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14.1: Water - Some Unique Properties

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    Describe the different properties of water as it relates to its polarity and ability to form hydrogen bonds.

    Looking Closer: Water, the Most Important Liquid

    Earth is the only known body in our solar system that has liquid water existing freely on its surface. That is a good thing because life on Earth would not be possible without the presence of liquid water. Water is a simple molecule consisting of one oxygen atom bonded to two different hydrogen atoms (Figure \(\PageIndex{1}\)). Because of the higher electronegativity of the oxygen atom, the bonds are polar covalent (polar bonds). 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 requires a partial negative charge \(\left( \delta - \right)\), while the hydrogen atoms each acquire a partial positive charge \(\left( \delta + \right)\). The molecule adopts a bent structure because of the two lone pairs of electrons on the oxygen atom. The \(\ce{H-O-H}\) bond angle is about \(105^\text{o}\), slightly smaller than the ideal \(109.5^\text{o}\) of an \(sp^3\) hybridized atomic orbital.The bent shape of the water molecule is critical because the polar \(\ce{O-H}\) bonds do not cancel one another and the molecule as a whole is polar.

    Figure \(\PageIndex{1}\) Polarity of the water molecule due to the uneven distribution of electrons in its covalent bond. C (From OpenStax Concepts of Biology text)

    Hydrogen Bonds

    Due to water’s polarity, each water molecule attracts other water molecules as oppositely charged ends of the molecules attract each other. When this happens, a weak interaction occurs between the positive hydrogen end from one molecule and the negative oxygen end of another molecule. This interaction is called a hydrogen bond. This hydrogen bonding contributes to the following water’s unique properties.

    1. Water is the universal solvent

    2. Exists in nature as a solid, liquid, and gas

    3. The density of ice is less than liquid water

    4. Water has a high heat capacity

    5. Water has a high heat of vaporization

    6. Water exists as a liquid at room temperature It is important to note here that even we are only focusing on water in this text book, hydrogen bonding also occurs in other substances that have polar molecules.

    Density of Water

    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 (Table \(\PageIndex{1}\)) . It actually reaches its highest density at about \(4^\text{o} \text{C}\).

    Temperature \(\left( ^\text{o} \text{C} \right)\) Density \(\left( \text{g/cm}^3 \right)\)
    Table \(\PageIndex{1}\) Density of Water and Ice
    100 (liquid) 0.9584
    50 0.9881
    25 0.9971
    10 0.9997
    4 1.0000
    0 (liquid) 0.9998
    0 (solid) 0.9168

    Between \(4^\text{o} \text{C}\) and \(0^\text{o} \text{C}\), the density 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 \(\PageIndex{2}\) ).

    Figure \(\PageIndex{2}\) Hydrogen bonding makes ice less dense than liquid water. The lattice structure water is more condensed (left structure) than that of ice (right structure). The lattice structure of ice makes it less dense than freely flowing molecules of liquid water, enabling ice to float on liquid water. ( Image credit: Lynn Yarris,

    Ice is less dense than liquid water and so it floats. Ponds or lakes begin to freeze at the surface, closer to the cold air. A layer of ice forms, but does not sink as it would if water did not have this unique structure 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. Ice is one of only a very few solids that is less dense than its liquid form.

    Solvation Ability of Water

    Water typically dissolves many ionic compounds and polar molecules. Nonpolar molecules such as those found in grease or oil do not dissolve in water. We will first examine the process that occurs when an ionic compound such as table salt (sodium chloride) dissolves in water.

    Water is attracted to the sodium chloride crystal because water is polar and has both a positive and a negative end. The positively charged sodium ions in the crystal attract the oxygen end of the water molecules because they are partially negative. The negatively charged chloride ions in the crystal attract the hydrogen end of the water molecules because they are partially positive. The action of the polar water molecules takes the crystal lattice apart (see Figure \(\PageIndex{3}\)) .

    Figure \(\PageIndex{3}\) The dissolving of sodium chloride in water.

    After coming apart from the crystal, the individual ions are then surrounded by solvent particles in a process called solvation. Note that the individual \(\ce{Na^+}\) ions are surrounded by water molecules with the oxygen atom oriented near the positive ion. Likewise, the chloride ions are surrounded by water molecules with the opposite orientation. Hydration is the process of solute particles being surrounded by water molecules arranged in a specific manner. Hydration helps to stabilize aqueous solutions by preventing the positive and negative ions from coming back together and forming a precipitate.

    Table sugar is sucrose \(\left( \ce{C_{12}H_{22}O_{11}} \right)\) and is an example of a molecular compound. Solid sugar consists of individual sugar molecules held together by intermolecular attractive forces. When water dissolves sugar, it separates the individual sugar molecules by disrupting the attractive forces, but does not break the covalent bonds between the carbon, hydrogen, and oxygen atoms. Dissolved sugar molecules are also hydrated, but without as distinct an orientation to the water molecules as in the case of the ions. The sugar molecules contain many \(\ce{-OH}\) groups that can form hydrogen bonds with the water molecules, helping form the sucrose solution.

    The Amphoteric Nature of Water

    Water is amphoteric: it has the ability to act as either an acid or a base in chemical reactions.[80] According to the Brønsted-Lowry definition, an acid is a proton (H+) donor and a base is a proton acceptor.[81] When reacting with a stronger acid, water acts as a base; when reacting with a stronger base, it acts as an acid.[81] For instance, water receives an H+ ion from HCl when hydrochloric acid is formed:

    HCl(acid) + H2O(base) ⇌ H3O+ + Cl

    In the reaction with ammonia, NH3, water donates a H+ ion, and is thus acting as an acid:

    NH3(base) + H2O(acid) ⇌ NH+4 + OH

    High Heat Capacity and Specific Heat of Water

    Different substances respond to heat in different ways. If a metal chair sits in the bright sun on a hot day, it may become quite hot to the touch. An equal mass of water in the same sun will not become nearly as hot. Table \(\PageIndex{2}\) list specific heats of various substances compared to water.Water has the highest specific heat capacity of any liquid. Water’s high heat capacity is a property caused by hydrogen bonding among the water molecules. Specific heat is defined as the amount of heat one gram of a substance must absorb or lose to change its temperature by one degree Celsius. For water, this amount is 1 cal/goC. The units for specific heat can either be in the SI units of joules per gram per degree \(\left( \text{J/g}^\text{o} \text{C} \right)\) or calories per gram per degree \(\left( \text{cal/g}^\text{o} \text{C} \right)\). This text will use \(\text{J/g}^\text{o} \text{C}\) for specific heat.

    Figure \(\PageIndex{4}\) This power plant in West Virginia, like many others, is located next to a large lake so that the water from the lake can be used as a coolant. Cool water from the lake is pumped into the plant, while warmer water is pumped out of the plant and back into the lake.

    It takes water a long time to heat up and a long time to cool down. In fact, the specific heat capacity of water is about five times more than that of sand. This explains why land cools faster than the sea. Coastal climates are much more moderate than inland climates because of the presence of the ocean. Water in lakes or oceans absorbs heat from the air on hot days and releases it back into the air on cool days. Water is used as a coolant for machinery because it is able to absorb large quantities of heat (Figure \(\PageIndex{4}\)). Due to its high heat capacity, warm-blooded animals use water to disperse heat more evenly and maintain temperature in their bodies: it acts in a similar manner to a car’s cooling system, transporting heat from warm places to cool places, causing the body to maintain a more even temperature.


    Specific Heat \(\left( \text{J/g}^\text{o} \text{C} \right)\)
    Table \(\PageIndex{2}\) Specific Heats of Some Common Substances
    Water (l) 4.18
    Water (s) 2.06
    Water (g) 1.87
    Ammonia (g) 2.09
    Ethanol (l) 2.44
    Aluminum (s) 0.897
    Carbon, graphite (s) 0.709
    Copper (s) 0.385
    Gold (s) 0.129
    Iron (s) 0.449
    Lead (s) 0.129
    Mercury (l) 0.140
    Silver (s) 0.233

    High Heat of Vaporization of Water

    Water in its liquid form has an unusually high boiling point temperature, a value close to 100°C. As a result of the network of hydrogen bonding present between water molecules, a high input of energy is required to transform one gram of liquid water into water vapor, an energy requirement called the heat of vaporization. Water has a heat of vaporization value of 40.65 kJ/mol. A considerable amount of heat energy (586 calories) is required to accomplish this change in water. This process occurs on the surface of water. As liquid water heats up, hydrogen bonding makes it difficult to separate the water molecules from each other, which is required for it to enter its gaseous phase (steam). As a result, water acts as a heat sink, or heat reservoir, and requires much more heat to boil than does a liquid such as ethanol (grain alcohol), whose hydrogen bonding with other ethanol molecules is weaker than water’s hydrogen bonding.

    The fact that hydrogen bonds need to be broken for water to evaporate means that a substantial amount of energy is used in the process. As the water evaporates, energy is taken up by the process, cooling the environment where the evaporation is taking place. In many living organisms, including humans, the evaporation of sweat, which is 90 percent water, allows the organism to cool so that homeostasis of body temperature can be maintained.


    • The polarity of water and its ability to hydrogen bond contributes to its unique properties.
    • Ionic solute molecules are hydrated (surrounded by solvent molecules in a specific orientation).
    • Ice is less dense than liquid water due to spaces in the intermolecular structure of ice not present in water.
    • Heat capacity is the amount of heat required to raise the temperature of an object by \(1^\text{o} \text{C}\)).
    • The specific heat of a substance is the amount of energy required to raise the temperature of 1 gram of the substance by \(1^\text{o} \text{C}\).
    • The dissociation of liquid water molecules, which changes the substance to a gas, requires a lot of energy.

    Contributors and Attributions

    • CK-12 Foundation by Sharon Bewick, Richard Parsons, Therese Forsythe, Shonna Robinson, and Jean Dupon.

    • Boundless: The Chemical Foundation of Life
    • Libretext: Introduction to Environmental Science (Zendher et al.)
    • TextMap: Introductory Chemistry (Tro)
    • Marisa Alviar-Agnew (Sacramento City College)

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