5.1 Properties of Water

The most obvious feature of the oceans is that they contain water. Water is so ubiquitous that it may not seem like a very interesting substance, but it has many unique properties that impact global oceanographic and climatological processes. Many of these processes are due to hydrogen bonds forming between water molecules.

The water molecule consists of two hydrogen atoms and one oxygen atom. The electrons responsible for the bonds between the atoms are not distributed equally throughout the molecule, so that the hydrogen ends of water molecules have a slight positive charge, and the oxygen end has a slight negative charge, making water a polar molecule. The negative oxygen side of the molecule forms an attraction to the positive hydrogen end of a neighboring molecule. This rather weak force of attraction is called a hydrogen bond (Figure 5.1.1). If not for hydrogen bonds, water would vaporize at -68 o C, meaning liquid water (and thus life) could not exist on Earth. These hydrogen bonds are responsible for some of water’s unique properties:

1. Water is the only substance to naturally exist in a solid, liquid, and gaseous form under the normal range of temperatures and pressures found on Earth. This is due to water’s relatively high freezing and vaporizing points (see below).

2. Water has a high heat capacity, which is the amount of heat that must be added to raise its temperature. Specific heat is the heat required to raise the temperature of 1 g of a substance by 1 o C. Water has the highest specific heat of any liquid except ammonia (Table 5.1.1).

Table 5.1.1 Specific heat values for a number of common substances

Water is therefore one of the most difficult liquids to heat or cool; it can absorb large amounts of heat without increasing its temperature. Remember that temperature reflects the average kinetic energy of the molecules within a substance; the more vigorous the motion, the higher the temperature. In water, the molecules are held together by hydrogen bonds, and these bonds must be overcome to allow the molecules to move freely. When heat is added to water the energy must first go to breaking the hydrogen bonds before the temperature can begin to rise. Therefore, much of the added heat is absorbed by breaking H bonds, not by increasing the temperature, giving water a high heat capacity.

Hydrogen bonds also give water a high latent heat; the heat required to undergo a phase change from solid to liquid, or liquid to gas. The latent heat of fusion is the heat required to go from solid to liquid; 80 cal/g in the case of ice melting to water. Ice is a solid because hydrogen bonds hold the water molecules into a solid crystal lattice (see below). As ice is heated, the temperature rises up to 0 o C. At that point, any additional heat goes to melting the ice by breaking the hydrogen bonds, not to increasing the temperature. So as long as ice is present, the water temperature will not increase. This is why your drink will remain cold as long as it contains ice; any heat absorbed goes to melting the ice, not to warming the drink.

When all of the ice is melted, additional heat will increase the temperature of the water 1 o C for each calorie of heat added, until it reaches 100 o C. At that point, any additional heat goes to overcoming the hydrogen bonds and turning the liquid water into water vapor, rather than increasing the water temperature. The heat required to evaporate liquid water into water vapor is the latent heat of vaporization which has a value of 540 cal/g (Figure 5.1.2).

The high heat capacity of water helps regulate global climate, as the oceans slowly absorb and release heat, preventing rapid swings in temperature (see section 8.1). It also means that aquatic organisms aren’t as subjected to the same rapid temperature changes as terrestrial organisms. A deep ocean organism may not experience more than a 0.5 o C change in temperature over its entire life, while a terrestrial species may encounter changes of more than 20 o C in a single day!

3. Water dissolves more substances than any other liquid; it is a “universal solvent”, which is why so many substances are dissolved in the ocean. Water is especially good at dissolving ionic salts; molecules made from oppositely charged ions such as NaCl (Na + and Cl – ). In water, the charged ions attract the polar water molecules. The ions get surrounded by a layer of water molecules, weakening the bond between the ions by up to 80 times. With the bonds weakened between ions, the substance dissolves (Figure 5.1.3).

4. The solid phase is less dense than the liquid phase. In other words, ice floats. Most substances are denser in the solid form than in the liquid form, as their molecules are more closely packed together as a solid. Water is an exception: the density of fresh water is 1.0 g/cm 3 , while the density of ice is 0.92 g/cm 3 , and once again, this is due to the action of hydrogen bonds.

As water temperature cools the molecules slow down, eventually slowing enough that hydrogen bonds can form and hold the water molecules in a crystal lattice. The molecules in the lattice are spaced farther apart than the molecules in liquid water, which makes ice less dense than liquid water (Figure 5.1.4). This is familiar to anyone who has ever left a full water bottle in the freezer, only to have it burst as the water freezes and expands.

But the relationship between temperature and water density is not a simple linear one. As water cools, its density increases as expected, as the water molecules slow down and get closer together. However, fresh water reaches its maximum density at a temperature of 4 o C, and as it cools beyond that point its density declines as the hydrogen bonds begin to form and the intermolecular spacing increases (Figure 5.1.5 inset). The density continues to decline until the temperature reaches 0 o C and ice crystals form, reducing the density dramatically (Figure 5.1.5).

There are a number of important implications to ice being less dense than water. Ice floating on the surface of the ocean helps regulate ocean temperatures, and therefore global climate, by influencing the amount of sunlight that is reflected rather than absorbed (see section 8.1). On a smaller scale, surface ice can prevent lakes and ponds from freezing solid during the winter. As fresh surface water cools, the water gets denser, and sinks to the bottom. The new surface water then cools and sinks, and the process is repeated in what is referred to as overturning, with denser water sinking and less dense water moving to the surface only to be cooled and sink itself. In this way, the entire body of water is cooled somewhat evenly. This process continues until the surface water cools below 4 o C. Below 4 o C, the water becomes less dense as it cools, so it no longer sinks. Instead, it remains as the surface, getting colder and less dense, until it freezes at 0 o C. Once fresh water freezes, the ice floats and insulates the rest of the water beneath it, reducing further cooling. The densest bottom water is still at 4 o C, so it does not freeze, allowing the bottom of a lake or pond to remain unfrozen (which is good news for the animals living there) no matter how cold it gets outside.

The dissolved salts in seawater inhibit the formation of the crystal lattice, and therefore make it harder for ice to form. So seawater has a freezing point of about -2 o C (depending on salinity), and freezes before a temperature of maximum density is reached. Thus seawater will continue to sink as it gets colder, until it finally freezes.

5. Water has a very high surface tension, the highest of any liquid except mercury (Table 5.1.2). Water molecules are attracted to each other by hydrogen bonds. For molecules not at the water surface, they are surrounded by other water molecules in all directions, so the attractive forces are evenly distributed in all directions. But for molecules at the surface there are few adjacent molecules above them, only below, so all of the attractive forces are directed inwards, away from the surface (Figure 5.1.6). This inwards force is what causes water droplets to take on a spherical shape, and water to bead up on a surface, as the spherical shape provides the minimum possible surface area. These attractive forces also cause the surface of the water to act like an elastic “skin” which allows things like insects to sit on the water’s surface without sinking.

Table 5.1.2 Surface tensions of various liquids

a weak bond between two molecules due to the electrostatic attraction of a proton in one molecule to the negative polar end of the other molecule (5.1)

a molecule where the electrons are not distributed equally, leading to a charge imbalance across the molecule; portions of the molecule are slightly positive while other portions are slightly negative (5.1)

the amount of heat needed to change a substance’s temperature by one degree (5.1)

the heat required to raise the temperature of 1 g of a substance by 1 degree C (5.1)

the energy that an object possesses due to its motion (5.1)

the change of state between a solid, liquid, or gas (8.1)

the heat required to change a substance from solid to liquid; 80 cal/g in the case of ice melting to water (5.1)

the heat required change a substance from liquid to gas; 540 cal/g to turn water into vapor (5.1)

an atom or molecule that has either gained or lost electrons and has thus become charged (5.1)

the vertical cycling within a body of water, where denser water sinks and less dense water floats to the surface (5.1)

where a cohesive layer forms on the water surface due to attraction between water molecules (5.1)

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Introduction to Oceanography Copyright © by Paul Webb is licensed under a Creative Commons Attribution 4.0 International License, except where otherwise noted.