From the bonds module you know that atoms with 4 valence electrons prefer to share rather than lose or gain. Carbon is the most important such atom in the universe, and this module is about why.
Carbon's four bonds
Established Carbon has 4 valence electrons. This puts it right in the middle — losing 4 electrons or gaining 4 electrons are both energetically expensive. So carbon shares, forming four covalent bonds, and forming them strongly.
Four bonds in 3D space arrange themselves into a tetrahedron, with 109.5° between each pair, to minimise repulsion between the shared electron pairs. This tetrahedral geometry is at the heart of why organic molecules have the complex 3D shapes they do — and from the previous modules you know that shape determines function in biology. Carbon's geometry is the substrate for biological function.
Catenation — carbon's superpower
Established Carbon can bond strongly to itself, forming long chains, branches, and rings. This is called catenation. Almost no other element does this nearly as well.
Catenation lets carbon build long chains (fatty acids, polymers); branches (most amino acid side chains); rings (benzene, sugars, DNA bases, steroids); and combinations of all three (proteins, DNA). The number of possible carbon molecules is essentially unbounded. Tens of millions of distinct organic compounds are known, with billions more theoretically constructible.
Carbon's chemical palette
Carbon forms strong bonds with the most biologically relevant elements: C-H for chains and branches, C-O and C=O for sugars and acids, C-N for proteins and DNA, C-S for some amino acids and enzymes, C-P (via phosphate groups) for the DNA backbone.
This means carbon-based molecules can have a huge range of properties: greasy or water-loving, acidic or basic, reactive or inert, rigid or flexible. This is the chemical palette that biology paints with — and it's what gives biology the expressive range it has.
Why silicon isn't carbon
Silicon is right below carbon in the periodic table and has the same 4 valence electrons. So why isn't silicon-based life common in the universe, as far as we know?
Frontier Four real reasons:
First, Si-Si bonds are weaker than C-C bonds. Silicon chains fall apart more easily. They cannot sustain the long polymers that biology depends on.
Second, Si-O bonds are extremely strong. Too strong. Silicon reacted with oxygen forms SiO₂ — quartz, sand — which is a rock, not a versatile biomolecule. Carbon with oxygen forms CO₂, a gas that's easy to use and release. Silicon-based metabolism has nowhere comparable to release its "exhaust."
Third, silicon molecules don't dissolve well in water. They don't engage in the dynamic biochemistry that water enables. Hydrogen bonding, polarity, and hydrophobic exclusion (M-Chem-04) are all keyed to small polar carbon-and-water systems.
Fourth, silicon doesn't form stable double or triple bonds the way carbon does. This narrows the chemical vocabulary substantially.
This doesn't completely rule out exotic silicon biology in exotic environments (very cold, non-aqueous solvents like liquid methane, which Saturn's moon Titan has). But it explains why carbon is overwhelmingly favoured wherever we look in our solar system.
Bringing it back to the protomolecule
Now you can see why the protomolecule, whatever it is, almost certainly operates on carbon chemistry. It interacts with terrestrial biology, hijacks cellular machinery, and reorganises biomolecules. To do that, it has to speak the same chemical language — carbon-based, water-soluble, hydrogen-bond-rich, protein-and-nucleic-acid compatible.
The handwave-heavy advanced part isn't what the protomolecule is made of. It's how its instructions are encoded and executed with such precision. That's a topic for later modules.