Water is the most familiar substance on Earth and one of the strangest. For such a small molecule it boils at a remarkably high temperature; hydrogen sulfide, a similar molecule, is a gas at room temperature. Its solid form floats on its liquid form, which is rare. It soaks up heat slowly and releases it slowly, which helps steady coastal climates. Most of these oddities trace back to the polar water molecule described in the bonds level.

The hydrogen bond

Each water molecule has a slightly negative oxygen end and slightly positive hydrogen ends. Neighbors line up so that opposite ends attract. A hydrogen bond is an attraction between a hydrogen atom on one molecule and a strongly electronegative atom, such as oxygen, on a neighboring molecule. It is much weaker than the bonds inside a molecule, but water has so many of them that together they hold the liquid tightly, raising its boiling point and giving it high surface tension.

When water freezes, hydrogen bonds lock each molecule into position with four neighbors. Ice is less dense than liquid water because hydrogen bonds hold its molecules in an open lattice with more space between them. So ice floats. Lakes freeze from the top down, and the ice layer insulates the water beneath, where fish survive the winter.

Water as a solvent

Water dissolves more substances than almost any other liquid, which is why it is called a nearly universal solvent. But it has limits, summed up in a rule chemists use every day: "like dissolves like." Polar and ionic substances dissolve in polar water, while nonpolar substances such as oils do not. Water molecules crowd around ions and polar molecules and pull them into solution. Oil molecules offer no charges to attract water, so water molecules cling to one another and squeeze the oil out into a separate layer.

Acids and bases

Lemon juice and vinegar taste sour. Soap and baking soda feel slippery and taste bitter. Chemists long sorted substances this way before they understood why. In 1884 the Swedish chemist Svante Arrhenius proposed an explanation based on ions. In Arrhenius's terms, acids release hydrogen ions in water, and bases release hydroxide ions. A hydrogen ion is simply a hydrogen atom without its electron, which is a lone proton.

In 1923 Johannes Brønsted in Denmark and Thomas Lowry in England, working separately, gave a broader definition: an acid is a substance that donates a proton, and a base is a substance that accepts one. This covers bases such as ammonia, which contain no hydroxide but readily grab protons.

Acids also differ in strength, which is not the same as how concentrated they are. A strong acid breaks apart almost completely in water, while a weak acid, such as the acetic acid in vinegar, does so only partly. A dilute strong acid can be more dangerous than a concentrated weak one.

The pH scale

In 1909 Søren Sørensen, a chemist at the Carlsberg brewery's laboratory in Copenhagen, introduced a simple way to describe acidity. The pH scale runs from about 0 to 14: 7 is neutral, lower numbers are acidic and higher numbers are basic. The scale is built so that each step of one unit means a tenfold change in acidity. A solution at pH 3 is ten times as acidic as one at pH 4 and a hundred times as acidic as one at pH 5. Lemon juice is around pH 2, coffee about 5, pure water 7 and household bleach about 12 or 13. Human blood is kept within a narrow band, from about 7.35 to 7.45.

Neutralizing and buffering

When an acid meets a base, they cancel each other's properties. Neutralization is the reaction of an acid with a base, producing water and a salt. Here "salt" means any ionic compound, not only table salt.

Keeping a narrow pH, as blood must, needs something more. A buffer is a solution that resists changes in pH by soaking up added acid or base. Blood uses dissolved carbon dioxide and bicarbonate as a buffer: if acid enters, bicarbonate takes up the extra protons; if base enters, the weak acid supplies protons to meet it. The oceans' own chemistry is shifting too, as they absorb carbon dioxide and grow more acidic.

Worked example: heartburn and the antacid tablet

Your stomach makes hydrochloric acid, a strong acid, and its contents usually sit somewhere around pH 1.5 to 3.5. That acidity helps break down food and kills many germs. The stomach lining protects itself with a layer of mucus. The esophagus, the tube above it, has no such protection, so when stomach contents splash upward, they burn. That is heartburn.

A common antacid tablet contains calcium carbonate, a mild base. In the stomach it reacts with the acid to make calcium chloride, a salt, along with water and carbon dioxide gas, which is why antacids can cause a burp. Baking soda works the same way. Both raise the pH enough to ease the burning.

Why not use a strong base such as lye, which is far more powerful? Because a strong base would attack the tissues as fiercely as the acid. The goal is not to drive the pH as high as possible but to nudge it to a safer level, and a weak base does that gently. You can see the same chemistry at home with red cabbage juice, a natural indicator that turns pink in vinegar and green or blue in baking soda solution.