The last time you enjoyed a bit of tart dressing on a salad, or perhaps just a splash of balsamic vinegar, you were probably too busy relishing the tingly gustatory experience to appreciate the underlying chemistry.

An impressive range of flavors and types of vinegar pepper the marketplace, and by the end of the 2010s, even "drinking vinegar" had made its way onto the shelves of health-food and grocery stores across the United States. But all of these have at least one thing in common: The ingredient that gives these dressings and sauces their distinctive "zing" is a molecule called acetic acid.

Acetic acid uses are not limited to the condiment world, although this is certainly the first aisle in the supermarket in which to look if you suddenly fund yourself in need of this compound. In terms of its acid-base chemistry, it is not an especially strong acid, so acetic acid hazards are more mundane than those of more corrosive acids such as sulfuric acid.

But before diving into acetic acid specifically (with a wetsuit!), you should be familiar with acid-base chemistry in general and how acids and bases can be used to manipulate each other, water and the pH (acidity or basicity) of solutions. Then, you'll get examples of how acetic acid is used and prepared and where it shows up in the world. When you're done, the last thing you should feel is a bitter taste in your mouth!

Various definitions of acids and bases have been proposed across the centuries, and for the most part, these are complementary by building on previous knowledge rather than supplanting it.

These compounds were identified as having unique properties many centuries ago (some acids notably have the ability to corrode metal), but not until the late 1800s was a formal definition proposed. At that time, Svante Arrhenius defined an acid as a substance that increased the hydrogen ion concentration in water.

When an acid is added to water, it dissociates into a proton and whatever is left (more on that in a moment). Because water does not exist solely as a sea of intact H₂O molecules, but rather as a combination of H₂O and some number of "free" H⁺ and OH^– ions.

This means that in effect it can serve as both an acid and a base. H₂O itself can act as a base by accepting a proton to become what is called a hydronium ion (H₃O⁺). You can see that adding a hydronium ion to a hydroxide ion yields the right raw material for 2 molecules of H₂O to form.

Other definitions of acids and bases help account for special cases that do not appear to make sense at a glance, like the fact that ammonia (NH₃) can serve as a base despite not being able to donate a hydroxyl group.

This is because acids can alternatively be seen as proton donors and bases as proton acceptors; better yet, acids can be treated as electron-pair acceptors and bases as electron-pair donors.

All of this talk of solutions assumes that readers know what these are. Regardless, it never hurts to review a fundamental concept of chemistry that is relevant to acetic acid and innumerable other compounds.

Most reactions you will read about or even try in the lab occur in an aqueous solution, which is a fancy name for a solid compound (a solute) dissolved in a water (more generally, a solution requires a liquid solvent, but it doesn't have to be water).

When certain solids, especially ionic compounds, are placed in solution, they dissolve readily, and often this is a consequence of the specific properties of the solute and the solvent. Water, for example, is a polar molecule, and also contains strong hydrogen bonds.

When table salt, or NaCl, is placed in water, its ionic bonds are no match for the electrochemical properties of water, and they come apart. The Na⁺ and Cl^– ions then find their way into spaces among the intact water molecules.

With acids and bases, the driving forces for dissolution are different, but the result is still the formation of ions. The hydronium ion (from the donated proton) represents the cation, while the anion is called the conjugate base. In nomenclature, this is where the suffix "ate" comes from: When acetic acid dissociates into its component ions, the conjugate base left in solution is called acetate.

Acetic acid is also known as ethanoic acid and less commonly as methane carboxylic acid. It has the chemical formula C₂H₄O₂, though it is usually written CH₃COOH to indicate that it is a carboxylic acid.

These are acids containing a carboxyl group, which is a terminal carbon atom double-bonded to oxygen as well as to a hydroxyl group. The H atom of the hydroxyl group is the acidic proton of the compound.

Acetic acid has a molecular weight of 60.05 grams per mole (g/mol). Acetic acid's density is 1.053 g/mol at room temperature in the liquid form, though it can also exist as a solid. The pKa of acetic acid is 4.76, which is the pH value at which half of the acid will be intact and the other half in the ionic form.

• The formula for the acetate ion (the conjugate base of acetic acid) is CH₃COO–. 

Acetic acid can be combined with sugars, spices and other foodstuffs to make various vinegars, but it is important outside the culinary world as well. Polymer compounds such as vinyl acetate are used in the manufacture of plastics, while cellulose acetate is used in the photographic industry.

Acetate is an important compound in biochemistry because it can be combined with a molecule called coenzyme A (CoA) to create acetyl-CoA, an important chemical in cellular respiration (in particular the Krebs cycle or citric acid cycle that occurs in the mitochondria).

Acetic acid is made in a number of ways: via the oxidation of acetaldehyde, by the oxidation of ethanol (ethyl alcohol) and by the oxidation of butane or butene. It can also be made on a large scale from the one-carbon alcohol methanol.

Acids are corrosive and can damage skin, eyes and other organic tissues. Do not treat the fact that vinegar is drinkable or that acetic acid is termed "weak" as an excuse to be careless. If only 1 part in 20 of most vinegars is acetic acid and the rest water, imagine what it would feel like at full strength.

Acids can damage more than just skin because some are volatile and readily evaporate; this means you could wind up inhaling chemicals that could irritate the lining of your nasal passages and throat.

As a general guideline, always wear eye and hand protection while working with acids and bases, no matter the molarity or the identity of the acid or base. In fact, not to end on a "sour" note, but you should always use safety precautions in chemistry labs, especially if you want to keep doing more of them!