Narration script

Unit two, chemical bonding. Atoms bond by sharing or transferring their valence electrons. The type of bond, covalent, ionic or metallic, decides whether a material is a semiconductor, an insulator or a conductor; whether it is hard or soft, brittle or bendable; and how strongly it absorbs light. Silicon is a covalent network. Silver is metallic. Glass is a mixed covalent and ionic network. And halide perovskites are strongly ionic.

You will learn to distinguish the bond types, use electronegativity to predict bond character, explain silicon's tetrahedral bonding, connect bond strength to band gap and other properties, and explain how hydrogen passivates defects.

Why silicon? Why does hydrogen in silicon nitride improve cells? Why do silver contacts conduct while glass does not? All of these are bonding questions. And passivation, the main lever on a cell's voltage, is bond chemistry at surfaces.

In a covalent bond, two atoms share a pair of electrons. Silicon mixes its outer s and p orbitals into four equivalent hybrid orbitals pointing to the corners of a tetrahedron, about one hundred and nine and a half degrees apart. Each overlaps with a neighbour to form a shared pair. So every silicon atom bonds to exactly four neighbours, making a rigid, directional three-dimensional network. The consequences: a high melting point of fourteen hundred and fourteen degrees Celsius, hardness, brittleness, because bonds cannot slide, they break, which is why wafers crack rather than bend, and a band gap, because electrons are locked into bonds unless they get at least one point one two electron-volts.

Here are the three bond types side by side. Covalent: a shared electron pair, pointing in a definite direction, as in silicon. Ionic: an electron transferred from one atom to another, and the resulting ions attract, as in table salt. Metallic: positive ion cores sitting in a sea of free electrons, as in silver, copper and aluminium.

In ionic bonding, one atom pulls so much harder that an electron is effectively transferred. Sodium gives one to chlorine, and the two ions attract in all directions. Ionic solids are brittle, have high melting points, usually insulate, with large band gaps, and often dissolve in water.

In metals, the valence electrons are not tied to any atom. They form a sea around positive ion cores. Those free electrons make metals excellent conductors, and make them shiny. Because the bonding is not directional, planes of atoms can slide, so metals bend. Copper ribbon bends; silicon snaps. Silver has the lowest resistivity of any metal, about one point six micro-ohm centimetres, just ahead of copper. Its cost is why the industry works so hard to reduce silver per cell.

Real bonds lie on a spectrum, and the electronegativity difference between the two atoms tells you where.

This chart shows Pauling electronegativities for elements that matter in solar. Silicon bonded to silicon is purely covalent. Gallium-arsenic and cadmium-tellurium, with differences around zero point four, are polar covalent. Silicon-oxygen, with a difference of about one and a half, is strongly polar. Caesium-iodine is largely ionic.

The table maps these bonds to devices. Silicon-hydrogen bonds passivate surfaces. Silicon-nitrogen makes the anti-reflection coating. Silicon-oxygen makes glass and the native oxide. Perovskites are listed as ionic despite modest electronegativity differences, which helps explain their soft lattice and low-temperature processing.

Now the key link between chemistry and electronics. Stronger, shorter bonds push the bonding states further below the antibonding states, so the band gap grows.

This chart shows it down group four. Same diamond structure, but longer bonds as the atoms get bigger. Carbon, as diamond, has a gap of about five and a half electron-volts. Silicon, one point one two. Germanium, zero point six six. And grey tin, about zero. Adding ionic character at a similar bond length widens the gap: germanium, then gallium arsenide at one point four two, then zinc selenide at about two point seven. That is the logic of band-gap engineering in compound semiconductors.

Bond energies matter for manufacturing too. A silicon-oxygen bond is about four point seven electron-volts, much stronger than silicon-silicon at about two point three. That is why turning sand into silicon costs so much energy, and why any silicon surface exposed to air instantly grows a thin oxide.

At a surface or grain boundary, silicon atoms are missing neighbours. They have dangling bonds, half-filled orbitals with energies inside the band gap. These trap electrons and holes and make them recombine, which kills voltage. Passivation ties them off. Chemical passivation bonds hydrogen, oxygen, or a matched layer to the dangling bonds, removing their states from the gap. Field-effect passivation uses fixed charge in a coating, negative in aluminium oxide and positive in silicon nitride, to repel one carrier type from the surface, so recombination has no partners. And hydrogen-rich silicon nitride releases hydrogen during firing, which diffuses into the wafer and passivates defects in the bulk as well.

Finally, weak bonds: van der Waals forces and hydrogen bonds, ten to a hundred times weaker than covalent bonds. They govern polymers. They hold encapsulant chains together, set their softness, and control adhesion to glass, often helped by silane coupling agents that form true covalent links.

Misconceptions. Bonds are not simply ionic or covalent; it is a continuum. A stronger material does not conduct better; diamond is an insulator. Passivation is not a protective paint; it is electronic, working at the atomic scale. And metals do not automatically make good contacts to silicon; contacts need careful engineering.

For the job: passivation layers, deposited by atomic layer deposition or plasma-enhanced chemical vapour deposition, are the core of every modern cell design. Wafer brittleness drives breakage-yield engineering. And metallization choices balance conductivity, cost and contact chemistry.

Try the review questions. Next unit, we see how these bonds arrange atoms into crystals.