Hydroxide ion concentration

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 ... 

A similar approach is used for the hydroxide ion concentration. The pOH of a solution is defined as... 

A measured volume, 10.00 liters, of the waste process water from a cotton mill require 23.62 ml of 0.1000 M hydrochloric acid to produce a neutral solution. What is the hydroxide ion concentration in the waste ... 

The addition of half a mole of ammonium chloride to 1 litre of a 0.1M solution of aqueous ammonia has decreased the degree of ionisation from 1.35 to 0.0036 per cent, and the hydroxide ion concentration from 0.00135 to 0.000 0036 mol L-1. 

In all cases the reaction of the solution can be quantitatively expressed by the magnitude of the hydrogen ion (or hydroxonium ion) concentration, or, less frequently, of the hydroxide ion concentration, since the following simple relations between [H + ] and [OH-] exist ... 

The hydrogen ions required for this reaction can be obtained only from the further dissociation of the water this dissociation produces simultaneously an equivalent quantity of hydroxyl ions. The hydrogen ions are utilised in the formation of HA consequently the hydroxide ion concentration of the solution will increase and the solution will react alkaline. 

If the solution were completely hydrolysed, the concentration of acetic (ethanoic) acid produced would be 0.01 molL-1. But the degree of hydrolysis is 0.0238 percent, therefore the concentration of acetic acid is 2.38 x 10 6 mol L 1. This is also equal to the hydroxide ion concentration produced, i.e. pOH = 5.62. 

B. Precipitation and separation of hydroxides at controlled hydrogen ion concentration or pH. The underlying theory is very similar to that just given for sulphides. Precipitation will depend largely upon the solubility product of the metallic hydroxide and the hydroxide ion concentration, or since pH + pOH = pKw (Section 2.16), upon the hydrogen ion concentration of the solution. 

As shown above the sulphide ion concentration of a saturated aqueous solution of hydrogen sulphide may be controlled within wide limits by suitably changing the concentration of hydrogen ions—a common ion—of the solution. In a like manner the hydroxide ion concentration of a solution of a weak base, such as aqueous ammonia (Kb = 1.8 x 10-5), may be regulated by the addition of a common ion, e.g. ammonium ions in the form of the completely dissociated ammonium chloride. The magnitude of the effect is best illustrated by means of an example. In a 0.1M ammonia solution, the degree of dissociation is given (Section 2.13) approximately by. 

An immediate application of the use of the aqueous ammonia-ammonium chloride mixture may be made to the familiar example of the prevention of precipitation of magnesium hydroxide (S.P. 1.5 x 0-11). We can first calculate the minimum hydroxide ion concentration necessary to prevent precipitation in, say, 0.1M magnesium solution. 

Now consider the conditions necessary for the practically complete precipitation of magnesium hydroxide from a 0.1M solution of, say, magnesium chloride. A pOH slightly in excess of4.9(i.e. pH = 9.1) might fail to precipitate the hydroxide owing to supersaturation. Let us suppose the hydroxide ion concentration is increased ten-fold, i.e. to pOH 3.9 or pH 10.1, then, provided no supersaturation is present ... 

Lewis and Suhr (1958b) observed the disappearance of the 4-nitrobenzenediazo-nium ion by spectrophotometry, and found that the rate increased with increasing hydroxide ion concentration in the pH range 7.5-10. At pH 10-14 and in concentrated NaOH solution the rate was independent of pH. Subsequently the rates were... 

A reaction mechanism in which the ( )-diazoate is formed by attack of the diazonium ion by a hydroxide ion in such a way that the ( )-diazoate is the primary intermediate (i. e., reaction sequence 6 - 3 in Scheme 5-14) is not consistent with the observation that the isomerization rate constant is independent of the hydroxide ion concentration. 

Analysis of the decay of the sum of the diazonium ion and (jF)-diazoate concentrations as a function of time reveals that there are two reactions. The first is observed only at the beginning and at relatively low temperatures (20 °C) it is first order in relation to the above sum of concentrations and to the hydroxide ion concentration. The second is a very complex function of the hydroxide ion concentration, so that a mechanistic interpretation was not possible. 

If one limits the consideration to only that limited number of reactions which clearly belong to the category of nucleophilic aromatic substitutions presently under discussion, only a few experimental observations are pertinent. Bunnett and Bernasconi30 and Hart and Bourns40 have studied the deuterium solvent isotope effect and its dependence on hydroxide ion concentration for the reaction of 2,4-dinitrophenyl phenyl ether with piperidine in dioxan-water. In both studies it was found that the solvent isotope effect decreased with increasing concentration of hydroxide ion, and Hart and Bourns were able to estimate that fc 1/ for conversion of intermediate to product was approximately 1.8. Also, Pietra and Vitali41 have reported that in the reaction of piperidine with cyclohexyl 2,4-dinitrophenyl ether in benzene, the reaction becomes 1.5 times slower on substitution of the N-deuteriated amine at the highest amine concentration studied. 

First, consider the base hydrolysis of pentaammine(bromo)cobalt(2+) ion, Eq. (I -27).4 The rate is directly proportional to the hydroxide ion concentration,... 

The composition of the transition state is [H2O 4-1- + OC1" - OH-], or [IOC1H- H20]. This and the inverse dependence on hydroxide ion concentration suggest the following equations ... 

The absolute (a) and relative (b) amounts of Pu(IV) hydroxide ion concentration (or pH), calculated by Equation 4 based on the data given in Table I. The region of interest for the present investigation is marked by dotted lines. 

Find the mole ratio of hydroxide ion concentration to solute concentration. 

Calculate the hydroxide ion concentration from the solute concentration. 

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. 

The calculation of x can often be simplified, as explained in Toolbox 10.1. We ignore contributions from the autoprotolysis of water to the hydroxide ion concentration if the concentration of hydroxide ions is greater than 10 h mol-L-. Step 5 Determine the pOH of the solution and then calculate the pH from the pOH by using Eq. 6b. 

Below is the titration curve for the neutralization of 25 mL of a base with a strong monoprotic acid. Answer the following questions about the reaction and explain your reasoning in each case, (a) Is the base strong or weak (b) What is the initial hydroxide ion concentration of the base (c) What is Kh for the base (d) What is the initial concentration of the base (e) What is the concentration of acid in the titrant (f) Use Table 11.3 to select an indicator for the titration. 

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. ... 

Just as for solutions of strong acids, the water equilibrium contributes negligibly to the total hydroxide ion concentration in any solution of strong base whose concentration is greater than 10 M. Consequently, the... 

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... 

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