Solutions with a high concentration of h+ ions are

17.16.1 Hydrogen Ion Concentration

Hydrogen ion concentration (the pH) is one of the important factors that affect growth and multiplication of algae and hence the oil and biodiesel production. Most algal growth occurs in the region of neutral pH, although optimum pH is the pH of initial culture in which an alga is adapted to grow [145]. Bartley et al. [146] found that pH of around 8 seems most beneficial for maximum growth rate and lipid accumulation of Nannochloropsis salina and to minimize invading organisms. However, adding buffers will not be cost-effective or realistic at a large scale. They also demonstrated that higher pH values per se do not slow Nannochloropsis production. Thus, the addition of CO2 at large scales is mostly valuable for providing an inorganic carbon source for algae.

Moheimani [147] found pH 7 and 7.5 to be ideal for lipid accumulation in Tetraselmis suecica and Chlorella sp. While, Bartley et al. [146] found no significant effect of pH change on lipid accumulation, the treatment with a pH change to 8 exhibited the greatest overall accumulation (averaging 24.75% by mass) of N. salina. Rodolfi et al. [148] found the lipid content (% biomass) for different Nannochloropsis spp. to be 24.4–35.7%. The earlier results indicate that pH may not be an important stress factor that triggers increased lipid accumulation in microalgae. Acidic pH of culture media can alter nutrient uptake or induce metal toxicity and therefore have an effect on algal growth and oil production [149]. The green microalga Chlamydomonas acidophila and the diatom Pinnularia braunii accumulate storage lipids, such as triacylglycerides, under extremely acidic environment (pH 1) [150]. However, basic pH decreases membrane-associated polar lipids due to cell cycle inhibition. In basic pH conditions, membrane lipids in Chlorella were observed to be less unsaturated [151].

6.5.2.1 Chemical biosensors

Hydrogen ion concentration is a critical biomarker to determine the wound status. For healthy skin, the pH value is approximately in the range of 5.5 but for infected wounds, the pH value is in the range of 7-8.5. The alkaline nature of pH in the wound is due to the presence of bacterial colonies and enzymes. When a wound is kept in an acidic condition, the fibroblasts proliferate more actively and the wound healing process is stimulated more while an infected wound shows a slightly basic pH environment due to certain enzyme activities, bacterial colonization, and formation of protein structures [83]. Despite the conflicting information on wound pH, the pH has been identified as a significant parameter in determining the phase of the healing process and bacterial colonization [84]. Consequently, several research groups have developed dressings which incorporate pH-sensitive materials. For example, a hydrogel-based wireless pH sensor embedded into a wound dressing has been reported to continuously monitor pH [85]. The device consists of a pH-sensitive polyvinyl alcohol-polyacrylic acid hydrogel placed between two planar spiral coils that act as an inductive transducer (Fig. 6.6A). As pH changes, the hydrogel swells and deswells and the distance separating the two planar coils change which results in a change of inductance of the coil and the frequency response of the transducer. It was observed that a linear inductance responds, over the pH ranging from 2 to 7, with a change in the coil separation distance. A network-spectrum analyzer was also used with an antenna to allow for wireless measurement of the coil gap, and hence pH value. Likewise a potentiometric pH sensor embedded on a commercial adhesive bandage has been reported for wound pH monitoring [81]. In this case, the silver/silver chloride (Ag/AgCl) and carbon electrodes have been screen-printed onto the bandage. The Ag/AgCl reference electrode has been partially coated with a polyvinyl butyral polymer (PVB) whereas the carbon electrode serves as a working electrode and has been electropolymerized with polyaniline (PANi). Fig. 6.6B shows the fabrication process of the printed potentiometric sensor on an adhesive bandage. This wearable pH sensor showed a response in a limited pH range of 5.5-8 and relatively long-time intervals (up to 100 minutes) for the detection of pH fluctuations at a wound site (Fig. 6.6B). The pH bandage sensor exhibited the pH sensitivity close to the theoretical Nernstian response (59.2 mV/pH) (Fig. 6.6B) in the pH range 4.35-8. Further, less interference to other ions was observerd along with fast response time, good repeatability, reproducibility, and lack of hysteresis effect. This sensor also shows a minimal impact on the sensing performance during different bending cycles. Due to its similarity in the chemical environment in the vicinity of a wound, the sensor could also find application in human serum [81].

Another example relates to the bandage for monitoring irregular bleeding, pH levels, and external pressure at the wound site [78]. This bandage has low-cost sensor integrated with a wireless monitoring system for continuous monitoring of wound healing, as shown in Fig. 6.6C. The output of the sensor shows ‘Bandage OK’ when there is no bleeding. When there is bleeding, the bandage communicates with the receiver which displays the ‘Change Bandage’ sign [78]. Recently another new work reported thread-based sensors for the in vitro and in vivo analysis of glucose, pH, strain, and temperature levels [77]. The possibilities of simultaneous monitoring of pH and glucose in wound fluid by using fluorescence changes for different chemical methods have been explored as well [84]. There is not much work in the area of bandage-based glucose sensors for wound monitoring. In addition to this, the measurement of tissue oxygenation is another major wound healing inhibitors in chronic wounds [87]. A detailed study is required for the development and embedding of such sensors in bandages for wound monitoring.

11.10 Effect of pH

A hydrogen-ion concentration (pH) affects the activity of enzymes and therefore the microbial growth rate. The optimal pH for growth may be different from that for product formation. Generally, the acceptable pH range varies about the optimum by 1 to 2 pH units. Different organisms have different pH optima: the pH optimum for many bacteria ranges from pH 3–8; for yeast, pH 3–6; for molds, pH 3–7; for plant cells, pH 5–6; and for animal cells, pH 6.5–7.5. Many organisms have mechanisms to maintain intracellular pH at a relatively constant level in the presence of fluctuations in environmental pH. When pH differs from the optimal value, the maintenance-energy requirements increase. One consequence of different pH optima is that the pH of the medium can be used to select one organism over another.

In most fermentations, pH can vary substantially. Often, the nature of the nitrogen source can be important. If ammonium is the sole nitrogen source, hydrogen ions are released into the medium as a result of the microbial utilization of ammonia, resulting in a decrease in pH. If nitrate is the sole nitrogen source, hydrogen ions are removed from the medium to reduce nitrate to ammonia, resulting in an increase in pH. Also, pH can change because of the production of organic acids, the utilization of acids (particularly amino acids), or the production of bases. The evolution or supply of CO2 can alter pH greatly in some systems (eg, seawater or animal cell culture). Thus, pH control by means of a buffer or an active pH control system is important. Variation of specific growth rate with pH is depicted in Fig. 11.8, indicating a pH optimum.

The pH scale

10

The hydrogen ion concentration of one molar hydrochloric acid, pure water and one molar sodium hydroxide are 1, 10−7 and 10−14 mol dm−3 respectively. It is convenient to construct a scale of simpler numbers to represent these values. This can be achieved by taking the reciprocal of the logarithm to the base ten of the hydrogen ion concentration of the solution. When the conversion is made the value is called the pH value of the solution as shown by the conversions overpage.

When [H3O+] = 1 mol dm−3,

pHvalue=1log10[H3O+]=−log10[H3O+]=−log101=0

When [H3O+] = 10−7 mol dm−3,

pHvalue=1log10[H3O +]=−log10[H3O+]=−log1010−7=7

When [H3O+] = 10−14 mol dm−3,

pH value=1log10[H3O+]= −log10[H3O+]=−log10 10−14=14

the pH scale has been selected with values between 0 and 14 corresponding to hydrogen ion concentration of 1 mol dm−3 and 10−14 mol dm−3. A knowledge of the pH value of a solution gives a value for the hydrogen ion concentration of that solution. For example, if a solution of hydrochloric acid, HCl(aq), has a pH of 4. this means that for the solution, using the equation

pHvalue=−log10[H3O+]then4=−log10[H3O+]

Rearranging this expression gives

and taking antilogs gives

Hence the concentration of the acid must be 10−4 M.

The calculation of the hydrogen ion concentrations in weak acids from pH values does not give the concentration of the weak acid directly but if the degree of ionisation is known then the concentration can be found. For example, the pH of a solution of chloroethanoic acid is 1.72 when α = 0.064. This means that the hydrogen ion concentration of the solution given by the equation

and in the form [H3+O] = antilog (-pH) is [H3+O] = antilog (-1.72) or [H3+O] = 0.019 mol dm−3. By using the relationship

[H3+O]equil =[H3+O]solution×αthen0.019=[H3+O]solution×0.064or [H3+O]solution=0.0190.064=0.3

Hence the concentration of the acid must be 0.3 mol dm–3.

11

An equivalent scale can be applied to the concentration of hydroxide ions using the definition

pOH=1log10[OH−] =-log10[OH−]

12

On these combined scales of pH and pOH it can be shown that because for water when pH = pOH = 7 that pH + pOH = 14. This relationship is useful in the interconversion of values. For example, the pOH at a 0.01 M solution of sodium hydroxide is 2, the pH of the same solution must be 14-2 = 12.

6.06.4.2.2 pH-responsive molecular brushes

Potentiometric hydrogen ion concentration (pH) is an important stimulus, which can be addressed through pH-responsive materials.311,330–335 Upon pH variation, ionizable polymers with a pKa between 3 and 10 (weak acids and bases) exhibit a change in the ionization state leading to conformational changes. The classical monomers are acrylic acid (AA), methacrylic acid (MAA), maleic anhydride (MA), and N,N-dimethylaminoethyl methacrylate (DMAEMA). For example, the pKa of PAA depends on molecular weight and ranges within 6.8–7 for molecular weights on the order of 100 kDa.336 When the pH of the solution is below the pKa value, the polymer is in a compact collapsed form. As the pH increases above the pKa, the polymer exhibits fully stretched conformation due to the electrostatic repulsion between the segments.337 The unique properties of pH-responsive polymers arise from the facile pH adjustment, which induces ionic interaction and hydrogen bonding, resulting in a reversible microphase separation or self-organization phenomenon. Thus, pH-responsive polymeric systems provide the possibility of preparation of smart functional materials that can be used for potential therapeutic applications, for example, controlled drug delivery based on pH-triggered release.

Several strategies have been employed to demonstrate the effect of pH on the conformation and aggregation behavior of molecular brushes. One strategy was to mimic proteoglycans – polyelectrolyte brush-like macromolecules present in the body (Figure 65). They consist of a core protein with loosely grafted glycosaminoglycan chains, which are long, linear carbohydrate polymers that are negatively charged under physiological conditions.162–165,339 The first example of synthetic substitutes for proteoglycans has been provided by Lienkamp et al.338,340 using cylindrical polyelectrolyte brushes from poly(styrenesulfonate) (PSS). Cylindrical polyelectrolyte brushes from PSS were synthesized by polymer analogous hydrolysis from the corresponding dodecyl and ethyl ester brushes. It has been found that the aggregation behavior, size, and shape of the aggregates in solution depend on the side-chain length and the degree of saponification. The end-functionalized PSS polyelectrolyte brushes with a positively charged linker were synthesized and their complexation behavior toward negatively charged latex particles was investigated.

To investigate the effect of grafting density, Lee et al.341 prepared a series of water-soluble loosely grafted PAA brushes with four different grafting densities by the ‘grafting from’ approach using ATRP. AFM was used to study the conformation of adsorbed brushes as a function of pH. As shown in Figure 66, the adsorbed molecules undergo a globule-to-extended conformational transition as the solution is changed from acidic to basic. This transition was monitored on a mica surface by imaging individual molecules with AFM. The conformational behavior was compared with 100% grafted PAA brushes. Unlike the loosely grafted brushes, the 100% grafted molecules remained fully extended in a broad range of pH values (pH 2–9) due to steric repulsion between the densely grafted side chains, which is strongly enhanced upon adsorption to a substrate.

PDMAEMA is a unique stimuli-responsive polymer since it responds to temperature and also to pH in aqueous solution. It can also be permanently quaternized and converted to zwitterionic structures (via reaction with propanesultone), forming materials with UCST properties, as described in the previous section. Xu et al.115 prepared molecular brushes with PDMAEMA side chains and the corresponding quaternized analog from a PBIEM backbone, and studied pH response of these polymers. As expected, the structural changes were induced by variation of pH, ranging from 2 to 10. At pH 7, the PDMAEMA brushes formed worm-like structures that can be quite curved. At pH 2, most of the brushes are protonated and ionized, showing more stretched morphologies. More remarkably, at pH 10, the brushes are strongly contracted, with an average length around 110 nm, which is attributed to a collapse of the nonionized PDMAEMA side chains. pH-responsive PDMAEMA brushes were also synthesized from a conductive PT backbone by Wang et al.127 They observed conformational transitions of PT-g-PDMAEMA with a change in pH, which contributed to spectral shifts. In dilute aqueous solution, the absorption and fluorescence spectra of the polymer brush show sensitive and reversible pH responses. As shown schematically in Figure 67, the polymer brush forms a more extended conformation with a decrease in pH from 8 to 2. Protonation of the Me2N groups and increased repulsive interactions among the PDMAEMA side chains drive the redshift of the absorption and fluorescence spectra of the PT backbone.

Hydrogen Ion Concentration (pH)

The significant influence of the hydrogen ion concentration on the properties of water-based drilling fluids has long been recognized and has been the subject of numerous studies. Hydrogen ion concentration is more conveniently expressed as pH, which is the logarithm of the reciprocal of the hydrogen ion concentration in gram moles per liter. Thus, in a neutral solution the hydrogen ion (H+) and the hydroxyl ion (OH−) concentrations are equal, and each is equal to 10−7. A pH of 7 is neutral. A decrease in pH below 7 shows an increase in acidity (hydrogen ions), while an increase in pH above 7 shows an increase in alkalinity (hydroxyl ions). Each pH unit represents a 10-fold change in concentration.

Two methods for the measurement of pH are in common use: (1) a colorimetric method using paper test strips impregnated with indicators; and (2) an electrometric method using a glass electrode instrument.

Colorimetric method. Paper test strips impregnated with organic dyes, which develop colors characteristic of the pH of the liquid with which they come in contact, afford a simple and convenient method of pH measurement. The rolls of indicator paper are taken from a dispenser that has the reference comparison colors mounted on its sides. Test papers are available in both a wide-range type, which permits estimation of pH to 0.5 units, and a narrow-range type, which permits estimation to 0.2 units of pH. The test is made by placing a strip of the paper on the surface of the mud (or filtrate), allowing it to remain until the color has stabilized (usually <30 s), and comparing the color of the paper with the color standards. High concentrations of salt in the sample may alter the color developed by the dyes and cause the estimate of pH to be unreliable.

Glass electrode pH meter. When a thin membrane of glass separates two solutions of differing hydrogen ion concentrations, an electrical potential difference develops that can be amplified and measured. The pH meter consists of (1) a glass electrode made of a thin-walled bulb of special glass within which is sealed a suitable electrolyte and electrode; (2) the reference electrode, a saturated calomel cell; (3) a means of amplifying the potential difference between the external liquid (mud) and the glass electrode; and (4) a meter reading directly in pH units. Provision is made for calibrating with standard buffer solutions and for compensating for variations in temperature. A special glass electrode (less affected by sodium ions) should be used in measuring the pH of solutions containing high concentrations of sodium ions (high salinity or very high pH).

Intracellular pH During Respiratory Alkalosis

“Whole-body” intracellular hydrogen-ion concentration, as assessed by the DMO (5,5-dimethyl-2,4-oxazolidinedione) method, has been found to fall in parallel with extracellular hydrogen-ion concentration when healthy human subjects hyperventilate voluntarily to achieve a PaCO2 of 15–20 mm Hg. Similar results have been obtained from studies in dog and rat muscle and rat brain. On the other hand, 31P-nuclear magnetic resonance (NMR) spectroscopy has revealed much smaller changes in canine heart intracellular pH in response to acute hypocapnia as compared with the extracellular compartment (49). The response of intracellular acidity to chronic hypocapnia has not been studied.

7.38 pH Value or Hydrogen Ion

The pH value, or hydrogen ion concentration, determines the acidity of a water. It is one of the most important determinations in water chemistry as many of the processes involved in water treatment are pH dependent. Pure water is very slightly ionized into positive hydrogen (H+) ions and negative hydroxyl (OH−) ions. In very general terms a solution is said to be neutral when the numbers of hydrogen ions and hydroxyl ions are equal, each corresponding to an approximate concentration of 10−7 moles/l. This neutral point is temperature dependent and occurs at pH 7.0 at 25°C. When the concentration of hydrogen ions exceeds that of the hydroxyl ions (i.e. at pH values less than 7.0) the water has acidic characteristics. Conversely, when there is an excess of hydroxyl ions (i.e. the pH value is greater than 7.0) the water has basic characteristics and is described as being on the alkaline side of neutrality.

The pH value of unpolluted water is mainly determined by the inter-relationship between free carbon dioxide and the amounts of carbonate and bicarbonate present (Section 10.41). The pH values of most natural waters are in the range 4–9, with soft acidic waters from moorland areas generally having lower pH values and hard waters which have percolated through chalk or limestone generally having higher pH values.

Most water treatment processes, but particularly clarification and disinfection, require careful pH control to optimize the efficacy of the process fully. The pH of the water entering distribution must also be controlled to minimize the corrosion potential of the water (Section 7.21).

Publisher Summary

This chapter discusses the measurement of hydrogen ion concentration and carbon dioxide in clinical aspects. pH is a measure of the hydrogen ion activity in a liquid. Hydrogen ion activity is not exactly the same as hydrogen ion concentration (H+), but for practical purposes in the clinical situation, these may be regarded as equivalent. This chapter illustrates a (H+) electrode assembly. The (H+) electrode is an example of an ion-selective electrode, and it depends for its operation on an hydrogen-ion sensitive glass at its tip. A potential develops across this glass, which depends on the difference of (H+) across it. The (H+) within the (H+) electrode is maintained at a constant value by a buffer solution so that the potential across the glass is dependent on the (H+) in the blood sample in the channel.

4.10.6 The pH and p(x) Concept

In 1909 Sorensen proposed to express the hydrogen-ion concentration in an aqueous solution in terms of its negative logarithm and designated such values as pH+. His symbol has been superseded by the simple designation pH. The terms may be represented by

(4.27)pH=−logH+orpH=log1H+

With water and in the absence of foreign materials, activity equals molar concentration and [H+] equals [OH−] as required by electroneutrality, and the product at 25°C equals Kw or 10–14. These conditions mean that {H+} = {OH−} = 10− 7, and the pH equals 7, which is considered the “neutral” pH for water. The pH scale is usually represented as ranging from 0 to 14. Values of pH lower than 7 indicate that the hydrogen ion concentration is greater than the hydroxide ion concentration, and the aqueous solution is termed acidic. The opposite condition is implied when the pH exceeds 7, and the aqueous solution is termed basic.

pH is the chemical property that affects various chemical processes in soils. In general, potentially toxic metals cations are most mobile under acid conditions and increasing the pH by liming usually reduces their bioavailability. However molybdate anions become more available with increasing pH (Alloway, 1995).

The method of expressing hydrogen ion activity or concentration as pH is also useful for expressing other small numbers such as the concentration of other ions or ionization constants for solutions of weak acids and bases. For this purpose, the p(x) notation is used, with p(x) defined as

(4.28)px=−log10x=log10 1x

Here the quantity x may be the concentration of a given chemical species, an equilibrium constant, or the like. Thus, just as pH is the negative logarithm of the hydrogen ion activity, pOH signifies the negative logarithm of the hydroxide ion activity, and the pKw the negative logarithm of the ionization constant for water. From the mass reaction equation for water,

(4.29)H+OH−=Kw

It follows that

(4.30)−logH+−logOH−=−logKw

and that

(4.31)pH+pOH=pKw

Since Kw = 1 × 10− 14 at 25°C, it follows that at this temperature pKw = 14.

For weak acids and bases, pKA is the negative logarithm of the ionization constant for weak acids, and pKB is the negative logarithm of the ionization constant for weak bases. The ionization constants and pKA and pKB values for several weak acids and bases of interest to geoenvironmental engineering are listed in Tables 4.11 and 4.12 (Sawyer et al., 1967).

Table 4.11. Typical Ionization Constants for Weak Acids at 25°C

Adapted from Sawyer, C.N., McCarty, P.L., Parkin, G.F., 1967. Chemistry for Environmental Engineering. McGraw Hill, Inc., New York, 658p.

Table 4.12. Typical Ionization Constants for Weak Bases and Salts of Weak Acids at 25°C

Adapted from Sawyer, C.N., McCarty, P.L., Parkin, G.F., 1967. Chemistry for Environmental Engineering. McGraw Hill, Inc., New York, 658p.