1Why a water molecule is polar

Oxygen attracts shared electrons more strongly than hydrogen does. In each O–H bond the electron pair spends more of its time near the oxygen. So the bond is covalent (electrons are shared) but unevenly shared: a polar covalent bond. The oxygen end gains a partial negative charge, written δ−, and each hydrogen a partial positive charge, δ+. These are partial charges, not the full charges of an ion; no electron has been handed over.

Water is polar because it is also bent: the two hydrogens sit on the same side of the oxygen, about 104.5° apart, so both δ+ charges are on one side and the δ− on the other. In an exam, say both halves: polar covalent bonds, and a bent shape.

Polar covalent bonds hold the atoms of one water molecule together; hydrogen bonds hold different water molecules to each other.

2Hydrogen bonds: between molecules, and within big ones

A hydrogen bond is the attraction between a δ+ hydrogen that is covalently bonded to oxygen (or nitrogen) and a δ− oxygen or nitrogen atom somewhere else. In water, each molecule can hydrogen-bond to as many as four neighbours: its two hydrogens point at two other oxygens, and its oxygen attracts the hydrogens of two more. Figure 1 shows that arrangement and labels the two kinds of bond. The solid lines are covalent bonds, inside one molecule. The dashed lines are hydrogen bonds, between molecules.

A single hydrogen bond is far weaker than a covalent bond and lasts only a split second in liquid water before re-forming with a new partner. Their strength is in their number.

The same attraction works within large biological molecules and between them. Hydrogen bonds hold the two strands of DNA together (topic 1.6) and fold parts of a protein chain into helices and sheets (topic 1.7). Water also forms hydrogen bonds with any other polar or charged molecule. That gives you two words used throughout the course:

• Hydrophilic ("water-loving"): polar or charged substances, such as glucose and salts. Water molecules cluster around them, so they dissolve.
• Hydrophobic ("water-fearing"): nonpolar substances, such as fats and oils. Water cannot hydrogen-bond to them, so it bonds to itself and leaves them clumped together, which is why membranes hold together (topic 1.5).

3High specific heat capacity: slow to warm, slow to cool

Specific heat capacity is the energy needed to raise the temperature of 1 gram of a substance by 1 °C. For water it is about 4.2 J per gram per °C, higher than almost any other common liquid (ethanol needs about 2.4 J).

Why so high? Temperature measures how fast molecules move. Much of the heat put into water goes into breaking hydrogen bonds, not speeding molecules up, so its temperature rises only a little. On cooling, hydrogen bonds form and release energy, which slows the fall.

Organisms are mostly water, so a burst of heat from exercise or a hot afternoon changes body temperature only slowly, giving the body time to correct it: water helps maintain a homeostatic body temperature. In the same way a lake warms slowly in summer and cools slowly in winter, so what lives in it faces small swings.

4High heat of vaporization: evaporative cooling

Heat of vaporization is the energy needed to turn a given mass of liquid into gas. For a water molecule to escape from the liquid, every hydrogen bond holding it must break. So evaporating water takes a lot of energy: about 2.4 kJ for each gram at body temperature.

That energy comes from the liquid and the surface it sits on, so what is left behind is cooler: evaporative cooling. Sweat on skin, saliva on a panting dog's tongue and water evaporating from leaves all carry heat away. Because the heat of vaporization is high, a little evaporated water removes a lot of heat.

A careful exam answer links three steps: evaporation needs energy to break hydrogen bonds, that energy is heat taken from the body or leaf, so the body or leaf cools.

5Reading cooling data: means and standard error

Experiments on evaporative cooling give you numbers, and AP expects you to judge whether a difference is real. Here is an invented experiment. A student measured the surface temperature of 6 leaves on plants in full sun with their stomata open (so water was evaporating from the leaves), and 6 leaves on plants given a treatment that closes the stomata.

The formula sheet gives the mean, the standard deviation and the standard error of the mean, SE = s ÷ √n, where s is the standard deviation and n the sample size.

The two ranges of mean ± 2SE do not overlap, so the difference of about 3.9 °C is very likely real, not chance. Now write the claim, evidence and reasoning: leaves losing water by evaporation were cooler (claim); mean 31.07 °C against 35.00 °C, with non-overlapping ±2SE ranges (evidence); evaporating water takes heat from the leaf to break hydrogen bonds, because water has a high heat of vaporization (reasoning).

6Cohesion, adhesion and surface tension

All three come from hydrogen bonding.

• Cohesion is water molecules sticking to each other through hydrogen bonds. It lets a thin, continuous column of water be pulled up through the narrow tubes of a plant's xylem from roots to leaves: as water evaporates from the leaf, the molecules behind are pulled up after it, and the column does not snap.
• Adhesion is water molecules sticking to other polar or charged surfaces, such as the cellulose walls of plant cells or the glass of a tube. Adhesion is why water creeps up a narrow glass tube or a paper towel, a movement called capillary action. In a narrow tube, adhesion pulls water up the walls and cohesion drags the rest of the column with it.
• Surface tension is how hard it is to break the surface of a liquid. Surface molecules have no water above them, so they hydrogen-bond to those beside and below, forming a taut film that a pond skater can stand on.

Keep the two words apart: cohesion is water to water; adhesion is water to something else.