The self-ionization of water

Water itself is ionized to a very small extent (equation 6.1) and the value of the self-ionization constant, (equation 6.2), shows that the equilibrium lies well to the left-hand side. The self-ionization in equation 6.1 is also called autoprotolysis. 

Although we use concentrations in equation 6.2, this is an approximation, and we return to this in Section 6.3. 

If a pure liquid partially dissociates into ions, it is self- 

A Bronsted acid can act as a proton donor, and a Brensted base can function as a proton acceptor. 

In aqueous solution, protons are solvated and so it is more correct to write [H30]+(aq) than H (aq). Even this is oversimplified because the oxonium ion is further hydrated and species such as [H502] (see Fig. 10.1), [H703] and [H904] are also present. 

Equilibrium 7.1 illustrates that water can function as both a Brpnsted acid and a Brpnsted base. In the presence of other Br0nsted acids or bases, the role of water depends on the relative strengths of the various species in solution. When HCl is bubbled into water, the gas dissolves and equilibrium 7.3 is established. 

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Reaction (5.N) describes the nylon salt nylon equilibrium. Reactions (5.0) and (5.P) show proton transfer with water between carboxyl and amine groups. Since proton transfer equilibria are involved, the self-ionization of water, reaction (5.Q), must also be included. Especially in the presence of acidic catalysts, reactions (5.R) and (5.S) are the equilibria of the acid-catalyzed intermediate described in general in reaction (5.G). The main point in including all of these equilibria is to indicate that the precise concentration of A and B... 

At the first stoichiometric point of the titration, aii the diprotic acid has been converted to its conjugate base, H A. This amphiprotic anion can react with itseif, analogous to the self-ionization of water ... 

That the self-ionization of water is slight and is consistent with the poor electrical conductance of the liquid... 

So, the pH of pure water is 7.00 because that is the negative logarithm of the hydrogen ion concentration that is generated by the self ionization of water. It is a value based on the inherent acid-base properties of water. 

Let s see, in pure water the hydrogen ion concentration is 1.0 x 10-7 M. Okay, you say. So let s just add the 1.0 x 10"8 M from the HC1 to the 1.0 x 10-7 M from the water. But that doesn t work, because the introduction of H+ from HC1 impacts the self-ionization of water. According to Le Chatelier s principle, the position of the equilibrium will be shifted to the left, because we are adding a product. Many biological systems are coupled equilibria, so if you change one, you change them all. If you want to solve one system, you have to solve all of them simultaneously, because they are all interconnected. 

What role does the self-ionization of water play in acid-base chemistry ... 

A sample of pure water will contain a small quantity of ions (H+ and OH-) produced from the self-ionization of water. These ions exist only for a brief time period before rejoining to form water molecules. At any given moment, only a very small proportion of the sample of water exists as ions (only about one of every billion particles). The following equation describes this process ... 

The second variation is to determine either the pH or the hydrogen ion concentration of a solution when given the hydroxide ion concentration, [OH-], for the solution. To solve these problems you need to utilize the equilibrium constant expression for the self-ionization of water (Kw). This expression will allow you to convert from the hydroxide ion concentration, [OH-], to the hydrogen ion concentration, [H+], The [H+] can then be used to calculate pH if necessary. One of the free-response questions on the 1999 test required this calculation. 

In the calculations, we have omitted the self-ionization of water. Since the equilibrium concentration of hydrogen ion [H+] is so small (1.0 X 10 7), it is negligible compared to the molarity of the acetic acid. 

Perhaps you recognize this equation from the beginning of the chapter. If you don t, it is the equation for the self-ionization of water, from which Kw was derived. What you ve now seen are three equations. The first is an equation that would be used to calculate the Ka of acetic acid. The second is the equation used to calculate Kh for the conjugate base. The third, which is derived from these two, is the formula for calculating Kw. There is a very clear relationship between Ka, Ku, and Kw. The three constants are all related in Equation 14.9, shown below ... 

As also shown in Figure 9, a pair of water molecules are in equilibrium with two ions—a hydronium ion and a hydroxide ion—in a reaction known as the self-ionization of water. 

Explain the relationship between the self-ionization of water and... 

The hydronium ion is a hydrated hydrogen ion, which means that a water molecule is attached to a hydrogen ion by a covalent bond. However, the symbols H+ and H3O+ can be used interchangeably in chemical equations to represent a hydrogen ion in aqueous solution. Thus, a simplified version of the equation for the self-ionization of water is... 

Water not only serves as the solvent in solutions of acids and bases, it also plays a role in the formation of the ions. In aqueous solutions of acids and bases, water sometimes acts as an acid and sometimes as a base. You can think of the self-ionization of water as an example of water assuming the role of an acid and a base in the same reaction. 

The result is a special equilibrium constant expression that applies only to the self-ionization of water. The constant, K, is called the ion product constant for water. The ion product constant for water is the value of the equilibrium constant expression for the self-ionization of water. Experiments show that in pure water at 298 K, [H+] and [OH ] are both equal to 1.0 X 10 M. Therefore, at 298 K, the value of K is 1.0 X 10 ... 

The product of [H+] and [OH ] always equals 1.0 X 10 at 298 K. This means that if the concentration of H+ ion increases, the concentration of OH ion must decrease. Similarly, an increase in the concentration of OH ion causes a decrease in the concentration of H+ ion. You can think about these changes in terms of Le ChStelier s principle, which you learned about in Chapter 18. Adding extra hydrogen ions to the self-ionization of water at equilibrium is a stress on the system. The system reacts in a way to relieve the stress. The added H+ ions react with OH ions to form more water molecules. Thus, the concentration of OH ion decreases. Example Problem 19-1 shows how you can use to calculate the concentration of either the hydrogen ion or the hydroxide ion if you know the concentration of the other ion. 

In this way, two neutral water molecules can react and produce a positively charged hydronium (H30+) ion and a negatively charged hydroxide (OH ) ion. This type of reaction, called the self-ionization of water, is actually a relatively rare occurrence, but it does happen. You may have measured this occurrence in the laboratory, if you ever measured the pH of pure water. Pure water at 25 °C has a pH of 7, which means that the concentration of hydronium ions, [H30+] is 1.0 X 10 7 moles/liter (remember molarity ), or 1.0 X 10 7 M. This pure water, with a pH of 7, is said to be neutral, because it has the same number of hydronium ions and hydroxide ions, which makes sense, because of the chemical equation shown preceding this paragraph. 

Using the equilibrium expression that you reviewed in Chapter 12, the equilibrium constant for the self-ionization of water can be expressed as... 

Storage of reducing power an electron from chlorophyll may be captured by a hydrogen ion (H+), produced from the self-ionization of water, to yield a hydrogen atom, which is immediately taken up by... 

NADP (nicotinamide adenine dinucleotide phosphate), storing the reducing power in the form of NADPH. The chlorophyll ion can regain an electron from a hydroxyl ion (OH ), which is also formed during the self-ionization of water, and the resulting hydroxyl radical (OH-, which has no charge) combines with others to form oxygen and water. 

Here the ammonium ion is the- conjugate acid of the base NH3, and H2O is the conjugate acid of the base OH . The fact that H2O is now acting as an acid, whereas in the HA ionization it acts as a base, is of particular importance. On account of its ability to act both as an acid and a base, H2O is said to be oinphoteric (Greek amphoteros, in both ways). In the self-ionization of water one molecule is acting as an acid.and another as a base, as is clear if we write the reaction as... 

Figure 18.13 In the self-ionization of water, one water molecule acts as an acid, and the other acts as a base. 

Write a balanced chemical equation that represents the self-ionization of water. 

Assume that the self-ionization of water makes a negligible contribution to the concentrations of hydronium ions and of hydroxide ions, which will typically be true unless the solutions are extremely dilute. 

Now let s consider a strong base such as potassium hydroxide. Let s say the solution is 1 x 10 M. Since KOH is a strong base, it dissociates completely. So we can confidently say that the concentration of hydroxide-ion formed is also 1 X 10 M. Just like the strong acid, we can ignore the possibility of hydroxide-ion formation from water, since the self-ionization of water is negligible. Since the OH" concentration is 1 x 10 , the pOH is 3. From this, we can say that the pH of this solution is close to 11, since pH + pOH =14. 

In aqueous solutions this is not a major distinction, but it was known that other acid base reactions can occur in solvents that are not water. The Arrhenius theory applies to water. But as with the water example, where the self-ionization of water occurs, hydronium ion concentration increases above a value of 10" M for acids, and those that decrease the value are bases. 

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