pH

In chemistry, pH (/piːˈeɪtʃ/) is a scale used to specify how acidic or basic a water-based solution is. Acidic solutions have a lower pH, basic solutions have a higher pH. At room temperature, pure water is neither acidic nor basic and has a pH of 7.

The scale is logarithmic. It is approximately the negative of the base 10 logarithm of the molar concentration (measured in units of moles per liter) of hydrogen ions. More precisely it is the negative of the base 10 logarithm of the activity of the hydrogen ion.[1] At 25 °C, solutions with a pH less than 7 are acidic and solutions with a pH greater than 7 are basic. The neutral value of the pH depends on the temperature, being lower than 7 if the temperature increases. Pure water is neutral (pH 7) at 25 °C. Contrary to popular belief, the pH value can be less than 0 or greater than 14 for very strong acids and bases respectively.[2]

Measurements of pH are important in agronomy, medicine, chemistry, water treatment, and many other applications.

The pH scale is traceable to a set of standard solutions whose pH is established by international agreement.[3] Primary pH standard values are determined using a concentration cell with transference, by measuring the potential difference between a hydrogen electrode and a standard electrode such as the silver chloride electrode. The pH of aqueous solutions can be measured with a glass electrode and a pH meter, or an indicator.

There are three current theories used to describe acid–base reactions: Arrhenius, Bronsted-Lowry and Lewis when determining pH.

216 pH Scale-01
pH values of some common substances

History

The concept of pH was first introduced by the Danish chemist Søren Peder Lauritz Sørensen at the Carlsberg Laboratory in 1909[4] and revised to the modern pH in 1924 to accommodate definitions and measurements in terms of electrochemical cells. In the first papers, the notation had the "H" as a subscript to the lowercase "p", as so: pH.

The exact meaning of the "p" in "pH" is disputed, but according to the Carlsberg Foundation, pH stands for "power of hydrogen".[5] It has also been suggested that the "p" stands for the German Potenz (meaning "power"), others refer to French puissance (also meaning "power", based on the fact that the Carlsberg Laboratory was French-speaking). Another suggestion is that the "p" stands for the Latin terms pondus hydrogenii (quantity of hydrogen), potentia hydrogenii (capacity of hydrogen), or potential hydrogen. It is also suggested that Sørensen used the letters "p" and "q" (commonly paired letters in mathematics) simply to label the test solution (p) and the reference solution (q).[6] Currently in chemistry, the p stands for "decimal cologarithm of", and is also used in the term pKa, used for acid dissociation constants.[7]

Bacteriologist Alice C. Evans, famed for her work's influence on dairying and food safety, credited William Mansfield Clark and colleagues (of whom she was one) with developing pH measuring methods in the 1910s, which had a wide influence on laboratory and industrial use thereafter. In her memoir, she does not mention how much, or how little, Clark and colleagues knew about Sørensen's work a few years prior.[8]:10 She said:

In these studies [of bacterial metabolism] Dr. Clark's attention was directed to the effect of acid on the growth of bacteria. He found that it is the intensity of the acid in terms of hydrogen-ion concentration that affects their growth. But existing methods of measuring acidity determined the quantity, not the intensity, of the acid. Next, with his collaborators, Dr. Clark developed accurate methods for measuring hydrogen-ion concentration. These methods replaced the inaccurate titration method of determining acid content in use in biologic laboratories throughout the world. Also they were found to be applicable in many industrial and other processes in which they came into wide usage.[8]:10

The first electronic method for measuring pH was invented by Arnold Orville Beckman, a professor at California Institute of Technology in 1934.[9] It was in response to local citrus grower Sunkist that wanted a better method for quickly testing the pH of lemons they were picking from their nearby orchards.[10]

Definition and measurement

pH

pH is defined as the decimal logarithm of the reciprocal of the hydrogen ion activity, aH+, in a solution.[3]

For example, for a solution with a hydrogen ion activity of 5×10−6 (at that level, this is essentially the number of moles of hydrogen ions per liter of solution) we get 1/(5×10−6) = 2×105, thus such a solution has a pH of log10(2×105) = 5.3. For a commonplace example based on the facts that the masses of a mole of water, a mole of hydrogen ions, and a mole of hydroxide ions are respectively 18 g, 1 g, and 17 g, a quantity of 107 moles of pure (pH 7) water, or 180 tonnes (18×107 g), contains close to 1 g of dissociated hydrogen ions (or rather 19 g of H3O+ hydronium ions) and 17 g of hydroxide ions.

Note that pH depends on temperature. For instance at 0 °C the pH of pure water is 7.47. At 25 °C it's 7.00, and at 100 °C it's 6.14.

This definition was adopted because ion-selective electrodes, which are used to measure pH, respond to activity. Ideally, electrode potential, E, follows the Nernst equation, which, for the hydrogen ion can be written as

where E is a measured potential, E0 is the standard electrode potential, R is the gas constant, T is the temperature in kelvins, F is the Faraday constant. For H+ number of electrons transferred is one. It follows that electrode potential is proportional to pH when pH is defined in terms of activity. Precise measurement of pH is presented in International Standard ISO 31-8 as follows:[11] A galvanic cell is set up to measure the electromotive force (e.m.f.) between a reference electrode and an electrode sensitive to the hydrogen ion activity when they are both immersed in the same aqueous solution. The reference electrode may be a silver chloride electrode or a calomel electrode. The hydrogen-ion selective electrode is a standard hydrogen electrode.

Reference electrode | concentrated solution of KCl || test solution | H2 | Pt

Firstly, the cell is filled with a solution of known hydrogen ion activity and the emf, ES, is measured. Then the emf, EX, of the same cell containing the solution of unknown pH is measured.

The difference between the two measured emf values is proportional to pH. This method of calibration avoids the need to know the standard electrode potential. The proportionality constant, 1/z is ideally equal to the "Nernstian slope".

To apply this process in practice, a glass electrode is used rather than the cumbersome hydrogen electrode. A combined glass electrode has an in-built reference electrode. It is calibrated against buffer solutions of known hydrogen ion activity. IUPAC has proposed the use of a set of buffer solutions of known H+ activity.[3] Two or more buffer solutions are used in order to accommodate the fact that the "slope" may differ slightly from ideal. To implement this approach to calibration, the electrode is first immersed in a standard solution and the reading on a pH meter is adjusted to be equal to the standard buffer's value. The reading from a second standard buffer solution is then adjusted, using the "slope" control, to be equal to the pH for that solution. Further details, are given in the IUPAC recommendations.[3] When more than two buffer solutions are used the electrode is calibrated by fitting observed pH values to a straight line with respect to standard buffer values. Commercial standard buffer solutions usually come with information on the value at 25 °C and a correction factor to be applied for other temperatures.

The pH scale is logarithmic and therefore pH is a dimensionless quantity.

p[H]

This was the original definition of Sørensen,[5] which was superseded in favor of pH in 1909. [H] is the concentration of hydrogen ions, denoted [H+] in modern chemistry, which appears to have units of concentration. More correctly, the thermodynamic activity of H+ in dilute solution should be replaced by [H+]/c0, where the standard state concentration c0 = 1 mol/L. This ratio is a pure number whose logarithm can be defined.

However, it is possible to measure the concentration of hydrogen ions directly, if the electrode is calibrated in terms of hydrogen ion concentrations. One way to do this, which has been used extensively, is to titrate a solution of known concentration of a strong acid with a solution of known concentration of strong alkaline in the presence of a relatively high concentration of background electrolyte. Since the concentrations of acid and alkaline are known, it is easy to calculate the concentration of hydrogen ions so that the measured potential can be correlated with concentrations. The calibration is usually carried out using a Gran plot.[12] The calibration yields a value for the standard electrode potential, E0, and a slope factor, f, so that the Nernst equation in the form

can be used to derive hydrogen ion concentrations from experimental measurements of E. The slope factor, f, is usually slightly less than one. A slope factor of less than 0.95 indicates that the electrode is not functioning correctly. The presence of background electrolyte ensures that the hydrogen ion activity coefficient is effectively constant during the titration. As it is constant, its value can be set to one by defining the standard state as being the solution containing the background electrolyte. Thus, the effect of using this procedure is to make activity equal to the numerical value of concentration.

The glass electrode (and other ion selective electrodes) should be calibrated in a medium similar to the one being investigated. For instance, if one wishes to measure the pH of a seawater sample, the electrode should be calibrated in a solution resembling seawater in its chemical composition, as detailed below.

The difference between p[H] and pH is quite small. It has been stated[13] that pH = p[H] + 0.04. It is common practice to use the term "pH" for both types of measurement.

pH indicators

Universal indicator paper
Chart showing the variation of color of universal indicator paper with pH

Indicators may be used to measure pH, by making use of the fact that their color changes with pH. Visual comparison of the color of a test solution with a standard color chart provides a means to measure pH accurate to the nearest whole number. More precise measurements are possible if the color is measured spectrophotometrically, using a colorimeter or spectrophotometer. Universal indicator consists of a mixture of indicators such that there is a continuous color change from about pH 2 to pH 10. Universal indicator paper is made from absorbent paper that has been impregnated with universal indicator. Another method of measuring pH is using an electronic pH meter.

pOH

PHscalenolang
Relation between p[OH] and p[H] (red = acidic region, blue = basic region)

pOH is sometimes used as a measure of the concentration of hydroxide ions. OH. pOH values are derived from pH measurements. The concentration of hydroxide ions in water is related to the concentration of hydrogen ions by

where KW is the self-ionisation constant of water. Taking logarithms

So, at room temperature, pOH ≈ 14 − pH. However this relationship is not strictly valid in other circumstances, such as in measurements of soil alkalinity.

Extremes of pH

Measurement of pH below about 2.5 (ca. 0.003 mol dm−3 acid) and above about 10.5 (ca. 0.0003 mol dm−3 alkaline) requires special procedures because, when using the glass electrode, the Nernst law breaks down under those conditions. Various factors contribute to this. It cannot be assumed that liquid junction potentials are independent of pH.[14] Also, extreme pH implies that the solution is concentrated, so electrode potentials are affected by ionic strength variation. At high pH the glass electrode may be affected by "alkaline error", because the electrode becomes sensitive to the concentration of cations such as Na+ and K+ in the solution.[15] Specially constructed electrodes are available which partly overcome these problems.

Runoff from mines or mine tailings can produce some very low pH values.[16]

Non-aqueous solutions

Hydrogen ion concentrations (activities) can be measured in non-aqueous solvents. pH values based on these measurements belong to a different scale from aqueous pH values, because activities relate to different standard states. Hydrogen ion activity, aH+, can be defined[17][18] as:

where μH+ is the chemical potential of the hydrogen ion, is its chemical potential in the chosen standard state, R is the gas constant and T is the thermodynamic temperature. Therefore, pH values on the different scales cannot be compared directly due to different solvated proton ions such as lyonium ions, requiring an intersolvent scale which involves the transfer activity coefficient of hydronium/lyonium ion.

pH is an example of an acidity function. Other acidity functions can be defined. For example, the Hammett acidity function, H0, has been developed in connection with superacids.

Unified absolute pH scale

The concept of "unified pH scale" has been developed on the basis of the absolute chemical potential of the proton. This model uses the Lewis acid–base definition. This scale applies to liquids, gases and even solids.[19] In 2010, a new "unified absolute pH scale" has been proposed that would allow various pH ranges across different solutions to use a common proton reference standard.[20]

Applications

Pure water is neutral. When an acid is dissolved in water, the pH will be less than 7 (25 °C). When a base, or alkali, is dissolved in water, the pH will be greater than 7. A solution of a strong acid, such as hydrochloric acid, at concentration 1 mol dm−3 has a pH of 0. A solution of a strong alkali, such as sodium hydroxide, at concentration 1 mol dm−3, has a pH of 14. Thus, measured pH values will lie mostly in the range 0 to 14, though negative pH values and values above 14 are entirely possible. Since pH is a logarithmic scale, a difference of one pH unit is equivalent to a tenfold difference in hydrogen ion concentration.

The pH of neutrality is not exactly 7 (25 °C), although this is a good approximation in most cases. Neutrality is defined as the condition where [H+] = [OH] (or the activities are equal). Since self-ionization of water holds the product of these concentration [H+]×[OH] = Kw, it can be seen that at neutrality [H+] = [OH] = Kw, or pH = pKw/2. pKw is approximately 14 but depends on ionic strength and temperature, and so the pH of neutrality does also. Pure water and a solution of NaCl in pure water are both neutral, since dissociation of water produces equal numbers of both ions. However the pH of the neutral NaCl solution will be slightly different from that of neutral pure water because the hydrogen and hydroxide ions' activity is dependent on ionic strength, so Kw varies with ionic strength.

If pure water is exposed to air it becomes mildly acidic. This is because water absorbs carbon dioxide from the air, which is then slowly converted into bicarbonate and hydrogen ions (essentially creating carbonic acid).

pH in soil

Classification of soil pH ranges

The United States Department of Agriculture Natural Resources Conservation Service, formerly Soil Conservation Service classifies soil pH ranges as follows: [21]

Soil pH effect on nutrient availability
Nutritional elements availability within soil varies with pH. Light blue color represents the ideal range for most plants.
Denomination pH range
Ultra acidic < 3.5
Extremely acidic 3.5–4.4
Very strongly acidic 4.5–5.0
Strongly acidic 5.1–5.5
Moderately acidic 5.6–6.0
Slightly acidic 6.1–6.5
Neutral 6.6–7.3
Slightly alkaline 7.4–7.8
Moderately alkaline 7.9–8.4
Strongly alkaline 8.5–9.0
Very strongly alkaline > 9.0

pH in nature

Lemon
Lemon juice tastes sour because it contains 5% to 6% citric acid and has a pH of 2.2. (high acidity)

pH-dependent plant pigments that can be used as pH indicators occur in many plants, including hibiscus, red cabbage (anthocyanin) and red wine. The juice of citrus fruits is acidic mainly because it contains citric acid. Other carboxylic acids occur in many living systems. For example, lactic acid is produced by muscle activity. The state of protonation of phosphate derivatives, such as ATP, is pH-dependent. The functioning of the oxygen-transport enzyme hemoglobin is affected by pH in a process known as the Root effect.

Seawater

The pH of seawater is typically limited to a range between 7.5 and 8.4.[22] It plays an important role in the ocean's carbon cycle, and there is evidence of ongoing ocean acidification caused by carbon dioxide emissions.[23] However, pH measurement is complicated by the chemical properties of seawater, and several distinct pH scales exist in chemical oceanography.[24]

As part of its operational definition of the pH scale, the IUPAC defines a series of buffer solutions across a range of pH values (often denoted with NBS or NIST designation). These solutions have a relatively low ionic strength (≈0.1) compared to that of seawater (≈0.7), and, as a consequence, are not recommended for use in characterizing the pH of seawater, since the ionic strength differences cause changes in electrode potential. To resolve this problem, an alternative series of buffers based on artificial seawater was developed.[25] This new series resolves the problem of ionic strength differences between samples and the buffers, and the new pH scale is referred to as the 'total scale', often denoted as pHT. The total scale was defined using a medium containing sulfate ions. These ions experience protonation, H+ + SO2−
4
⇌ HSO
4
, such that the total scale includes the effect of both protons (free hydrogen ions) and hydrogen sulfate ions:

[H+]T = [H+]F + [HSO
4
]

An alternative scale, the 'free scale', often denoted 'pHF', omits this consideration and focuses solely on [H+]F, in principle making it a simpler representation of hydrogen ion concentration. Only [H+]T can be determined,[26] therefore [H+]F must be estimated using the [SO2−
4
] and the stability constant of HSO
4
, K*
S
:

[H+]F = [H+]T − [HSO
4
] = [H+]T ( 1 + [SO2−
4
] / K*
S
)−1

However, it is difficult to estimate K*
S
in seawater, limiting the utility of the otherwise more straightforward free scale.

Another scale, known as the 'seawater scale', often denoted 'pHSWS', takes account of a further protonation relationship between hydrogen ions and fluoride ions, H+ + F ⇌ HF. Resulting in the following expression for [H+]SWS:

[H+]SWS = [H+]F + [HSO
4
] + [HF]

However, the advantage of considering this additional complexity is dependent upon the abundance of fluoride in the medium. In seawater, for instance, sulfate ions occur at much greater concentrations (>400 times) than those of fluoride. As a consequence, for most practical purposes, the difference between the total and seawater scales is very small.

The following three equations summarise the three scales of pH:

pHF = − log [H+]F
pHT = − log ( [H+]F + [HSO
4
] ) = − log [H+]T
pHSWS = − log ( [H+]F + [HSO
4
] + [HF] ) = − log [H+]SWS

In practical terms, the three seawater pH scales differ in their values by up to 0.12 pH units, differences that are much larger than the accuracy of pH measurements typically required, in particular, in relation to the ocean's carbonate system.[24] Since it omits consideration of sulfate and fluoride ions, the free scale is significantly different from both the total and seawater scales. Because of the relative unimportance of the fluoride ion, the total and seawater scales differ only very slightly.

Living systems

pH in living systems[27]
Compartment pH
Gastric acid 1.5-3.5[28]
Lysosomes 4.5
Human skin 4.7[29]
Granules of chromaffin cells 5.5
Urine 6.0
Cytosol 7.2
Blood (natural pH) 7.34–7.45
Cerebrospinal fluid (CSF) 7.5
Mitochondrial matrix 7.5
Pancreas secretions 8.1

The pH of different cellular compartments, body fluids, and organs is usually tightly regulated in a process called acid-base homeostasis. The most common disorder in acid-base homeostasis is acidosis, which means an acid overload in the body, generally defined by pH falling below 7.35. Alkalosis is the opposite condition, with blood pH being excessively high.

The pH of blood is usually slightly basic with a value of pH 7.365. This value is often referred to as physiological pH in biology and medicine. Plaque can create a local acidic environment that can result in tooth decay by demineralization. Enzymes and other proteins have an optimum pH range and can become inactivated or denatured outside this range.

Calculations of pH

The calculation of the pH of a solution containing acids and/or bases is an example of a chemical speciation calculation, that is, a mathematical procedure for calculating the concentrations of all chemical species that are present in the solution. The complexity of the procedure depends on the nature of the solution. For strong acids and bases no calculations are necessary except in extreme situations. The pH of a solution containing a weak acid requires the solution of a quadratic equation. The pH of a solution containing a weak base may require the solution of a cubic equation. The general case requires the solution of a set of non-linear simultaneous equations.

A complicating factor is that water itself is a weak acid and a weak base (see amphoterism). It dissociates according to the equilibrium

with a dissociation constant, Kw defined as

where [H+] stands for the concentration of the aqueous hydronium ion and [OH] represents the concentration of the hydroxide ion. This equilibrium needs to be taken into account at high pH and when the solute concentration is extremely low.

Strong acids and bases

Strong acids and bases are compounds that, for practical purposes, are completely dissociated in water. Under normal circumstances this means that the concentration of hydrogen ions in acidic solution can be taken to be equal to the concentration of the acid. The pH is then equal to minus the logarithm of the concentration value. Hydrochloric acid (HCl) is an example of a strong acid. The pH of a 0.01M solution of HCl is equal to −log10(0.01), that is, pH = 2. Sodium hydroxide, NaOH, is an example of a strong base. The p[OH] value of a 0.01M solution of NaOH is equal to −log10(0.01), that is, p[OH] = 2. From the definition of p[OH] above, this means that the pH is equal to about 12. For solutions of sodium hydroxide at higher concentrations the self-ionization equilibrium must be taken into account.

Self-ionization must also be considered when concentrations are extremely low. Consider, for example, a solution of hydrochloric acid at a concentration of 5×10−8M. The simple procedure given above would suggest that it has a pH of 7.3. This is clearly wrong as an acid solution should have a pH of less than 7. Treating the system as a mixture of hydrochloric acid and the amphoteric substance water, a pH of 6.89 results.[30]

Weak acids and bases

A weak acid or the conjugate acid of a weak base can be treated using the same formalism.

First, an acid dissociation constant is defined as follows. Electrical charges are omitted from subsequent equations for the sake of generality

and its value is assumed to have been determined by experiment. This being so, there are three unknown concentrations, [HA], [H+] and [A] to determine by calculation. Two additional equations are needed. One way to provide them is to apply the law of mass conservation in terms of the two "reagents" H and A.

C stands for analytical concentration. In some texts, one mass balance equation is replaced by an equation of charge balance. This is satisfactory for simple cases like this one, but is more difficult to apply to more complicated cases as those below. Together with the equation defining Ka, there are now three equations in three unknowns. When an acid is dissolved in water CA = CH = Ca, the concentration of the acid, so [A] = [H]. After some further algebraic manipulation an equation in the hydrogen ion concentration may be obtained.

Solution of this quadratic equation gives the hydrogen ion concentration and hence p[H] or, more loosely, pH. This procedure is illustrated in an ICE table which can also be used to calculate the pH when some additional (strong) acid or alkaline has been added to the system, that is, when CA ≠ CH.

For example, what is the pH of a 0.01M solution of benzoic acid, pKa = 4.19?

  • Step 1:
  • Step 2: Set up the quadratic equation.
  • Step 3: Solve the quadratic equation.

For alkaline solutions an additional term is added to the mass-balance equation for hydrogen. Since addition of hydroxide reduces the hydrogen ion concentration, and the hydroxide ion concentration is constrained by the self-ionization equilibrium to be equal to

In this case the resulting equation in [H] is a cubic equation.

General method

Some systems, such as with polyprotic acids, are amenable to spreadsheet calculations.[31] With three or more reagents or when many complexes are formed with general formulae such as ApBqHr,the following general method can be used to calculate the pH of a solution. For example, with three reagents, each equilibrium is characterized by an equilibrium constant, β.

Next, write down the mass-balance equations for each reagent:

Note that there are no approximations involved in these equations, except that each stability constant is defined as a quotient of concentrations, not activities. Much more complicated expressions are required if activities are to be used.

There are 3 non-linear simultaneous equations in the three unknowns, [A], [B] and [H]. Because the equations are non-linear, and because concentrations may range over many powers of 10, the solution of these equations is not straightforward. However, many computer programs are available which can be used to perform these calculations. There may be more than three reagents. The calculation of hydrogen ion concentrations, using this formalism, is a key element in the determination of equilibrium constants by potentiometric titration.

See also

References

  1. ^ Bates, Roger G. Determination of pH: theory and practice. Wiley, 1973.
  2. ^ Lim, Kieran F. (2006). "Negative pH Does Exist". Journal of Chemical Education. 83 (10): 1465. Bibcode:2006JChEd..83.1465L. doi:10.1021/ed083p1465.
  3. ^ a b c d Covington, A. K.; Bates, R. G.; Durst, R. A. (1985). "Definitions of pH scales, standard reference values, measurement of pH, and related terminology" (PDF). Pure Appl. Chem. 57 (3): 531–542. doi:10.1351/pac198557030531. Archived (PDF) from the original on 24 September 2007.
  4. ^ Sørensen, S. P. L. (1909). "Über die Messung und die Bedeutung der Wasserstoffionenkonzentration bei enzymatischen Prozessen". Biochem. Zeitschr. 21: 131–304. Two other publications appeared in 1909 one in French and one in Danish
  5. ^ a b "Carlsberg Group Company History Page". Carlsberggroup.com. Archived from the original on 18 January 2014. Retrieved 7 May 2013.
  6. ^ Myers, Rollie J. (2010). "One-Hundred Years of pH". Journal of Chemical Education. 87 (1): 30–32. Bibcode:2010JChEd..87...30M. doi:10.1021/ed800002c.
  7. ^ Nørby, Jens (2000). "The origin and the meaning of the little p in pH". Trends in Biochemical Sciences. 25 (1): 36–37. doi:10.1016/S0968-0004(99)01517-0. PMID 10637613.
  8. ^ a b Evans, Alice C. (1963). "Memoirs" (PDF). NIH Office of History. National Institutes of Health Office of History. Retrieved 2018-03-27.
  9. ^ "ORIGINS: BIRTH OF THE PH METER". CalTech Engineering & Science Magazine. Retrieved 11 March 2018.
  10. ^ Tetrault, Sharon (June 2002). "The Beckmans". Orange Coast. Orange Coast Magazine. Retrieved 11 March 2018.
  11. ^ Quantities and units – Part 8: Physical chemistry and molecular physics, Annex C (normative): pH. International Organization for Standardization, 1992.
  12. ^ Rossotti, F.J.C.; Rossotti, H. (1965). "Potentiometric titrations using Gran plots: A textbook omission". J. Chem. Educ. 42 (7): 375–378. Bibcode:1965JChEd..42..375R. doi:10.1021/ed042p375.
  13. ^ Mendham, J.; Denney, R. C.; Barnes, J. D.; Thomas, M. J. K. (2000), Vogel's Quantitative Chemical Analysis (6th ed.), New York: Prentice Hall, ISBN 0-582-22628-7, Section 13.23, "Determination of pH"
  14. ^ Feldman, Isaac (1956). "Use and Abuse of pH measurements". Analytical Chemistry. 28 (12): 1859–1866. doi:10.1021/ac60120a014.
  15. ^ Mendham, J.; Denney, R. C.; Barnes, J. D.; Thomas, M. J. K. (2000), Vogel's Quantitative Chemical Analysis (6th ed.), New York: Prentice Hall, ISBN 0-582-22628-7, Section 13.19 The glass electrode
  16. ^ Nordstrom, D. Kirk; Alpers, Charles N. (March 1999). "Negative pH, efflorescent mineralogy, and consequences for environmental restoration at the Iron Mountain Superfund site, California". Proceedings of the National Academy of Sciences of the United States of America. 96 (7): 3455–62. Bibcode:1999PNAS...96.3455N. doi:10.1073/pnas.96.7.3455. PMC 34288. PMID 10097057.
  17. ^ IUPAC, Compendium of Chemical Terminology, 2nd ed. (the "Gold Book") (1997). Online corrected version:  (2006–) "activity (relative activity), a". doi:10.1351/goldbook.A00115
  18. ^ International Union of Pure and Applied Chemistry (1993). Quantities, Units and Symbols in Physical Chemistry, 2nd edition, Oxford: Blackwell Science. ISBN 0-632-03583-8. pp. 49–50. Electronic version.
  19. ^ Himmel, D.; Goll, S. K.; Leito, I.; Krossing, I. (2010). "A Unified pH Scale for all Phases". Angew. Chem. Int. Ed. 49 (38): 6885–6888. doi:10.1002/anie.201000252. PMID 20715223.
  20. ^ Himmel, Daniel; Goll, Sascha K.; Leito, Ivo; Krossing, Ingo (2010-08-16). "A Unified pH Scale for All Phases". Angewandte Chemie International Edition. 49 (38): 6885–6888. doi:10.1002/anie.201000252. ISSN 1433-7851. PMID 20715223.
  21. ^ Soil Survey Division Staff. "Soil survey manual.1993. Chapter 3, selected chemical properties". Soil Conservation Service. U.S. Department of Agriculture Handbook 18. Archived from the original on 14 May 2011. Retrieved 12 March 2011.
  22. ^ Chester, Jickells, Roy, Tim (2012). Marine Geochemistry. Blackwell Publishing. ISBN 978-1-118-34907-6.
  23. ^ Royal Society (2005). Ocean acidification due to increasing atmospheric carbon dioxide (PDF). ISBN 978-0-85403-617-2. Archived (PDF) from the original on 16 July 2010.
  24. ^ a b Zeebe, R. E. and Wolf-Gladrow, D. (2001) CO2 in seawater: equilibrium, kinetics, isotopes, Elsevier Science B.V., Amsterdam, Netherlands ISBN 0-444-50946-1
  25. ^ Hansson, I. (1973). "A new set of pH-scales and standard buffers for seawater". Deep-Sea Research. 20 (5): 479–491. Bibcode:1973DSROA..20..479H. doi:10.1016/0011-7471(73)90101-0.
  26. ^ Dickson, A. G. (1984). "pH scales and proton-transfer reactions in saline media such as sea water". Geochim. Cosmochim. Acta. 48 (11): 2299–2308. Bibcode:1984GeCoA..48.2299D. doi:10.1016/0016-7037(84)90225-4.
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  29. ^ Lambers, H.; Piessens, S.; Bloem, A.; Pronk, H.; Finkel, P. (2006-10-01). "Natural skin surface pH is on average below 5, which is beneficial for its resident flora". International Journal of Cosmetic Science. 28 (5): 359–370. doi:10.1111/j.1467-2494.2006.00344.x. ISSN 1468-2494. PMID 18489300.
  30. ^ Maloney, Chris. "pH calculation of a very small concentration of a strong acid". Archived from the original on 8 July 2011. Retrieved 13 March 2011.
  31. ^ Billo, E.J. (2011). EXCEL for Chemists (3rd ed.). Wiley-VCH. ISBN 978-0-470-38123-6.

External links

Acid

An acid is a molecule or ion capable of donating a hydron (proton or hydrogen ion H+), or, alternatively, capable of forming a covalent bond with an electron pair (a Lewis acid).The first category of acids is the proton donors or Brønsted acids. In the special case of aqueous solutions, proton donors form the hydronium ion H3O+ and are known as Arrhenius acids. Brønsted and Lowry generalized the Arrhenius theory to include non-aqueous solvents. A Brønsted or Arrhenius acid usually contains a hydrogen atom bonded to a chemical structure that is still energetically favorable after loss of H+.

Aqueous Arrhenius acids have characteristic properties which provide a practical description of an acid. Acids form aqueous solutions with a sour taste, can turn blue litmus red, and react with bases and certain metals (like calcium) to form salts. The word acid is derived from the Latin acidus/acēre meaning sour. An aqueous solution of an acid has a pH less than 7 and is colloquially also referred to as 'acid' (as in 'dissolved in acid'), while the strict definition refers only to the solute. A lower pH means a higher acidity, and thus a higher concentration of positive hydrogen ions in the solution. Chemicals or substances having the property of an acid are said to be acidic.

Common aqueous acids include hydrochloric acid (a solution of hydrogen chloride which is found in gastric acid in the stomach and activates digestive enzymes), acetic acid (vinegar is a dilute aqueous solution of this liquid), sulfuric acid (used in car batteries), and citric acid (found in citrus fruits). As these examples show, acids (in the colloquial sense) can be solutions or pure substances, and can be derived from acids (in the strict sense) that are solids, liquids, or gases. Strong acids and some concentrated weak acids are corrosive, but there are exceptions such as carboranes and boric acid.

The second category of acids are Lewis acids, which form a covalent bond with an electron pair. An example is boron trifluoride (BF3), whose boron atom has a vacant orbital which can form a covalent bond by sharing a lone pair of electrons on an atom in a base, for example the nitrogen atom in ammonia (NH3). Lewis considered this as a generalization of the Brønsted definition, so that an acid is a chemical species that accepts electron pairs either directly or by releasing protons (H+) into the solution, which then accept electron pairs. However, hydrogen chloride, acetic acid, and most other Brønsted-Lowry acids cannot form a covalent bond with an electron pair and are therefore not Lewis acids. Conversely, many Lewis acids are not Arrhenius or Brønsted-Lowry acids. In modern terminology, an acid is implicitly a Brønsted acid and not a Lewis acid, since chemists almost always refer to a Lewis acid explicitly as a Lewis acid.

Acid dissociation constant

An acid dissociation constant, Ka, (also known as acidity constant, or acid-ionization constant) is a quantitative measure of the strength of an acid in solution. It is the equilibrium constant for a chemical reaction known as dissociation in the context of acid–base reactions.

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The chemical species HA, A, and H+ are said to be in equilibrium when their concentrations (written above in square brackets) do not change with the passing of time, because both forward and backward reactions are occurring at the same very fast rate. The chemical equation for acid dissociation can be written symbolically as:

where HA is a generic acid that dissociates into A, the conjugate base of the acid and a hydrogen ion, H+. It is implicit in this definition that the quotient of activity coefficients, Γ,

is a constant that can be ignored in a given set of experimental conditions.

For many practical purposes it is more convenient to discuss the logarithmic constant, pKa

The more positive the value of pKa, the smaller the extent of dissociation at any given pH (see Henderson–Hasselbalch equation)—that is, the weaker the acid. A weak acid has a pKa value in the approximate range −2 to 12 in water. Acids with a pKa value of less than about −2 are said to be strong acids; the dissociation of a strong acid is effectively complete such that concentration of the undissociated acid is too small to be measured. pKa values for strong acids can, however, be estimated by theoretical means.

Acid rain

Acid rain is a rain or any other form of precipitation that is unusually acidic, meaning that it has elevated levels of hydrogen ions (low pH). It can have harmful effects on plants, aquatic animals and infrastructure. Acid rain is caused by emissions of sulfur dioxide and nitrogen oxide, which react with the water molecules in the atmosphere to produce acids. Some governments have made efforts since the 1970s to reduce the release of sulfur dioxide and nitrogen oxide into the atmosphere with positive results. Nitrogen oxides can also be produced naturally by lightning strikes, and sulfur dioxide is produced by volcanic eruptions. Acid rain has been shown to have adverse impacts on forests, freshwaters and soils, killing insect and aquatic life-forms, causing paint to peel, corrosion of steel structures such as bridges, and weathering of stone buildings and statues as well as having impacts on human health.

Buffer solution

A buffer solution (more precisely, pH buffer or hydrogen ion buffer) is an aqueous solution consisting of a mixture of a weak acid and its conjugate base, or vice versa. Its pH changes very little when a small amount of strong acid or base is added to it. Buffer solutions are used as a means of keeping pH at a nearly constant value in a wide variety of chemical applications. In nature, there are many systems that use buffering for pH regulation. For example, the bicarbonate buffering system is used to regulate the pH of blood.

Calcium carbonate

Calcium carbonate is a chemical compound with the formula CaCO3. It is a common substance found in rocks as the minerals calcite and aragonite (most notably as limestone, which is a type of sedimentary rock consisting mainly of calcite) and is the main component of pearls and the shells of marine organisms, snails, and eggs. Calcium carbonate is the active ingredient in agricultural lime and is created when calcium ions in hard water react with carbonate ions to create limescale. It is medicinally used as a calcium supplement or as an antacid, but excessive consumption can be hazardous.

Civil Services Examination (India)

The Civil Services Examination (CSE) is a nationwide competitive examination in India conducted by the Union Public Service Commission for recruitment to various Civil Services of the Government of India, including the Indian Administrative Service (IAS), Indian Foreign Service (IFS), Indian Police Service (IPS) among others. Also simply referred as UPSC examination, it is conducted in three phases - a preliminary examination consisting of two objective-type papers (General Studies Paper I and General Studies Paper II also popularly known as Civil Service Aptitude Test or CSAT), and a main examination consisting of nine papers of conventional (essay) type, in which two papers are qualifying and only marks of seven are counted followed by a personality test (interview).

Doctor of Philosophy

A Doctor of Philosophy (PhD, Ph.D., or DPhil; Latin philosophiae doctorem or doctorem philosophiae) is the highest university degree that is conferred after a course of study by universities in most English-speaking countries. PhDs are awarded for programs across the whole breadth of academic fields. As an earned research degree, those studying for a PhD are usually required to produce original research that expands the boundaries of knowledge, normally in the form of a thesis or dissertation, and defend their work against experts in the field. The completion of a PhD is often a requirement for employment as a university professor, researcher, or scientist in many fields.

Individuals who have earned a Doctor of Philosophy degree may, in many jurisdictions, use the title Doctor (often abbreviated "Dr" or "Dr.") or, in non-English-speaking countries, variants such as "Dr. phil." with their name, although the proper etiquette associated with this usage may also be subject to the professional ethics of their own scholarly field, culture, or society. Those who teach at universities or work in academic, educational, or research fields are usually addressed by this title "professionally and socially in a salutation or conversation." Alternatively, holders may use post-nominal letters such as "Ph.D.", "PhD", or "DPhil" (depending on the awarding institution). It is, however, considered incorrect to use both the title and post-nominals at the same time.The specific requirements to earn a PhD degree vary considerably according to the country, institution, and time period, from entry-level research degrees to higher doctorates. During the studies that lead to the degree, the student is called a doctoral student or PhD student; a student who has completed all of their coursework and comprehensive examinations and is working on their thesis/dissertation is sometimes known as a doctoral candidate or PhD candidate (see: all but dissertation). A student attaining this level may be granted a Candidate of Philosophy degree at some institutions, or may be granted a master's degree en route to the doctoral degree. Sometimes this status is also colloquially known as "Ph.D. ABD", meaning "All But Dissertation."A PhD candidate must submit a project, thesis or dissertation often consisting of a body of original academic research, which is in principle worthy of publication in a peer-reviewed journal. In many countries, a candidate must defend this work before a panel of expert examiners appointed by the university. Universities sometimes award other types of doctorate besides the PhD, such as the Doctor of Musical Arts (D.M.A.) for music performers and the Doctor of Education (Ed.D.) for professional educators. In 2005 the European Universities Association defined the Salzburg Principles, ten basic principles for third-cycle degrees (doctorates) within the Bologna Process. These were followed in 2016 by the Florence Principles, seven basic principles for doctorates in the arts laid out by the European League of Institutes of the Arts, which have been endorsed by the European Association of Conservatoires, the International Association of Film and Television Schools, the International Association of Universities and Colleges of Art, Design and Media, and the Society for Artistic Research.In the context of the Doctor of Philosophy and other similarly titled degrees, the term "philosophy" does not refer to the field or academic discipline of philosophy, but is used in a broader sense in accordance with its original Greek meaning, which is "love of wisdom". In most of Europe, all fields (history, philosophy, social sciences, mathematics, and natural philosophy/sciences) other than theology, law, and medicine (the so-called professional, vocational, or technical curriculum) were traditionally known as philosophy, and in Germany and elsewhere in Europe the basic faculty of liberal arts was known as the "faculty of philosophy".

Doctorate

A doctorate (from Latin docere, "to teach") or doctor's degree (from Latin doctor, "teacher") or doctoral degree, is an academic degree awarded by universities, derived from the ancient formalism licentia docendi ("licence to teach") In most countries, it is a research degree that qualifies the holder to teach at university level in the degree's field, or to work in a specific profession. There are a variety of names for doctoral degrees; the most common is the Doctor of Philosophy (PhD), which is awarded in many different fields, ranging from the humanities to scientific disciplines.

In the United States and some other countries, there are also some types of vocational, technical, or professional degrees that are referred to as doctorates. Professional doctorates historically came about to meet the needs of practitioners in a variety of disciplines. However, the aims and means of these degrees vary greatly across disciplines, making their significance unclear.

Many universities also award honorary doctorates to individuals deemed worthy of special recognition, either for scholarly work or for other contributions to the university or to society.

Isoelectric point

The isoelectric point (pI, pH(I), IEP), is the pH at which a particular molecule carries no net electrical charge or is electrically neutral in the statistical mean. The standard nomenclature to represent the isoelectric point is pH(I), although pI is also commonly seen, and is used in this article for brevity. The net charge on the molecule is affected by pH of its surrounding environment and can become more positively or negatively charged due to the gain or loss, respectively, of protons (H+).

Surfaces naturally charge to form a double layer. In the common case when the surface charge-determining ions are H+/OH−, the net surface charge is affected by the pH of the liquid in which the solid is submerged.

The pI value can affect the solubility of a molecule at a given pH. Such molecules have minimum solubility in water or salt solutions at the pH that corresponds to their pI and often precipitate out of solution. Biological amphoteric molecules such as proteins contain both acidic and basic functional groups. Amino acids that make up proteins may be positive, negative, neutral, or polar in nature, and together give a protein its overall charge. At a pH below their pI, proteins carry a net positive charge; above their pI they carry a net negative charge. Proteins can, thus, be separated by net charge in a polyacrylamide gel using either preparative gel electrophoresis, which uses a constant pH to separate proteins or isoelectric focusing, which uses a pH gradient to separate proteins. Isoelectric focusing is also the first step in 2-D gel polyacrylamide gel electrophoresis.

In biomolecules, proteins can be separated by ion exchange chromatography. Biological proteins are made up of zwitterionic amino acid compounds; the net charge of these proteins can be positive or negative depending on the pH of the environment. The specific pI of the target protein can be used to model the process around and the compound can then be purified from the rest of the mixture. Buffers of various pH can be used for this purification process to change the pH of the environment. When a mixture containing a target protein is loaded into an ion exchanger, the stationary matrix can be either positively-charged (for mobile anions) or negatively-charged (for mobile cations). At low pH values, the net charge of most proteins in the mixture is positive - in cation exchangers, these positively-charged proteins bind to the negatively-charged matrix. At high pH values, the net charge of most proteins is negative, where they bind to the positively-charged matrix in anion exchangers. When the environment is at a pH value equal to the protein's pI, the net charge is zero, and the protein is not bound to any exchanger, and therefore, can be eluted out.

KLM

KLM Royal Dutch Airlines, legally Koninklijke Luchtvaart Maatschappij N.V. (literal translation: Royal Aviation Company, Inc.), is the flag carrier airline of the Netherlands. KLM is headquartered in Amstelveen, with its hub at nearby Amsterdam Airport Schiphol. It is part of the Air France–KLM group, and a member of the SkyTeam airline alliance. Founded in 1919, KLM is the oldest airline in the world still operating under its original name and had 35,488 employees and a fleet of 119 as of 2015. KLM operates scheduled passenger and cargo services to 145 destinations.

Members of the Dewan Rakyat, 14th Malaysian Parliament

This is a list of the members of the Dewan Rakyat (House of Representatives) of the 14th Parliament of Malaysia.

PH indicator

A pH indicator is a halochromic chemical compound added in small amounts to a solution so the pH (acidity or basicity) of the solution can be determined visually. Hence, a pH indicator is a chemical detector for hydronium ions (H3O+) or hydrogen ions (H+) in the Arrhenius model. Normally, the indicator causes the color of the solution to change depending on the pH. Indicators can also show change in other physical properties; for example, olfactory indicators show change in their odor. The pH value of a neutral solution is 7.0 at 25°C (standard laboratory conditions). Solutions with a pH value below 7.0 are considered acidic and solutions with pH value above 7.0 are basic (alkaline). As most naturally occurring organic compounds are weak protolytes, carboxylic acids and amines, pH indicators find many applications in biology and analytical chemistry. Moreover, pH indicators form one of the three main types of indicator compounds used in chemical analysis. For the quantitative analysis of metal cations, the use of complexometric indicators is preferred, whereas the third compound class, the redox indicators, are used in titrations involving a redox reaction as the basis of the analysis.

PH meter

A pH meter is a scientific instrument that measures the hydrogen-ion activity in water-based solutions, indicating its acidity or alkalinity expressed as pH. The pH meter measures the difference in electrical potential between a pH electrode and a reference electrode, and so the pH meter is sometimes referred to as a "potentiometric pH meter". The difference in electrical potential relates to the acidity or pH of the solution. The pH meter is used in many applications ranging from laboratory experimentation to quality control.

Postgraduate education

Postgraduate education, or graduate education in North America, involves learning and studying for academic or professional degrees, academic or professional certificates, academic or professional diplomas, or other qualifications for which a first or bachelor's degree generally is required, and it is normally considered to be part of higher education. In North America, this level is typically referred to as graduate school (or sometimes colloquially as grad school).

The organization and structure of postgraduate education varies in different countries, as well as in different institutions within countries. This article outlines the basic types of courses and of teaching and examination methods, with some explanation of their history.

Provinces of the Philippines

The Provinces of the Philippines (Filipino: Mga Lalawigan ng Pilipinas/Mga Probinsya ng Pilipinas) are the primary political and administrative divisions of the Philippines. There are 81 provinces at present, further subdivided into component cities and municipalities. The National Capital Region, as well as independent cities, are independent of any provincial government. Each province is governed by an elected legislature called the Sangguniang Panlalawigan and by an elected governor.

The provinces are grouped into 17 regions based on geographical, cultural, and ethnological characteristics. Fourteen of these regions are designated with numbers corresponding to their geographic location in order from north to south. The Cordillera Administrative Region, National Capital Region, MIMAROPA Region and the Bangsamoro Autonomous Region in Muslim Mindanao do not have numerical designations.

Each province is a member of the League of Provinces of the Philippines, an organization which aims to address issues affecting provincial and metropolitan government administrations.

Soil pH

Soil pH is a measure of the acidity or basicity (alkalinity) of a soil. pH is defined as the negative logarithm (base 10) of the activity of hydronium ions (H+ or, more precisely, H3O+aq) in a solution. In soils, it is measured in a slurry of soil mixed with water (or a salt solution, such as 0.01 M CaCl2), and normally falls between 3 and 10, with 7 being neutral. Acid soils have a pH below 7 and alkaline soils have a pH above 7. Ultra-acidic soils (pH < 3.5) and very strongly alkaline soils (pH > 9) are rare.Soil pH is considered a master variable in soils as it affects many chemical processes. It specifically affects plant nutrient availability by controlling the chemical forms of the different nutrients and influencing the chemical reactions they undergo. The optimum pH range for most plants is between 5.5 and 7.5; however, many plants have adapted to thrive at pH values outside this range.

Thesis

A thesis or dissertation is a document submitted in support of candidature for an academic degree or professional qualification presenting the author's research and findings. In some contexts, the word "thesis" or a cognate is used for part of a bachelor's or master's course, while "dissertation" is normally applied to a doctorate, while in other contexts, the reverse is true. The term graduate thesis is sometimes used to refer to both master's theses and doctoral dissertations.The required complexity or quality of research of a thesis or dissertation can vary by country, university, or program, and the required minimum study period may thus vary significantly in duration.

The word "dissertation" can at times be used to describe a treatise without relation to obtaining an academic degree. The term "thesis" is also used to refer to the general claim of an essay or similar work.

Urine

Urine is a liquid by-product of metabolism in humans and in many animals. Urine flows from the kidneys through the ureters to the urinary bladder. Urination results in urine being excreted from the body through the urethra.

The cellular metabolism generates many by-products which are rich in nitrogen and must be cleared from the bloodstream, such as urea, uric acid, and creatinine. These by-products are expelled from the body during urination, which is the primary method for excreting water-soluble chemicals from the body. A urinalysis can detect nitrogenous wastes of the mammalian body.

Urine has a role in the earth's nitrogen cycle. In balanced ecosystems urine fertilizes the soil and thus helps plants to grow. Therefore, urine can be used as a fertilizer. Some animals use it to mark their territories. Historically, aged or fermented urine (known as lant) was also used for gunpowder production, household cleaning, tanning of leather and dyeing of textiles.

Human urine and feces are collectively referred to as human waste or human excreta, and are managed with a sanitation system. Livestock urine and feces also require proper management if the livestock population density is high.

Vagina

In mammals, the vagina is the elastic, muscular part of the female genital tract. In humans, it extends from the vulva to the cervix. The outer vaginal opening is normally partly covered by a membrane called the hymen. At the deep end, the cervix (neck of the uterus) bulges into the vagina. The vagina allows for sexual intercourse and birth. It also channels menstrual flow (menses), which occurs in humans and closely related primates as part of the monthly menstrual cycle.

Although research on the vagina is especially lacking for different animals, its location, structure and size is documented as varying among species. Female mammals usually have two external openings in the vulva, the urethral opening for the urinary tract and the vaginal opening for the genital tract. This is different from male mammals, who usually have a single urethral opening for both urination and reproduction. The vaginal opening is much larger than the nearby urethral opening, and both are protected by the labia in humans. In amphibians, birds, reptiles and monotremes, the cloaca is the single external opening for the gastrointestinal tract, the urinary, and reproductive tracts.

To accommodate smoother penetration of the vagina during sexual intercourse or other sexual activity, vaginal moisture increases during sexual arousal in human females and other female mammals. This increase in moisture provides vaginal lubrication, which reduces friction. The texture of the vaginal walls creates friction for the penis during sexual intercourse and stimulates it toward ejaculation, enabling fertilization. Along with pleasure and bonding, women's sexual behavior with others (which can include heterosexual or lesbian sexual activity) can result in sexually transmitted infections (STIs), the risk of which can be reduced by recommended safe sex practices. Other health issues may also affect the human vagina.

The vagina and vulva have evoked strong reactions in societies throughout history, including negative perceptions and language, cultural taboos, and their use as symbols for female sexuality, spirituality, or regeneration of life. In common speech, the word vagina is often used to refer to the vulva or to the female genitals in general. By its dictionary and anatomical definitions, however, vagina refers exclusively to the specific internal structure, and understanding the distinction can improve knowledge of the female genitalia and aid in healthcare communication.

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