Vitamin C (Ascorbic Acid): Chemistry and Function

Vitamin C is probably the most familiar vitamin of all: it’s in orange juice, it’s printed on cereal boxes, and it’s the one people reach for when they catch a cold. Chemically, it’s a small, water-soluble molecule with an unusual structure that makes it both an acid and an excellent reducing agent. Those two properties explain almost everything it does in the body, why it’s so easily destroyed in cooking, and how chemists measure it.

Identity

Property	Value
Chemical name	L-ascorbic acid
Formula	C₆H₈O₆
Molar mass	176.12 g mol⁻¹
Appearance	White crystalline solid
Melting point	about 190 °C (decomposes)
Solubility	Very soluble in water; poorly soluble in oils
First pKa	about 4.2
Food additive number	E300

You can confirm the molar mass from its formula with the molar mass calculator.

Structure

Ascorbic acid is closely related to the sugars, and in plants and most animals it’s made from glucose. Its structure has:

• a five-membered lactone ring (a cyclic ester) containing four carbons and one oxygen;
• a C=C double bond in the ring carrying two –OH groups (an arrangement called an enediol);
• a short side chain with two more –OH groups.

The molecule has two chiral centres; only the L form (L-ascorbic acid) has vitamin activity. Its many hydroxyl groups make it very polar, which is why it’s a water-soluble vitamin.

Why it’s an acid

Ascorbic acid has no carboxylic acid group, yet it’s roughly as acidic as acetic acid (the acid in vinegar). The acidic proton is on one of the enediol –OH groups. When it’s lost, the negative charge is spread by resonance over the C=C double bond and the neighbouring C=O group of the ring, stabilising the anion (ascorbate). This is the same kind of resonance stabilisation that makes carboxylic acids acidic (see acid strength and molecular structure).

At the pH of blood (about 7.4), well above its pKa of 4.2, vitamin C exists almost entirely as the ascorbate ion. Its acidity is also what gives citrus fruit part of its tartness, alongside citric acid.

Redox chemistry: an antioxidant

Vitamin C’s most important property is that it’s a good reducing agent: it readily gives away electrons.

Ascorbic acid loses two electrons and two protons (in two steps, via a relatively stable radical intermediate) to form dehydroascorbic acid:

C₆H₈O₆ → C₆H₆O₆ + 2H⁺ + 2e⁻

As a reducing agent, vitamin C:

• neutralises free radicals and reactive oxygen species, acting as an antioxidant in the watery parts of cells and blood;
• regenerates vitamin E, the antioxidant in cell membranes, after vitamin E has itself reacted with a radical;
• keeps metal ions reduced, which is the key to its enzyme roles below.

In the body, dehydroascorbic acid can be reduced back to ascorbate and reused. See oxidation and reduction for the general principles.

What vitamin C does in the body

1. Making collagen

Collagen, the most abundant protein in the body, forms a strong triple helix in skin, tendons, blood vessels, bones and gums. To be stable, collagen needs many of its proline and lysine amino acids converted to hydroxyproline and hydroxylysine. The enzymes that do this contain an iron ion that must be in the Fe²⁺ state. During the reaction, the iron sometimes gets stuck as Fe³⁺, and vitamin C reduces it back to Fe²⁺, keeping the enzyme working.

Without vitamin C, collagen is weak. That’s the cause of scurvy: bleeding gums, loosening teeth, easy bruising, poor wound healing and, eventually, death.

2. Absorbing iron

Iron in plant foods is mostly Fe³⁺, which is poorly absorbed. Vitamin C in the gut reduces Fe³⁺ to Fe²⁺, which is absorbed much more easily. Eating vitamin C-rich foods alongside plant sources of iron can significantly increase iron uptake (see iron in the body).

3. Other roles

Vitamin C is also a cofactor for enzymes that make carnitine (which carries fatty acids into mitochondria) and the hormone noradrenaline, and it supports immune function.

Why humans need it in the diet

Most mammals make their own vitamin C from glucose. Humans, other apes, guinea pigs and some bats and birds can’t, because the gene for the last enzyme in the pathway, L-gulonolactone oxidase, is broken. This mutation arose in our primate ancestors millions of years ago, probably causing no harm to fruit-eating animals with plenty of vitamin C in their diet.

Typical recommended intakes for adults range from about 40 mg a day (UK) to 75–90 mg a day (US). An orange provides roughly 50–70 mg. Good sources include citrus fruit, kiwi, strawberries, peppers, broccoli and potatoes.

A short history

• 1747: James Lind, a Scottish naval surgeon, gave different treatments to sailors with scurvy; those given oranges and lemons recovered quickly. It was one of the first controlled clinical trials.
• 1928–1932: the Hungarian biochemist Albert Szent-Györgyi isolated the substance (which he first called “hexuronic acid”) from adrenal glands and later paprika, and showed it prevented scurvy. He received the 1937 Nobel Prize in Physiology or Medicine.
• 1933: Norman Haworth worked out its structure and synthesised it (Nobel Prize in Chemistry 1937, shared). Its name, “a-scorbic”, means “against scurvy”.
• 1934 onwards: the Reichstein process allowed industrial production from glucose, combining chemical steps with a bacterial fermentation. Modern processes use more fermentation steps.

Vitamin C in food: why it disappears

Vitamin C is one of the least stable vitamins. It’s destroyed by:

• heat, especially prolonged boiling;
• oxygen, particularly when fruit or vegetables are cut, exposing more surface;
• alkaline conditions and traces of copper or iron ions, which catalyse its oxidation;
• leaching into cooking water, because it’s so water-soluble.

That’s why steamed or raw vegetables keep more vitamin C than boiled ones, and why orange juice loses vitamin C once opened. Food manufacturers add ascorbic acid (E300) as an antioxidant to stop cut fruit browning and to protect flavours.

Measuring vitamin C: titration

Because vitamin C is a reducing agent, it can be measured by redox titration (see titration calculations):

• DCPIP titration: the blue dye 2,6-dichlorophenolindophenol is reduced by vitamin C to a colourless form. Juice is added to a known volume of DCPIP until the blue colour disappears; the less juice needed, the more vitamin C it contains.
• Iodine titration: vitamin C reduces iodine (I₂) to iodide (I⁻). Iodine solution is added to the sample with starch indicator; the end point is the first permanent blue-black colour, when all the vitamin C has been used up (see the iodine test for starch).

C₆H₈O₆ + I₂ → C₆H₆O₆ + 2HI

Worked example: 25.0 cm³ of juice needed 12.5 cm³ of 0.00500 mol dm⁻³ iodine solution.

• Moles of I₂ = 0.00500 × 0.0125 = 6.25 × 10⁻⁵ mol
• Mole ratio 1:1, so moles of vitamin C = 6.25 × 10⁻⁵ mol
• Mass = 6.25 × 10⁻⁵ × 176.12 = 0.0110 g = 11.0 mg in 25.0 cm³, or about 44 mg per 100 cm³, a typical value for fresh orange juice.

Common misconceptions

• “Vitamin C prevents colds.” Large studies show regular supplements don’t reduce how often most people catch colds, though they may slightly shorten them.
• “Mega-doses are harmless.” Most excess is excreted, but very high doses can cause stomach upset and, in some people, kidney stones.
• “Natural vitamin C is better than synthetic.” The ascorbic acid molecule is identical whatever its source.

Key takeaways

• Vitamin C is L-ascorbic acid, C₆H₈O₆, a water-soluble lactone with an enediol group.
• It’s an acid (pKa ≈ 4.2) because the enediol anion is resonance-stabilised, and a strong reducing agent.
• It keeps the iron of collagen-making enzymes reduced; without it, collagen fails and scurvy results.
• It helps absorb iron by reducing Fe³⁺ to Fe²⁺ and acts as an antioxidant.
• It’s easily destroyed by heat, oxygen and leaching, and can be measured by DCPIP or iodine titration.