Autoionization of Water
Autoionization of water refers to the process by which water molecules can spontaneously dissociate into hydronium (H3O+) and hydroxide (OH-) ions. This occurs when two water molecules transfer a proton between them, resulting in the formation of these ions. The concentration of these ions in pure water is very low, but it is a crucial factor in determining the pH of a solution.
7 Key excerpts on "Autoionization of Water"
eBook - ePubAir Pollution Calculations
Quantifying Pollutant Formation, Transport, Transformation, Fate and Risks
- Daniel A. Vallero(Author)
- 2019(Publication Date)
- Elsevier(Publisher)
...The acidity and alkalinity of water are expressed as a solution’s hydrogen ion [H + ] concentration and hydroxide ion [OH − ] concentration, respectively. Unless deionized, water is the combination of the neutral molecular water and these ions. Water has the ability to act as either an acid or a base. In fact, water chemists sometimes show molecular water as HOH, perhaps to demonstrate ionization. The dissociation in water is a reversible reaction: HOH l ↔ H + aq + OH − aq (7.19) However, this equilibrium really involves the hydronium ion (H 3 O +), which is the chemical form of the hydrogen ion (H +) in an aqueous solution. When hydrogen loses its only electron, all that is left is a proton, which is highly reactive, so it immediately reacts with a water molecule, generating H 3 O +. By definition, an acid donates a proton to a solution, that is, H +, but this proton is extremely small, so it is more reasonable that it exists as hydronium. Therefore, the dissociation of water is a reversible reaction better described as 2HOH l ↔ H 3 O + aq + OH − aq (7.20) This dissociation is continuous, and the number of ions dissociating at any time is known as the water dissociation constant, K w, with the formula in the same format as any equilibrium constant discussed thus far, that is, the concentrations of products divided by the concentrations of reactants, all raised to the power of the coefficients of the balanced reaction: K w = H 3 O + OH − HOH 2 (7.21) However, the ions represent so small a fraction of the total mass in this reaction that the molecular water is assumed to be essentially pure water. Since pure substances are solvents without solutes, that is, concentration is meaningless, they are not included in equilibrium calculations. At 25°C, this dissociation reaction is. K w = H 3 O + OH − = 1.0 × 10 − 14 mol L − 1 (7.22) Recall that the brackets denote concentration. Because this equilibrium reaction is balanced, concentration proportions are 1:1...
eBook - ePub- J Fisher, J.R.P. Arnold, Julie Fisher, John Arnold(Authors)
- 2020(Publication Date)
- Taylor & Francis(Publisher)
...10 −7 M). If a source of H + or − OH is present other than the water itself, that is the ions are present at a concentration greater than 10 −7 M, then the solution is acidic. Conversely, if the concentration of H + is less than 10 −7 M (hence [ − OH ] is greater than 10 −7 M) the solution is referred to as alkaline or basic. There is, however, a more convenient way of expressing the acidity or basicity (alkalinity) of a solution and that is by using the pH scale. The pH of a solution is simply the logarithm (to the base 10) of the reciprocal of the hydrogen ion activity, which under conditions of dilute solutions may be approximated to the concentration, as follows; Therefore, Of course, as both K a and K b (Section N2) are related to the ionic product the pH scale may also be applied to these terms. Thus; The pH scale is thus a convenient system for specifying the acidity of a solution in terms of small, and for the most part, positive numbers, over the range 0 to 14, in general. It must be noted that as the ionic product of water varies with temperature, so too does the pH. Hence, when quoting pH values the temperature for the measurement must be specified. Neutralization and hydrolysis The original theory of neutralization, described by Arrhenius, stated that neutralization was the result of the reaction between an acid and base to produce a salt and water. Extending this to the more general situation when water may not necessarily be involved, neutralization may be described as follows; The extent to which the neutralization reaction proceeds depends not only on the strength of the acid and base components but also on the nature of the solvent. If the solvent is amphiprotic (i.e. able to act as a base or an acid) as is water for example, it may react with the products of neutralization as follows; The neutralization reaction is consequently reversed...
eBook - ePubThe Science For Conservators Series
Volume 2: Cleaning
- Matthew Cushman(Author)
- 2005(Publication Date)
- Routledge(Publisher)
...Moreover, because the ions are chemically more reactive than their parent molecules, their presence strongly influences the chemical interaction of water with other substances. A chemical equilibrium, like other forms of equilibrium or stability, can be upset by suitable external influences. The conditions of acidity and alkalinity are just this. The equilibrium is disturbed so that the concentrations of H 3 O + ions or OH – ions are no longer one ten-millionth of a mole per litre. In acidic solutions the concentration of H 3 O + is increased by hundreds, thousands or millions of times. Alkaline solutions, conversely, have the concentrations of OH – ions dramatically increased. Thus, the chemical behaviour of the solution becomes controlled by the behaviour of these ions. The compounds called acids and alkalies can bring about these remarkable changes in water when they go into solution. acidity and alkalinity A2 The pH Scale for Hydrogen Ion Concentrations The concentration of H 3 O + and OH – ions in pure water is one ten-millionth of a mole per litre. Written as a fraction this is which can be written more compactly as 10 –7, to be read as “ten to the minus seven”. The convention for describing numbers like this is simply to count how many noughts there are in the number. Numbers bigger than 1 are given a plus index; thus 1000 is 10 +3, ie “ten to the plus three” (normally just 10 3 or “ten to the power of three”). Fractions are indicated with a minus index. The fraction is 10 –3, “ten to the minus three”. It is long-winded to refer to concentrations in moles per litre when the numbers become awkward mouthfuls like “one ten-millionth” so a shorthand convention based on the “ten-to-the-something” system has been adopted...
eBook - ePub- Arnost Kotyk, Jan Slavik(Authors)
- 2020(Publication Date)
- CRC Press(Publisher)
...Thus, 2 H 2 O ⇌ H 3 O + + OH − (4) Even in an acid solution of pH = 0 the concentration of free protons was shown to be about 10 −130!2 However, this way of expressing the hydrogen ion concentration does not reflect its actual solvation degree in water which, in fact, shows highest stability for H 3 O + ⋅3H 2 O and for OH −⋅3H 2 O. From Equation 4, then, the dissociation constant of water. is K A = a H 3 O + ⋅ a OH − / a H 2 O ≃ c H 3 O + γ H 3 O + ⋅ c OH − γ OH − / c H 2 O γ H 2 O (5) The dissociation of water being negligible we may consider the concentration of H 2 O as constant and the activity coefficients equal to unity. The concentration of water itself is 55.35 mol dm −3 at 25°C and it is generally included in what is termed the ionization constant (ion product) of water, or K w = a H 3 O + ⋅ a OH − (6) The value of K w is almost exactly 10 −14 mol dm −3 at 25°C and depends on temperature as shown in Figure 1. The degree of dissociation of water α (= c H 3 O + / c H 2 O = c OH − / c H 2 O) at 25°C is, thus, 1.41·10 −9. FIGURE 1. Temperature dependence of the ionization constant (ion product) of water. Because of the temperature dependence, neutrality (identical number of plus and minus charges in solution) is achieved at 6.81 at human body temperature. For the sake of comparison, it should be noted that the self-ionization (autoprotolysis) constants of protogenic solvents (strong acids) are high (2.69 · 10 −4 M for sulfuric acid), but much like water for weaker acids or alcohols (3.16 · 10 −15 M for acetic acid and 2.51.10 −17 M for ethanol), and are extremely low for protophilic solvents (bases), e.g., 10 −29 M for ammonia. II...
eBook - ePub- Jean-Louis Burgot(Author)
- 2019(Publication Date)
- CRC Press(Publisher)
...(This remark anticipates a brief comparison between acid-base and redox processes which is done later see the following chapter). In aqueous solutions, Arrhenius and Brönsted-Lowry theories can be considered as being equivalent. 2) Quantitative aspects of acid-base reactions 2-1) Ionic product of water Water, even purissime, is slightly ionized. This has been discovered thanks to electrical measurements. Its ionization is the mark of the following equilibrium: H 2 O + H 2 O ⇌ H 3 O + (w) + OH − (w) The mass law permits to write: K = a H 3 O + a OH − / a H 2 O 2 where K is the equilibrium thermodynamic constant of the reaction above. Electrical measurements give the following concentrations, for pure water at 25°C: [ H 3 O + ] = [ OH − ] = 10 − 7 mol L − 1 Thus, pure water, even “purissime”, contains ions but very few. As a consequence, the molar fraction of molecular water can be safely taken to be equal to unity and its activity as well (see Chapter 19). As a result, one can write: K = a H 3 O + a OH − (152) The product a H3O+ a OH is named the ion-product constant for water or ionic product of water and is symbolized by K w : K w = a H 3 O + a OH − The ionic product varies with temperature. Dissociation increases with it: at 25 ° C K w = 10 − 14 at 50 ° C K w = 5, 6 10 − 14 at 100 ° C K w = 6, 0 10 − 13 In every aqueous solution, which is diluted in ions, the ionic product of water is a constant at a given temperature, no matter what the OH − and H 3 O + ions concentrations may be, provided the solution remains diluted. It is said: “the water ion product is satisfied”. Solutions for which: [ H 3 O + ] = [ OH − ] are said neutral, [ H 3 O + ] > [ OH − ] are said acidic, [ H 3 O + ] < [ OH − ] are said basic. 2-2) Constant of dissociation acid Ka (pKa) For one acid, in dilute aqueous solution, one defines the acid dissociation constant K a by: K a = a H + a A − / a HA (153) where a H+, a A, a HA are the activities of the species symbolized in subscript...
eBook - ePubPhysicochemical and Biomimetic Properties in Drug Discovery
Chromatographic Techniques for Lead Optimization
- Klara Valko(Author)
- 2013(Publication Date)
- Wiley(Publisher)
...Changing the acid/base character and the percentage ionized form of the drug molecule provides an excellent tool for medicinal chemists to tune the compound's distribution in the body. Before the experimental descriptions of the measurement of values, it is important to emphasize that the definition of refers to an aqueous environment. When the water is mixed with organic solvents or the compound is near the membrane or proteins, its shifts depending on the H-bond donor, acceptor property of the solvent, and its dielectric constant, dipolarity/polarizability. The Born equation describes the changes in the values of the compounds due to the electrostatic effects of the solvents [2], as is shown by Equation 8.6. 8.6 In Equation 8.6, is the number of charged species in the ionization process, is the common radius of all the ions, and is the dielectric constant of the solvent. Typically, when the water is mixed with organic solvents or proteins and membranes are in the aqueous compartment, the dielectric constant, H-bond acidity, and basicity of the surrounding are decreased relative to the net water that results in the weakening of both acids and bases. The of an acid increases and the of a base decreases, which means that ionization process is suppressed. Table 8.1 shows examples of the shift of simple basic and acidic compounds in increasing methanol concentration [2]. Another factor that affects the value is the ionic strength of the solution. The effect of ionic strength on the acid dissociation constant can be calculated using the Debye–Hückel equation [3]. The concentration of the compounds in ionic solutions cannot be considered as ideal even with low ionic strength. The concentration should be multiplied with the so-called activity coefficient that describes more precisely the activity of various species of molecules in solution, and the activity values should be used instead of concentration when describing chemical equilibrium constants...
eBook - ePub- James F. Pankow(Author)
- 2019(Publication Date)
- CRC Press(Publisher)
...For example, when discussing the aqueous chemistry of transition metal ions, specialists sometimes specifically note the association with water molecules; e.g., aqueous Cu 2+ is sometimes represented as Cu(H 2 O) 6 2 +. Usually, this is not done, though, because such representations obscure most discussions of reactions and equilibria. In any case, whenever we consider any species in aqueous solution, it should be implicitly understood that each species is surrounded and solvated at least to some extent by water molecules. For some species, like H +, Fe 3+, and Al 3+, that association with water is very close; for others, like I −, the association is much weaker. Many chemistry texts explicitly recognize the strong solvation of H + by water and argue that as soon as any H + is formed by a reaction such as Eq.(3.14), reaction with water proceeds according to H + + H 2 O = H 3 O + K = [H 3 O + ] [H + ][H 2 O] (3.16) with H 3 O + named the “hydronium” ion. The net reaction for Eqs.(3.14) and (3.16) is then HA + H 2 O = H 3 O + + A − (3.17) K a = [H 3 O + ] [ A − ] [HA][H 2 O] = [H 3 O + ] [ A − ] [HA]. (3.18) Equation (3.18) follows the usual convention that the concentration scale chosen for water (as the solvent) is the mole fraction scale so that in most solutions [H 2 O] ≈ 1. So, how can one text use the H + representation and another use the H 3 O + representation, and both still come out with the same answer for pH and other parameters for a given solution of interest? The answer is based on the fact that neither representation is exactly correct, and in fact, they both refer symbolically to all forms of H + in aqueous solution, with every possible combination of interactions with. H 2 O...
