons by one. This results in a higher effective nuclear charge — a higher average attraction by the nucleus on each electron, which shrinks the size of the electron shell as it is attracted inward with more force.

Neutral lithium (Li) atom (left) compared to the much smaller Li+ ion (right)

In water, pure lithium reacts mildly as it forms the 1+ ion and dissolves. Since lithium has a lower electronegativity that hydrogen, it is an easier source of electrons for the highly electronegative oxygen atom in the water. Each water molecule (H2O) therefore ejects a hydrogen atom in favor of a lithium atom, forming the alkaline lithium hydroxide solution. Bubbles of hydrogen gas (H2) are released from the water in the process:

2Li(s) + 2H2O(l) —> 2LiOH(aq) + H2(g)


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

2. Helium

Helium is the 2nd element on the periodic table because it has 2 protons in its nucleus. Since the protons carry a positive charge, they will attract 2 electrons in order to balance the charge. Helium is therefore a neutral atom overall, with two electrons surrounding two protons (and 2 neutrons) in the nucleus.

The full-color wireframe represents a pair of electrons — a di-electron. In this case they are forming a sphere-shaped s-orbital. Helium only has 1 electron shell (an s-orbital) containing 2 electrons. We describe it with an electron configuration of 1s2.

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The image below provides a different view of the spherical helium atom.

The size of the nucleus in the center is greatly exaggerated. If the electron cloud were the size of a large football stadium, the nucleus would be the size of a penny at the center of the field — barely visible

Helium is the smallest atom on the periodic table. It is smaller than hydrogen because it has twice as many protons in the nucleus attracting twice as many electrons inwards. This higher “effective nuclear charge” shrinks the atom’s size. Since helium’s electrons are the most closely bound to their nucleus, helium has the highest ionization energy (2,370 kJ/mol or 24.5 eV) and is consequently the most unreactive element on the periodic table. Like the di-electron that envelops and binds the hydrogen molecule (H2), helium’s two electrons form a very stable, spherical, di-electron state a boson state — where the two electron wave functions are completely superimposed upon one another. (See Understanding Electrons.)

Helium also contains the most stable nucleus — the alpha particle. Most other nuclei are composed of various combinations of alpha particles (see the Robinson Model of Nuclear Binding), which is why this is the only type of multi-nucleon structure that is ejected during radioactive decay.

Alpha particle structure according to the Robinson Model of Nuclear Binding

With two proton-neutron pairs, the helium nucleus is itself a di-boson, making the helium atom a very stable tri-boson state.


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SEE OTHER NOBLE GASES: Helium, Neon, Argon, Krypton, Xenon