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15.1: Properties of Water

Difficulty Level: At Grade Created by: CK-12
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Lesson Objectives

  • Describe the structure and polarity of a water molecule.
  • Describe the hydrogen bonding that occurs in water and ice.
  • Discuss the unique properties of water and ice.

Check Your Understanding

Recalling Prior Knowledge

  • What is the molecular geometry of the water molecule?
  • What is hydrogen bonding?

Water is one of the most abundant and critically important substances on Earth. In this lesson, we begin a thorough look at the physical properties of water.

Structure of Water

Water is everywhere on our planet. Oceans, rivers, and lakes cover about 75% of the Earth’s surface. Large quantities of water are frozen in the polar regions and in glaciers. Water is an essential component of all organisms: between 70% and 90% of the mass of living things is water. The chemical reactions of most biological processes take place in water-based solutions. The water cycle (Figure below) describes the continuous movement of water on, above, and below the surface of the Earth. Water changes between the solid, liquid, and vapor states at various points throughout the cycle. In order to better understand the importance of water, we begin by examining its structure.

The water cycle illustrates how water in all its forms constantly cycles through various systems of the Earth.

Structure of Water

As discussed in earlier chapters, water is a simple molecule consisting of one oxygen atom bonded to two different hydrogen atoms. Because of the higher electronegativity of the oxygen atom, the bonds are polar covalent. The oxygen atom attracts the shared electrons of the covalent bonds to a significantly greater extent than the hydrogen atoms. As a result, the oxygen atom acquires a partial negative charge (δ−), while the hydrogen atoms each acquire a partial positive charge (δ+). The molecule adopts a bent structure because of the two lone pairs of electrons on the oxygen atom. The H-O-H bond angle is about 105°, slightly smaller than the ideal 109.5° of an sp3 hybridized atomic orbital (see Figure below).

The water molecule, visualized in three different ways: ball-and-stick model, space-filling model, and structural formula with partial charges.

The bent shape of the water molecule is critical because the polar O-H bonds do not cancel one another, so the molecule as a whole is polar. Figure below illustrates the net polarity of the water molecule. The oxygen is the negative end of the molecule, while the area between the hydrogen atoms is the positive end of the molecule.

Water is a polar molecule, with a buildup of negative charge around the more electronegative oxygen atom.

Polar molecules attract one another by dipole-dipole forces; the positive end of one molecule is attracted to the negative end of a nearby molecule. In the case of water, the highly polar O-H bonds leave very little electron density around the hydrogen atoms. Each hydrogen atom is strongly attracted to the lone-pair electrons on the oxygen atom of an adjacent molecule. These are called hydrogen bonds, and they are stronger than conventional dipole-dipole forces (Figure below).

A hydrogen bond is the attraction between a lone pair of electrons on a highly electronegative atom in one molecule and an electron-deficient hydrogen atom in a nearby molecule.

Because each oxygen atom has two lone pairs, it can make hydrogen bonds to the hydrogen atoms of two other molecules. Figure below shows the result – an approximately tetrahedral geometry around each oxygen atom consisting of two covalent bonds and two hydrogen bonds.

As a result of two covalent bonds and two hydrogen bonds, the geometry around each oxygen atom is approximately tetrahedral.

Structure of Ice

Liquid water is a fluid. The hydrogen bonds in liquid water constantly break and reform as the water molecules tumble past one another. As water cools, its molecular motion slows and the molecules move gradually closer to one another. The density of any liquid increases as its temperature decreases. For most liquids, this continues as the liquid freezes, and the solid state is denser than the liquid state. However, water behaves differently. It actually reaches its highest density at about 4°C (Table below).

Densities of Water and Ice
Temperature (°C) Density (g/cm3)
100 (liquid) 0.9584
50 0.9881
25 0.9971
10 0.9997
4 1.000
0 (liquid) 0.9998
0 (solid) 0.9168

Between 4°C and 0°C, the density of water gradually decreases as the hydrogen bonds begin to form a network characterized by a generally hexagonal structure with open spaces in the middle of the hexagons (Figure below).

The structure of liquid water (left) consists of molecules connected by hydrogen bonds that are short-lived, since water is a fluid. In ice (right), the hydrogen bonds become rigid, resulting in an interconnected, hexagonally shaped molecular framework.

Ice is less dense than liquid water, so it floats. Ponds or lakes begin to freeze at the surface, closer to the cold air. A layer of ice forms, but it does not sink as it would if water did not freeze in this unique structure, which is dictated by its shape, polarity, and hydrogen bonding. If the ice were to sink as it froze, entire lakes would freeze solid. Since the ice does not sink, liquid water remains under the ice all winter long. This is important, as fish and other organisms are capable of surviving through winter (Figure below). Ice is one of only a very few solids that is less dense than its liquid form.

Ice fishing, a popular winter activity, would not be possible without the unusually low density of solid water, because ice would sink and lakes would freeze solid.

Properties of Water

Compared to other molecular compounds of relatively low molar mass, ice melts at a very high temperature. A great deal of energy is required to break apart the hydrogen-bonded network of ice and return it to the liquid state. Likewise, the boiling point of water is very high. Most molecular compounds of similar molar mass are gases at room temperature.

Surface Tension

Surface tension was discussed in the chapter States of Matter. Water has a high surface tension because of its strong intermolecular hydrogen bonds. Surface tension can be seen by the curved meniscus that forms when water is in a thin column, such as a graduated cylinder or a buret (Figure below).

The meniscus of water in a graduated cylinder (left) forms because of water’s hydrogen bonding. The diagram on the right shows the correct way to use a graduated cylinder to measure the volume of a liquid.

Vapor Pressure

The hydrogen bonding between liquid water molecules also explains water’s unusually low vapor pressure for a molecule of its size. Relatively few molecules of water are capable of escaping the surface of the liquid and entering the vapor phase. Because evaporation is slow, the resulting vapor exerts a low pressure in a closed container. Low vapor pressure is an important physical property of water that prevents lakes, oceans, and other large bodies of water from evaporating at a much faster rate.

Lesson Summary

  • Water is a molecular compound consisting of polar molecules that have a bent shape. The oxygen atom acquires a partial negative charge while the hydrogen atoms acquire partial positive charges.
  • Hydrogen bonding between water molecules is responsible for its high surface tension, low vapor pressure, and high melting and boiling points relative to other molecular substances with similar molar masses.
  • The rigid hydrogen bonds in ice form an open network that causes the solid form of water to be less dense than its liquid form. As a result, ice floats in liquid water.

Lesson Review Questions

Reviewing Concepts

  1. What is the shape of a water molecule?
  2. What physical property of an element determines whether it acquires a positive or negative partial charge when participating in a polar covalent bond?
  3. How many hydrogen bonds is each water molecule capable of making?
  4. Why is ice less dense than water?
  5. Discuss the environmental consequences if ice was denser than water.
  6. How does the vapor pressure of water at a given temperature compare to that of other substances with similar molar masses? Explain. How does this affect the rate of evaporation of bodies of water?


  1. Use Table above to answer the following.
    1. What is the volume of 100.0 g of water at 100°C?
    2. What is the volume of 100.0 g of water at 4°C?
    3. What is the volume of 100.0 g of water at 0°C?
    4. What is the volume of 100.0 g of ice at 0°C?
  2. A completely full bottle of boiling water is placed in the freezer. Discuss what happens as the water cools and eventually freezes.
  3. Explain which property of water is responsible for each of the following phenomena.
    1. An asphalt road breaks apart as water seeps into small cracks and freezes.
    2. A bottle of gasoline evaporates more quickly than a bottle of water.
    3. Water forms nearly spherical drops as it slowly drips out of a faucet.

Further Reading / Supplemental Links

  • Walter Wick, A Drop of Water: A Book of Science and Wonder. Scholastic Press, 1997.
  • The Hydrogen Bond, (http://www.wisc-online.com/Objects/ViewObject.aspx?ID=GCH1004)
  • Yves Marechal, The Hydrogen Bond and the Water Molecule: The Physics and Chemistry of Water, Aqueous and Bio Media. Elsevier Science, 2006.
  • John Gregory, Particles in Water: Properties and Processes. CRC Press, 2005.

Points to Consider

Water is often referred to as the universal solvent because it is capable of dissolving so many different substances.

  • How does the structure of water aid it in its ability to dissolve solids?
  • What is the nature of solutions made from ionic and molecular compounds?

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