Chemical Bonds — How Atoms Stick Together

Atoms rarely exist in isolation. They combine with other atoms to form molecules and compounds — and the forces holding them together are called chemical bonds. The driving force is stability. Most atoms have an incomplete outer electron shell. By sharing or transferring electrons with other atoms, they can fill that shell and reach a more stable, lower-energy state. Think of it like puzzle pieces clicking together — each atom achieves something it couldn't alone. The type of bond formed depends on the atoms involved — specifically, how strongly they attract electrons. This property is called electronegativity. Oxygen and nitrogen pull electrons strongly; sodium and potassium barely hold onto their outer electrons at all.

A covalent bond forms when two atoms share a pair of electrons. Both atoms "claim" the shared electrons, giving each a fuller outer shell. Covalent bonds are the foundation of all biological molecules. When carbon (4 outer electrons) bonds with four hydrogens to make methane (CH₄), or forms the backbone of glucose, fatty acids, amino acids, and DNA — these are all covalent bonds. Types of covalent bonds: - Single bond (—): One pair of electrons shared. Strong but can rotate. Found in carbon chains of fatty acids. - Double bond (=): Two pairs shared. Stronger and rigid. The C=O bond in ketones and aldehydes; the O=O in oxygen gas. - Triple bond (≡): Three pairs shared. Very strong and rigid. Nitrogen gas (N≡N) — which is why N₂ is so unreactive despite making up 78% of air. Polar vs non-polar covalent bonds: When atoms share electrons equally, the bond is non-polar (e.g. C–H bonds). When one atom attracts electrons more strongly — as in O–H or N–H bonds — electrons spend more time near that atom, creating slight positive and negative charges. These are polar covalent bonds, and they're crucial for how water behaves. Medical connection: The difference between saturated and unsaturated fats is entirely about carbon double bonds. Saturated fats have only single C–C bonds (flexible → pack tightly → solid at room temperature → associated with cardiovascular risk). Unsaturated fats have at least one C=C double bond (creates a kink → cannot pack as tightly → liquid at room temperature → healthier profile).

An ionic bond forms when one atom gives an electron to another. The atom that loses an electron becomes positively charged (cation); the atom that gains one becomes negatively charged (anion). Opposite charges attract — that's the ionic bond. The classic example is sodium chloride (NaCl — table salt): - Sodium (Na) has 1 outer electron — it gives it away happily, becoming Na⁺ - Chlorine (Cl) has 7 outer electrons — it needs 1 more to fill its shell, becoming Cl⁻ - Na⁺ and Cl⁻ attract each other strongly Ionic bonds create salts — crystalline structures where positive and negative ions alternate in a lattice. In water, ionic compounds dissolve because water molecules pull the ions apart. The body's ionic story — electrolytes: The ions dissolved in your blood and cells are electrolytes, and they are vital: - Na⁺ (sodium): Controls fluid balance, nerve impulse generation - K⁺ (potassium): Inside cells; essential for heart rhythm and nerve conduction - Ca²⁺ (calcium): Triggers muscle contraction, nerve signals, blood clotting - Cl⁻ (chloride): Balances positive charges; important in stomach acid (HCl) - HCO₃⁻ (bicarbonate): The body's main blood pH buffer Abnormal electrolyte levels cause serious conditions: low potassium (hypokalaemia) → dangerous heart arrhythmias; low sodium (hyponatraemia) → brain swelling, confusion, seizures; low calcium (hypocalcaemia) → uncontrolled muscle twitching (tetany).

A hydrogen bond isn't a true chemical bond — it's an attraction between a slightly positive hydrogen atom (attached to O or N) and a slightly negative O or N on a nearby molecule. Each hydrogen bond is about 20 times weaker than a covalent bond. So why does it matter? Because hydrogen bonds act in huge numbers, and their combined effect is enormous: Water's unique properties (covered in the next lesson) — all due to hydrogen bonds between water molecules. DNA's double helix — the two strands are held together by hydrogen bonds between base pairs. A–T: 2 hydrogen bonds. G–C: 3 hydrogen bonds. The helix is stable enough to maintain the genetic code, but weak enough for the strands to be separated during replication and transcription. If DNA strands were joined by covalent bonds, they could never be read. Protein structure — hydrogen bonds between amino acids in the peptide backbone create alpha-helices and beta-sheets (secondary structure). They help fold proteins into the precise 3D shapes that determine their function. Drug binding — many drugs work by forming hydrogen bonds (and other weak interactions) with their target proteins. The antibiotic amoxicillin forms hydrogen bonds with bacterial enzymes; aspirin forms a covalent bond with COX enzymes. Drug designers carefully engineer molecules to maximise these interactions.

Two more weak forces complete the picture: Van der Waals forces are tiny, fleeting attractions between all molecules — caused by momentary imbalances in electron distribution creating brief dipoles. Each is incredibly weak, but when two large molecules have complementary shapes (fitting together like a lock and key), thousands of van der Waals contacts add up to a strong overall attraction. This is fundamental to enzyme-substrate binding, antibody-antigen recognition, and receptor-ligand interactions. The specificity of biological interactions often comes down to shape complementarity at the molecular level. Hydrophobic interactions aren't a bond at all — they're an entropy-driven effect. Non-polar molecules (like the tails of fatty acids, or the hydrophobic core of proteins) don't form hydrogen bonds with water. When pushed together, they let water molecules arrange more freely — which is energetically favoured. This is why: - Fats don't dissolve in water - Cell membranes are bilayers (hydrophobic tails in, hydrophilic heads facing water) - Proteins fold with hydrophobic amino acids buried inside, away from water Understanding these forces is the foundation of pharmacology — every drug-receptor interaction depends on a combination of hydrogen bonds, ionic interactions, van der Waals contacts, and hydrophobic effects.