Introduction to Organic Chemistry

Pick up a bottle of paracetamol from your hospital pharmacy. Every white tablet is an organic compound — yet for most of recorded human history, no chemist alive could have manufactured it in a laboratory. The prevailing belief, held with complete certainty until 1828, was that organic substances required something living to produce them: a mysterious animating force that no flask, furnace or mineral acid could ever replicate. This belief was called the Vital Force Theory, and it was so deeply accepted that most early chemists did not even try to synthesise organic compounds from scratch.

The theory collapsed in an unremarkable Berlin laboratory when Friedrich Wöhler, attempting nothing more than to prepare ammonium cyanate, discovered that his flask contained something else entirely — urea, the unmistakably organic compound that mammals excrete in urine. An inorganic salt had rearranged itself, without any cells or living tissue, into an organic product. The boundary between the living and the non-living dissolved in an afternoon, and modern organic chemistry was born.

In 1828, the young German chemist Friedrich Wöhler (1800–1882) was heating ammonium cyanate (an inorganic salt) when it rearranged to give urea — an unmistakably organic compound, identical to the urea isolated from mammalian urine.

This single experiment disproved the Vital Force Theory: an organic compound could be made from an inorganic precursor, with no living tissue involved. Modern organic chemistry begins from this date.

Why does gasoline dissolve engine grease but water does not? Why does alcohol catch fire but saltwater is used to douse flames? Why does sugar melt at 186 °C while table salt requires 801 °C to liquefy? All of these everyday contrasts trace back to one fundamental difference: organic compounds are held together by covalent bonds (shared electrons), while many inorganic compounds are held together by ionic bonds (transferred electrons). This single distinction ripples out into every physical and chemical property listed in the table below.

Think of the difference this way: ionic compounds like NaCl are fully charged — positive Na⁺ and negative Cl⁻ ions are locked in a rigid crystal lattice by strong electrostatic forces. That lattice takes enormous heat to break, so m.p. is high. When water surrounds the ions and pulls them apart, the salt dissolves instantly. Organic compounds are electrically neutral, covalent molecules. There is no lattice, so less energy breaks them — low m.p. And because water molecules cannot grip them electrostatically, most organic compounds are insoluble in water but dissolve freely in other covalent organic solvents. This is the chemist’s rule: like dissolves like.

In ionic bonding one atom outright takes an electron from another. Covalent bonding is more cooperative: two atoms share a pair of electrons, both nuclei attracted to the same electron cloud. This sharing holds the atoms together. In organic chemistry, almost every bond you will ever encounter is covalent — the C–H bonds in methane, the C–C backbone of an amino acid, the C=O of a ketone. Understanding how electrons are shared, and how evenly or unevenly, is the foundation on which all organic reactivity is built.

Here is the puzzle that puzzled Pauling. Carbon’s ground-state electron configuration is 1s² 2s² 2p² — only two unpaired electrons sit in the 2p subshell. By simple orbital theory, carbon should form two bonds. But methane is unambiguously CH₄ with four equivalent C–H bonds, all of identical length (110 pm) and angle (109.5°). Worse, if carbon simply used one 2s and three 2p orbitals, those four bonds should be different — one shaped like an s-orbital, three shaped like p-orbitals, pointing in strange directions. No such difference is observed. Pauling’s solution in 1931 was orbital hybridization.

The idea is elegant: mathematically blend the s and p atomic orbitals before bonding occurs, producing a new set of equivalent hybrid orbitals that point in the directions geometry demands. Think of it like making a smoothie. You start with one banana (the s orbital) and three apples (the three p orbitals), blend them completely, and get four identical banana-apple smoothies (four sp³ hybrids). Each smoothie is the same — because the ingredients were blended equally — and they point apart as far as possible (tetrahedral, 109.5°) to minimise repulsion between the electron pairs.

Organic molecules are three-dimensional objects, but chemistry communicates on paper in two dimensions. Over the past two centuries, chemists have developed four increasingly compact ways to draw the same molecule. You will encounter all four in this course and need to convert between them fluently, because an exam might give you a skeletal formula and ask you to name it, or give you a name and ask you to draw the Lewis structure. The key insight is that each representation just hides more implicit information — the Lewis structure shows everything; the skeletal formula hides every carbon atom and every C–H hydrogen, relying on you to reconstruct them mentally.

The skeletal (bond-line) formula is particularly important in medicine because drug structures in pharmacology textbooks and in the BNF are always drawn this way. A zig-zag line means a chain of carbons; each bend or endpoint is a carbon atom. Heteroatoms (O, N, Cl, etc.) are written in explicitly. Every carbon in the picture carries enough implicit H atoms to bring its bond count to exactly 4. Once you internalise that rule, you can “read” any drug structure you encounter in your clinical years.

Four ways to draw the same molecule — you must be fluent in moving between them, because the exam will give you any one and demand any other.

There are more than 10 million known organic compounds, a number that grows every year as synthetic chemists build new drugs and materials. At first glance this seems impossibly complex to memorise. The saving grace is that these 10 million molecules are all built from a small number of functional groups — specific atoms or clusters of atoms that react in predictable, characteristic ways regardless of what carbon scaffold they sit on. Learn the ~13 functional groups in the table below, and you know the chemical personality of every member of each family. The carboxylic acid group (–COOH) in aspirin behaves like the –COOH group in penicillin; the amine (–NH₂) in dopamine behaves like the amine in amphetamine.

For a medical student, functional groups are not just a chemistry exercise — they are the vocabulary of pharmacology. When you see a drug structure in the formulary, the functional groups immediately tell you: is it acidic or basic in plasma? Will it ionise? Does it H-bond? How will it be metabolised? Aspirin’s –COOH is deprotonated at pH 7.4 (so it stays in plasma); morphine’s –NH is protonated at stomach pH (so it is absorbed in the small intestine, not the stomach). Every drug interaction begins with a functional group.

Same molecular formula, completely different compound. Isomerism is one of the most clinically important concepts in organic chemistry, because it means that two molecules built from exactly the same atoms can have radically different shapes, solubilities, and biological activities. The antibiotic ampicillin and a related compound with the same formula can have vastly different bacterial kill rates; the two enantiomers of thalidomide include one that treats morning sickness and one that caused thousands of birth defects. In Chapter 1 we begin with the simplest type: structural (constitutional) isomers — the atoms are connected in a different order, even though the overall count is the same.

Think of it with Lego bricks. Give someone 4 red and 10 white bricks and ask them to build something. Two people can build quite different structures with the same pieces — a long straight tower vs a branched one. Those are chain isomers. Or they could build the same tower shape, but put the red bricks at position 1 vs position 2 — those are positional isomers. Or they could use the same pieces to build a car (an alcohol) vs a boat (an ether) that look nothing alike functionally — functional-group isomers. In all cases, the brick count is identical.

In Chapter 1 we meet only structural (constitutional) isomers — atoms connected in a different order. Stereoisomers (cis/trans, R/S, enantiomers) are introduced in Chapters 3 & 8.

A medical student once asked: “Why do I need organic chemistry? I’m going to be a doctor, not a chemist.” The answer is that the human body is an organic chemistry laboratory running 24 hours a day. Every breath, every heartbeat, every thought is the product of thousands of simultaneous organic reactions — oxidations, reductions, condensations, hydrolyses, phosphorylations — all happening in aqueous solution at 37 °C without any glass flasks. When any one of those reactions goes wrong, disease follows. Understanding the chemistry is the same as understanding the disease at its molecular root, which is increasingly where modern medicine operates.

• The body is ~70% water, but its dry mass is overwhelmingly organic. Proteins, carbohydrates, triglycerides and nucleic acids are all organic polymers built from a few dozen monomer types — amino acids, monosaccharides, fatty acids, nucleotides.
• Every metabolic step is an organic reaction. Glycolysis, the Krebs cycle, fatty-acid oxidation, the synthesis of neurotransmitters — all are oxidations, reductions, condensations, hydrolyses, substitutions you will meet in this course.
• Most disease processes are organic-chemistry failures. Diabetes — impaired glucose metabolism (Chapter 9). Cancer — mutated proteins (Ch 8) and dysregulated organic signalling. Sickle-cell anaemia — a single amino-acid substitution (Glu→Val at position 6 of β-globin).
• Almost every drug is an organic molecule. Aspirin, ibrutinib, antibiotics, anaesthetics, anti-virals — you cannot understand pharmacology without organic chemistry.
• Modern medicine is moving to the molecular level. CRISPR gene editing, mRNA vaccines, monoclonal antibodies, targeted small-molecule oncology — all rest on organic and biochemistry.

All items below are taken or close-paraphrased from TMU past papers (2019–2022). True/False statements (Part III) are the most common Chapter-1 format.