Hydronium/hydroxide ions

Strong and Weak Bases Just as the acidity of an aqueous solution is a measure of the concentration of the hydronium ion, H3O+, the basicity of an aqueous solution is a measure of the concentration of the hydroxide ion, OH . The most common example of a strong base is an alkali metal hydroxide, such as sodium hydroxide, which completely dissociates to produce the hydroxide ion. 

The relative importance of the potential catalytic mechanisms depends on pH, which also determines the concentration of the other participating species such as water, hydronium ion, and hydroxide ion. At low pH, the general acid catalysis mechanism dominates, and comparison with analogous systems in which the intramolecular proton transfer is not available suggests that the intramolecular catalysis results in a 25- to 100-fold rate enhancement At neutral pH, the intramolecular general base catalysis mechanism begins to operate. It is estimated that the catalytic effect for this mechanism is a factor of about 10. Although the nucleophilic catalysis mechanism was not observed in the parent compound, it occurred in certain substituted derivatives. 

A second way to achieve constancy of a reactant is to make use of a buffer system. If the reaction medium is water and B is either the hydronium ion or the hydroxide ion, use of a pH buffer can hold Cb reasonably constant, provided the buffer capacity is high enough to cope with acids or bases generated in the reaction. The constancy of the pH required depends upon the sensitivity of the analytical method, the extent of reaction followed, and the accuracy desired in the rate constant determination. 

Throughout this section the hydronium ion and hydroxide ion concentrations appear in rate equations. For convenience these are written [H ] and [OH ]. Usually, of course, these quantities have been estimated from a measured pH, so they are conventional activities rather than concentrations. However, our present concern is with the formal analysis of rate equations, and we can conveniently assume that activity coefficients are unity or are at least constant. The basic experimental information is k, the pseudo-first-order rate constant, as a function of pH. Within a senes of such measurements the ionic strength should be held constant. If the pH is maintained constant with a buffer, k should be measured at more than one buffer concentration (but at constant pH) to see if the buffer affects the rate. If such a dependence is observed, the rate constant should be measured at several buffer concentrations and extrapolated to zero buffer to give the correct k for that pH. 

Except for those reactions whose characteristic rate constants vary linearly with the hydronium or hydroxide ion concentration, the most effective presentation of pH-rate data is a graphical one. Two kinds of plot pH-rate profiles) are commonly seen ... 

Now suppose that only the monoanionic form of the dibasic acid H2S undergoes reaction and that neither the hydronium nor the hydroxide ion is directly involved. The kinetic scheme is, therefore. 

Hydronium ion, electrostatic potential map of, 145 Hydrophilic, 63 Hydrophobic. 63 Hydroquinone, 631 from quinones, 631 Hydroxide ion, electrostatic potential map of. 53. 145... 

A species that can either accept or donate a proton is referred to as amphiprotic. An example is the H20 molecule, which can gain a proton to form the hydronium ion, H30+, or lose a proton, leaving the hydroxide ion, OH-. 

Hydronium ion, 187 concentration calculation, 192 concentration and pH, 190 model, 186 Hydroquinone, 345 Hydrosphere, 437 composition, 439 Hydroxide ion, 106, 180 Hydroxides of lhird row, 371 Hydroxylamine, 251 Hydroxyl group, 329 Hypobromiie ion, 422 Hypochlorite ion, 361 Hypochlorous acid, structure, 359 Hypophosphorous acid, 372 Hypothesis, Avogadro s, 25, 52... 

For water, the second-order rate coefficient was determined as 9.5 x 10 12 by extrapolation from data at higher temperatures and using the presence of hydroxide ion to suppress any reaction with hydronium ion. For reaction with solutions of biphosphate and ammonium ions, since reaction via hydronium ions in these media is negligible (ca. 1 % of the total rate), the second-order rate coefficients were evaluated from exchange data at a single acid concentration as k2 (H2PC>4 ) = 3.89 xlO-7 and (NH ) = 5.0 x 10-9, the latter value being corrected for the water-catalysed reaction. 

Recall that an Arrhenius acid is a compound that produces hydronium ions in water and an Arrhenius base is a compound that produces hydroxide ions in water. 

The concentrations of H30 + and OH are very low in pure water, which explains why pure water is such a poor conductor of electricity. To imagine the very tiny extent of autoprotolysis, think of each letter in this book as a water molecule. We would need to search through more than 50 books to find one ionized water molecule. The autoprotolysis reaction is endothermic (AH° = +56 kj-mol l), and so we can expect Kw to increase with temperature, and aqueous solutions to have higher concentrations of both hydronium and hydroxide ions at higher temperatures. Unless otherwise stated, all the calculations in this chapter will be for 25°C. 

FIGURE 10.10 The product of the concentrations of hydronium and hydroxide ions in water (pure water and aqueous solutions) is a constant. If the concentration of one type of ion increases, the other must decrease to keep the product of the ion concentrations constant. 

FIGURE 10.9 As a result of autoprotolysis, pure water consists of hydronium ions and hydroxide ions as well as water molecules. The concentration of ions that results from autoprotolysis is only about 10 mol-L and so only about I molecule in ZOO million is ionized. The overlay shows only the ions. 

Proton transfer equilibrium is established as soon as a weak base is dissolved in water, and so we can calculate the hydroxide ion concentration from the initial concentration of the base and the value of its basicity constant. Because the hydroxide ions are in equilibrium with the hydronium ions, we can use the pOH and pKw to calculate the pH. 

Sometimes we need to know how the concentrations of the ions present in a solution of a polyprotic acid vary with pH. This information is particularly important in the study of natural waters, such as rivers and lakes (Box 10.1). For example, if we were examining carbonic acid in rainwater, then, at low pH (when hydronium ions are abundant), we would expect the fully protonated species (H2C03) to be dominant at high pH (when hydroxide ions are abundant), we expect the fully deprotonated species (C032 ) to be dominant at intermediate pH, we expect the intermediate species (HC03, in this case) to be dominant (Fig. 10.20). We can verify these expectations quantitatively. 

Hydronium and hydroxide ions appear to move through water much faster than other kinds of ions. Explain this observation. 

The apparent motion of hydronium and hydroxide ions is not as dependent on the diffusion of individual ions as is that of other ions they are formed and reformed as protons are transferred from and to water molecules in solution. Autoprotolysis allows rapid proton transfer between water molecules. 

The ability of water to ionize, while shght, is of central importance for life. Since water can act both as an acid and as a base, its ionization may be represented as an intermolecular proton transfer that forms a hydronium ion (HjO ) and a hydroxide ion (OH ) ... 

One of the most fundamental chemical reactions is the combination of a hydroxide ion (OH ) and a hydronium ion (H3 0+) to produce two molecules of water OH" (a g) + H3 (a g) 2 H2 O (/) A molecular view of this reaction (Figure 4-7f shows that the hydroxide anion accepts one hydrogen atom from the hydronium cation. Taking account of charges, it is a hydrogen cation (H ) that is transferred. The reaction occurs rapidly when H3 O and OH ions collide. The hydroxide anion accepts a hydrogen cation from the hydronium cation, forming two neutral water molecules. 

In Figure the hydronium ion acts as an acid because it donates a proton to a base. The hydroxide anion acts as a base because it accepts a proton from an acid. When a hydronium ion with charge +1 transfers a proton to a hydroxide ion with charge -1, the two resulting water molecules have zero charges. The pair of charges becomes a neutral pair. A proton transfer reaction such as this one, in which water is one product and a pair of charges has been neutralized, is called a neutralization reaction. 

Weak acids are compounds that donate protons quantitatively to hydroxide ions but not to water. All acids are proton donors, but whereas a strong acid quantitatively donates protons to water, a weak acid does not. Aqueous solutions of weak acids contain small concentrations of hydronium ions, making the solutions acidic, but nearly all the weak acid molecules remain intact. Representative of strong acids, HCl generates H3 O and Cl" quantitatively when dissolved in water. Representative of weak acids, HF remains predominantly as HF molecules when dissolved in water. However, HF donates protons quantitatively to OH ions to give H2 O molecules and F ions, as shown in Figure 4-9. 

In pure water, the hydronium and hydroxide ion concentrations are equal We find these concentrations by taking the square root of K- [H3 0 ] = [OH jg = 1.0 X 10 M (pure water at 25 C) Equal concentrations of these two ions means that pure water is neither acidic nor basic. 

The water equilibrium describes an inverse relationship between [H3 0+] eq nd [OH-]gq. When an acid dissolves in water, the hydronium ion concentration increases, so the hydroxide ion concentration must decrease to maintain the product of the concentrations at 1.0 X 10. Similarly, the hydroxide ion concentration increases when a base dissolves in water, so the hydronium ion concentration must decrease. 

In any solution of an acid, the total hydronium and hydroxide ion concentrations include the 10" M contribution from the water reaction. This example illustrates, however, that the change in hydronium ion concentration due specifically to the water equilibrium is negligibly small in an aqueous solution of a strong acid. This is true for any strong acid whose concentration is greater than 10 M. Consequently, the hydronium ion concentration equals... 

For any aqueous strong base, the hydroxide ion concentration can be calculated directly from the overall solution molarity. As is the case for aqueous strong acids, the hydronium and hydroxide ion concentrations are linked through the water equilibrium, as shown by Example. ... 

For this example, we summarize the first four steps of the method The problem asks for the concentration of ions. Sodium hydroxide is a strong base that dissolves in water to generate Na cations and OH- anions quantitatively. The concentration of hydroxide ion equals the concentration of the base. The water equilibrium links the concentrations of OH" and H3 O" ", so an equilibrium calculation is required to determine the concentration of hydronium ion. What remains is to organize the data, carry out the calculations, and check for reasonableness. 

What are the concentrations of hydronium and hydroxide ions in a beverage whose pH = 3.05 ... 

A pH around 3 represents an acidic solution, so we expect the hydronium ion concentration to be much larger than the hydroxide ion concentration. Remember that although pH and. w are dimensionless, concentrations of solutes always are expressed In mol/L. Our results have two significant figures because the logarithm has two decimal places. 

A logarithmic scale is useful not only for expressing hydronium ion concentrations, but also for expressing hydroxide ion concentrations and equilibrium constants. That is, the pH definition can be generalized to other quantities pOH = - log [OH ] p Tg = - log Tg p log... 

Most acids and bases are weak. A solution of a weak acid contains the acid and water as major species, and a solution of a weak base contains the base and water as major species. Proton-transfer equilibria determine the concentrations of hydronium ions and hydroxide ions in these solutions. To determine the concentrations at equilibrium, we must apply the general equilibrium strategy to these types of solutions. 

To protect a solution against pH variations, a major species in the solution must react with added hydronium ions, and another major species must react with added hydroxide ions. The conjugate base of a weak acid will react readily with hydronium ions, and the weak acid itself will react readily with hydroxide ions. This means that a buffer solution can be defined in terms of its composition. 

Redraw the original figure to show the equilibrium condition that is established when (a) three hydroxide ions enter the region, and (b) seven hydronium ions enter the region. Include any water molecules that are part of the buffer chemistry. 

When small amounts of hydronium or hydroxide ions are added to a buffer solution, the pH changes are very small. There is a limit, however, to the amount of protection that a buffer solution can provide. After either buffering agent is consumed, the solution loses its ability to maintain near-constant pH. The buffer capacity of a solution is the amount of added H3 O or OH that the buffer solution can tolerate without exceeding a specified pH range. 

Now consider what happens if some acid is added to this saturated solution. Hydronium ions react with hydroxide ions to form water OH a q) + H3 O (a 2H2 0(/) By Le Chatelier s principle, the system responds In... 

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