STRUCTURE AND BONDING OF THE ELEMENTS #3

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Hello, dear readers. Today, I will continue from where I stopped in my last post on Structure and bonding of the elements. I will be discussing about semiconductors, the solar cell and simple molecule lattices.
So, starting with the semiconductors….

Semiconductors

The biggest use for silicon is in semiconductor components for electronic circuits. You may wonder how it is that a non-metallic element with a diamond-like structure can be made to conduct electricity. Semiconductors are materials that, to put it simply, have been altered to allow electrons to pass through them, but not as easily as through metals. We know that metals conduct because they have a sea of delocalised electrons and that an insulator such as diamond does not conduct because all its electrons are locked into covalent bonds.

Silicon crystals are the most common semiconducting materials used in microelectronics and photovoltaics.
Jurii, CC BY 3.0

In silicon, it is possible to excite electrons into vacant higher energy orbitals. When this happens, these electrons become free to move. But at room temperature there is not sufficient energy to promote more than a small fraction of electrons and therefore the electrical conductivity stays exceedingly low.

To improve the electrical conductivity of silicon is to provide extra electrons that are easier to excite. This improvement is made by adding trace amounts of an element such as arsenic, a process called doping. The added element is called a dopant. Doping is a powerful way to increase the electrical conductivity of certain materials in the solid state

Boron is also used to dope silicon, but the mechanism is not the same as for arsenic. With boron, the number of electrons is reduced, which causes the formation of electron vacancies, known as holes. This encourages electrons to move to fill the holes, and thereby generates new holes. These holes are filed by other electrons moving in, and so on. There is thus a flow of charge through the silicon and so it conducts electricity.

The solar cell

The Earth receives more energy in sunlight in two days than is stored in all the known energy resources. So, if just a fraction of this energy could be harnessed it would provide a major alternative supply of electricity. Many solar energy plants already exist around the world, but many more are needed if they are to become a real alternative to the other energy sources.

Solar cells are used to convert radiant energy into electrical energy. One form of solar cell consists of two joined layers of silicon. One layer is doped with arsenic or phosphorus and is called n-type silicon (arsenic and phosphorus each have five electrons in their outer shells). This layer has extra electrons available for conduction, generated by the small number of dopant atoms. The other layer is doped with boron and is called p-type silicon. This layer has a shortage of electrons (a surplus of holes) because boron has only three outer electrons. The cell is connected into an external circuit.

Electrons flow from the n-type to the p-type layer. Eventually, equilibrium is established with a potential diference between the two layers. When sunlight falls on the cell, the equilibrium is disturbed and electrons move from the p-type to the n-type layer. The electrons return to the p-type layer via the external circuit, and thereby generate an electric curent.

From a solar cell to a PV system. Diagram of the possible components of a photovoltaic system
Rfassbind, Public Domain

SIMPLE-MOLECULE LATTICES

Some non-metals, such as solid nitrogen, oxygen, chlorine and iodine, form crystal lattices that consist of the simple molecules (N2, O2, Cl2, I2) held in place by attractive forces between the molecules. These forces are known as intermolecular forces because they are forces between molecules rather than within a molecule. The intermolecular forces are weak between simple molecules, and therefore the non-metals mentioned have relatively low melting points and boiling points. (The intermolecular forces are often referred to as van der Waals forces.) These weak forces explain why many non-metals are gases at room temperature and atmospheric pressure. It is easy to separate the molecules, but difficult to atomise them, because the covalent bonds within each molecule (the intramolecular bonds) are so much stronger.

Elements with a simple molecular structure do not conduct electricity, since they do not have electrons or ions that are free to move. All the electrons are localised.

DIATOMIC MOLECULES

As mentioned before, many gases, including chlorine, oxygen, nitrogen and hydrogen, form diatomic molecules. They cannot have a permanent dipole as both atoms must have equal electronegativity. So it is difficult to imagine forces of attraction between these neutral molecules.

An induced dipole-induced dipole attraction is set up between molecules. A temporary dipole is set up in a molecule when the electrons in the covalent bond that links the atoms move more to one end of the molecule than to the other. It is important to realise that the dipole is temporary and the electrons may also move more to the other end of the molecule. As a result of the temporary dipole, one end of the molecule carries a very small positive charge and the other end a very small negative charge. This small charge separation is not permanent and so in another instant in time the charges can switch ends.

The figure below shows the way an intermolecular force may be set up in a hydrogen molecule. Once a temporary dipole is set up within a molecule, it induces a similar dipole in neighbouring molecules and a very weak intermolecular force is set up.

The intermolecular force is known as an induced dipole-induced dipole interaction. When the temperature is sufficiently low, this force may be strong enough to hold the hydrogen molecules in a solid crystalline lattice.

The melting point of hydrogen is very low, since the two electrons in a hydrogen molecule are firmly attracted to the nuclei of their atoms. The larger the molecule, the stronger the intermolecular force, because there are more electrons and the outer electrons are under a weaker attractive force from the nuclei. So it is easier to produce a temporary dipole and the induced dipole-induced dipole attraction becomes stronger. In Group 7 this is illustrated by the change in physical state at room temperature and pressure. Fluorine, F2, a small molecule with few electrons, has weak induced dipole-induced dipole attractions, whereas bromine, Br2, a liquid, has more electrons, a larger molecule and stronger induced dipole-induced dipole attractions. In the case of iodine, I2, the strength of the intermolecular force is sufficient to make iodine a solid at room temperature.

It is important to understand that in these types of simple molecular crystal lattice it is attraction between molecules that holds the lattice together. The atoms which the molecule are covalently bonded to one another. But the strength of this covalent bonding does not determine the magnitude of the melting point.

Hydrogen-bonding-in-water-2D
Benjah-bmm27, Public Domain

PHOSPHORUS AND SULFUR

Phosphorus and sulfur both form simple molecular lattices, but a discussion of the structures is complicated since both elements exhibit allotropy. Suffice to say that phosphorus forms a lattice in which P4, molecules are arranged in a fixed pattern, while sulfur has a molecule with the formula S8. The P4 and S8, molecules are held in position in their respective lattices by weak induced dipole-induced dipole interactions.

Phosphorus and sulfur are both solids at room temperature and atmospheric pressure because the induced dipole-induced dipole interactions are much stronger than those in the diatomic molecules of nitrogen and oxygen, since there are more electrons in molecules of phosphorus and sulfur. Sulfur has a higher melting point than phosphorus because it has larger molecules with more electrons, so the intermolecular force is stonger.

THE NOBLE GASES

The noble gases exist as monatomic molecules; that is, just as atoms. All the noble gases have low melting points. Helium has the lowest melting point of any element, and therefore it must have the weakest intermolecular forces between its atoms.

The electrons in an atom are in constant motion and do not occupy set positions or set orbits. Therefore, it is possible for both of the electrons in a helium atom to arrive simultaneously on the same side of the atom. This side of the helium atom thus becomes very slightly negatively charged, while the other side becomes very slightly positively charged.

The electron cloud around the helium nucleus becomes distorted (it is no longer spherical), and it causes a similar distortion in neighbouring helium atoms. In other words, a temporary or instantaneous dipole is induced in neighbouring atoms.

A very weak attraction then exists between the slightly positive side of one helium atom and the slightly negative side of a neighbouring helium atom. This is a very weak intermolecular force because it relies on an asymmetric distribution of electrons, which has a low probability since both electrons are close to the nucleus, so are firmly attracted to the nucleus. At the same time, they repel each other.

An endohedral fullerene compound containing a noble gas atom
Hajv01, CC BY-SA 4.0

Once this force between helium atoms in the solid phase has been broken, the attraction between helium atoms in the liquid phase is negligible, and so helium boils only 4 degrees higher than the temperature at which it melts.

The larger the noble gas atom (hence the greater the number of electrons), the greater the likelihood of this asymmetric distribution of electrons. This is because the outer electrons are further away from the nucleus and are shielded from the nuclear charge. This means that the outer electrons can occupy a larger region and so undergo a smaller electron – electron repulsion, even when asymmetrically distributed. Therefore, the melting and boiling points of the noble gases increase with increasing atomic (proton) number.

A SUMMARY OF THE STRUCTURE AND BONDING OF ELEMENTS IN PERIOD 3

The physical properties of elements in Periods 2 and 3 vary considerably.

The two tables below summarize their physical properties.

Table 1: Summary of the physical properties of Period 2 elements

LithiumBerylliumBoronCarbonNitrogenOxygenFluorineNeon
BondingMetallicMetallicCovalentCovalentCovalentCovalentCovalent-
StructureGiantGiantGiantGiantSimpleSimpleSimpleSimple
Melting pointlowhighVery highVery highVery lowVery lowVery lowVery low
Electrical conductivity of solidHighHighlowlowlowlowlowlow

Table 2: Summary of the physical properties of Period 3 elements

SodiumMagnesiumAluminiumSiliconPhosphorusSulfurChlorineArgon
BondingMetallicMetallicMetallicCovalentCovalentCovalentCovalent-
StructureGiantGiantGiantGiantSimpleSimpleSimpleSimple
Melting pointlowhighhighVery highlowlowVery lowVery low
Electrical conductivity of solidHighHighHighHighVery lowVery lowVery lowVery low

SUMMARY

After reading this post, you should know the following:

● A crystalline solid has a regular arrangement of particles called a lattice that can be produced by the repetition in three dimensions of a unit cell. Crystalline solids have definite melting points.
● An amorphous solid has a closely packed structure that is disordered, and therefore does not have a unit cell. Amorphous solids melt over a range of temperatures, and many soften on heating.
● The structure and bonding within a crystal lattice determine the physical properties of an element.
● The melting point of a solid is determined by the strength of the force of attraction between particles in a lattice.
● A metal consists of closely packed positive ions in a sea of delocalised electrons. It has a giant structure.
● Metals conduct electricity because of the movement of the delocalised electrons in the presence of a potential difference.
● Metals are typically hard, strong, shiny and good thermal conductors. Many have high melting points. These properties can be explained by the structure and bonding.
● Non-metals form molecular lattices.
● Non metals that form a molecular lattice have a regular arrangement of simple molecules held together by weak intermolecular forces, called van der Waals forces (induced dipole-induced dipole interactions). Such non-metals have low melting points and low boiling points and do not conduct electricity.
● The strength of the induced dipole-induced dipole interaction in a diatomic element depends on the molecular size (indicated by the number of electrons in the molecule).
● A non-metal that forms a giant molecular or macromolecular lattice has a regular arrangement of atoms held together by strong covalent bonds. The whole crystal is assumed to be the molecule. Such non-metals have high melting points.
● Non-metals are generally poor electrical conductors because they have no free electrons.
● Different forms of the same element in the same state are known as allotropes. Allotropes have different physical properties, but similar chemical properties.
● Diamond, graphite and fullerene (C60) are three allotropes of carbon. Carbon and diamond have a giant molecular structure. Graphite has mobile delocalised electrons, so it conducts electricity. Diamond does not conduct electricity.

Thanks for coming.

REFERENCES

https://en.wikipedia.org/wiki/Semiconductor
https://www.britannica.com/science/semiconductor
https://depts.washington.edu/matseed/mse_resources/Webpage/semiconductor/semiconductor.htm
https://www.britannica.com/technology/solar-cell
https://en.wikipedia.org/wiki/Solar_cell
https://www.bbc.co.uk/bitesize/guides/zqrxdxs/revision/4
https://sites.google.com/site/ellesmerealevelchemistry/module-2-foundations-in-chemistry/2-2-electrons-bonding-and-structure/2-2-2-bonding-and-structure/2-2-2-n-simple-molecular-lattices
https://alevelnotes.com/notes/chemistry/elements-of-life/bonding
https://www.britannica.com/science/diatomic-molecule
https://www.thoughtco.com/what-are-the-seven-diatomic-elements-606623
https://en.wikipedia.org/wiki/Diatomic_molecule
https://en.wikipedia.org/wiki/Sulfur
https://en.wikipedia.org/wiki/Phosphorus
https://www.tandfonline.com/loi/gpss20
https://courses.lumenlearning.com/introchem/chapter/the-noble-gases-group-18/
https://www.britannica.com/science/noble-gas
https://en.wikipedia.org/wiki/Noble_gas
https://en.wikipedia.org/wiki/Period_2_element
http://www.docbrown.info/page07/ASA2ptable4b.htm
https://www.chemguide.co.uk/inorganic/period3/elementsphys.html
https://chem.libretexts.org/Bookshelves/Inorganic_Chemistry/Modules_and_Websites_(Inorganic_Chemistry)/Descriptive_Chemistry/Elements_Organized_by_Period/Period_3_Elements/Physical_Properties_of_Period_3_Elements



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