2. Water has several important physical properties. Although these
properties are familiar because of the omnipresence of water, most of
the physical properties of water are quite atypical. Given the low molar
mass of its constituent molecules, water has unusually large values of
viscosity, surface tension, heat of vaporization, and entropy of
vaporization, all of which can be ascribed to the extensive hydrogen
bonding interactions present in liquid water. The open structure of ice
that allows for maximum hydrogen bonding explains why solid water
is less dense than liquid water—a highly unusual situation among
common substances.
3. ACID-BASE REACTIONS
Water undergoes various types of chemical reactions. One of the most important chemical
properties of water is its ability to behave as both an acid (a proton donor) and a base (a proton
acceptor), the characteristic property of amphoteric substances. This behaviour is most
clearly seen in the autoionization of water:H2O(l) + H2O(l) ⇌ H3O+(aq) + OH−(aq),where the (l)
represents the liquid state, the (aq) indicates that the species are dissolved in water, and the
double arrows indicate that the reaction can occur in either direction and an equilibrium
condition exists. At 25 °C (77 °F) the concentration of hydrated H+ (i.e., H3O+, known as the
hydronium ion) in water is 1.0 × 10−7 M, where M represents moles per litre. Since one OH− ion
is produced for each H3O+ ion, the concentration of OH− at 25 °C is also 1.0 × 10−7 M. In water at
25 °C the H3O+ concentration and the OH− concentration must always be 1.0 × 10−14:[H+][OH−] =
1.0 × 10−14,where [H+] represents the concentration of hydrated H+ ions in moles per litre and
[OH−] represents the concentration of OH− ions in moles per litre.
When an acid (a substance that can produce H+ ions) is dissolved in water, both the acid and
the water contribute H+ ions to the solution. This leads to a situation in which the
H+ concentration is greater than 1.0 × 10−7 M. Since it must always be true that [H+][OH−] = 1.0
× 10−14 at 25 °C, the [OH−] must be lowered to some value below 1.0 × 10−7. The mechanism for
reducing the concentration of OH− involves the reactionH+ + OH− → H2O,which occurs to the
extent needed to restore the product of [H+] and [OH−] to 1.0 × 10−14 M. Thus, when an acid is
added to water, the resulting solution contains more H+ than OH−; that is, [H+] > [OH−]. Such
a solution (in which [H+] > [OH−]) is said to be acidic.
The most common method for specifying the acidity of a solution is its pH, which is defined in
terms of the hydrogen ion concentration:pH = −log [H+],where the symbol log stands for a base-
10 logarithm. In pure water, in which [H+] = 1.0 × 10−7 M, the pH = 7.0. For an acidic solution, the
pH is less than 7. When a base (a substance that behaves as a proton acceptor) is dissolved in
water, the H+ concentration is decreased so that [OH−] > [H+]. A basic solution is characterized
by having a pH > 7. In summary, in aqueous solutions at 25 °C:
neutral solution
[H+] = [OH−]
pH = 7
acidic solution
[H+] > [OH−]
pH < 7
basic solution
[OH−] > [H+]
pH > 7
4. OXIDATION-REDUCTION REACTIONS
When an active metal such as sodium is placed in contact with
liquid water, a violent exothermic (heat-producing) reaction occurs
that releases flaming hydrogen gas.2Na(s) + 2H2O(l) → 2Na+(aq) +
2OH−(aq) + H2(g)This is an example of an oxidation-reduction
reaction, which is a reaction in which electrons are transferred
from one atom to another. In this case, electrons are transferred
from sodium atoms (forming Na+ ions) to water molecules to
produce hydrogen gas and OH− ions. The other alkali metals give
similar reactions with water. Less-active metals react slowly with
water. For example,iron reacts at a negligible rate with liquid
water but reacts much more rapidly with superheatedsteam to
form iron oxide and hydrogen gas.
Noble metals, such as gold and silver, do not react with water at
all.
