pH – Slurry Pump ELM – Vertical Slurry Pump EVM(R)
History
The concept of p[H] was first introduced by Danish chemist Sren Peder Lauritz Srensen at the Carlsberg Laboratory in 1909 and revised to the modern pH in 1924 after it became apparent that electromotive force in cells depended on activity rather than concentration of hydrogen ions. It is unknown what the exact definition of p is. Some references suggest the p stands for ower, others refer to the German word otenz (meaning power in German), still others refer to otential. Jens Norby published a paper in 2000 arguing that p is a constant and stands for egative logarithm; which has also been used in other works. H then stands for Hydrogen. According to the Carlsberg Foundation pH stands for “power of hydrogen”. Other suggestions that have surfaced over the years are that the p stands for puissance (also meaning power but then the Carlsberg Laboratory was French speaking) or that pH stands for the Latin terms pondus Hydrogenii or potentia hydrogenii. It is also suggested that Srensen used the letters p and q (commonly paired letters in mathematics) simply to label the test solution (p) and the reference solution (q).
Definitions
Mathematical Definition
pH is defined as minus the decimal logarithm of the hydrogen ion activity in a solution. By virtue of its logarithmic nature, pH is a dimensionless quantity.
where aH is the (dimensionless) activity of hydrogen ions. The reason for this definition is that aH is a property of a single ion, which can only be measured experimentally by means of an ion-selective electrode which responds, according to the Nernst equation, to hydrogen ion activity. pH is commonly measured by means of a combined glass electrode, which measures the potential difference, or electromotive force, E, between an electrode sensitive to the hydrogen ion activity and a reference electrode, such as a calomel electrode or a silver chloride electrode. The combined glass electrode ideally follows the Nernst equation:
where E is a measured potential , E0 is the standard electrode potential, that is, the electode potential for the standard state in which the activity is one. R is the gas constant, T is the temperature in kelvins, F is the Faraday constant and n is the number of electrons transferred, one in this instance. The electrode potential, E, is proportional to the logarithm of the hydrogen ion activity.
This definition, by itself, is wholly impractical, because the hydrogen ion activity is the product of the concentration and an activity coefficient. The single-ion activity coefficient of the hydrogen ion is a quantity which cannot be measured experimentally. To get around this difficulty, the electrode is calibrated in terms of solutions of known activity.
The operational definition of pH is officially defined by International Standard ISO 31-8 as follows: For a solution X, first measure the electromotive force EX of the galvanic cell
reference electrode | concentrated solution of KCl || solution X | H2 | Pt
and then also measure the electromotive force ES of a galvanic cell that differs from the above one only by the replacement of the solution X of unknown pH, pH(X), by a solution S of a known standard pH, pH(S). The pH of X is then
The difference between the pH of solution X and the pH of the standard solution depends only on the difference between two measured potentials. Thus, pH is obtained from a potential measured with an electrode calibrated against one or more pH standards; a pH meter setting is adjusted such that the meter reading for a solution of a standard is equal to the value pH(S). Values pH(S) for a range of standard solutions S, along with further details, are given in the IUPAC recommendations. The standard solutions are often described as standard buffer solution. In practice, it is better to use two or more standard buffers to allow for small deviations from Nernst-law ideality in real electrodes. Note that because the temperature occurs in the defining equations, the pH of a solution is temperature-dependent.
Measurement of extremely low pH values, such as some very acidic mine waters, requires special procedures. Calibration of the electrode in such cases can be done with standard solutions of concentrated sulfuric acid, whose pH values can be calculated with using Pitzer parameters to calculate activity coefficients.
pH is an example of an acidity function. Hydrogen ion concentrations can be measured in non-aqueous solvents, but this leads, in effect, to a different acidity function, because the standard state for a non-aqueous solvent is different from the standard state for water. Superacids are a class of non-aqueous acids for which the Hammett acidity function, H0, has been developed.
p[H]
This was the original definition of Srensen, which was superseded in favour of pH.[when?] 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 alkali in the presence of a relatively high concentration of background electrolyte. Since the concentrations of acid and alkali 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. The calibration yieds 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 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 difference between p[H] and pH is quite small. It has been stated that pH = p[H] + 0.04. Unfortunately it is common practice to use the term “pH” for both types of measurement.
pOH
pOH is sometimes used as a measure of the concentration of hydroxide ions, OH, or alkalinity. pOH is not measured independently, but is derived from pH. The concentration of hydroxide ions in water is related to the concentration of hydrogen ions by
[OH = KW /[H+]
where KW is the self-ionisation constant of water. Taking cologarithms
pOH = pKW pH.
So, at room temperature pOH 14 pH. However this relationship is not strictly valid in other circumstances, such as in measurements of soil alkalinity.
Applications
Some typical pH values
Pure water has a pH around 7; the exact values depends on the temperature. When an acid is dissolved in water the pH will be less than 7 and 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 dm3 has a pH of 0. A solution of a strong alkali, such as sodium hydroxide, at concentration 1 mol dm3 has a pH of 14. Thus, measured pH values will mostly lie in the range 0 to 14. Since pH is a logarithmic scale a difference of one pH unit is equivalent to a ten-fold difference in hydrogen ion concentration.
Because the glass electrode (and other ion selective electrodes) responds to activity, the electrode 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.
An approximate measure of pH may be obtained by using a pH indicator. A pH indicator is a substance that changes colour around a particular pH value. It is a weak acid or weak base and the colour change occurs around 1 pH unit either side of its acid dissociation constant, or pKa, value. For example, the naturally occurring indicator litmus is red in acidic solutions (pH<7) and blue in alkaline (pH>7) solutions. Universal indicator consists of a mixture of indicators such that there is a continuous colour change from about pH 2 to pH 10. Universal indicator paper is simple paper that has been impregnated with universal indicator.
Universal indicator components
Indicator
Low pH color
Transition pH range
High pH color
Thymol blue (first transition)
red
1.22.8
orange
Methyl red
red
4.46.2
yellow
Bromothymol blue
yellow
6.07.6
blue
Thymol blue (second transition)
yellow
8.09.6
blue
Phenolphthalein
colorless
8.310.0
purple
A solution whose pH is 7 is said to be neutral, that is, it is neither acidic nor basic. Water is subject to a self-ionisation process.
H2O H+ + OH
The dissociation constant, KW, has a value of about 10-14, so in neutral solution of a salt both the hydrogen ion concentration and hydroxide ion concentration are about 10-7 mol dm-3. The pH of pure water decreases with increasing temperatures. For example, the pH of pure water at 50 C is 6.55. Note, however, that water that has been exposed to air is mildly acidic. This is because water absorbs carbon dioxide from the air, which is then slowly converted into carbonic acid, which dissociates to liberate hydrogen ions:
CO2 + H2O H2CO3 HCO3 + H+
Calculation of pH for weak and strong acids
In the case of a strong acid, there is complete dissociation, so the pH is simply equal to minus the logarithm of the acid concentration. For example, a 0.01 molar solution of hydrochloric acid has a pH of og(0.01), that is, pH = 2.
The pH of a solution of a weak acid may be calculated by means of an ICE table. For acids with a pKa value greater than about 2,
pH = ( pKa log c0),
where c0 is the concentration of the acid. This is equivalent to Burrows’ weak acid pH equation
A more general method is as follows. Consider the case of dissolving a weak acid, HA, in water. First write down the equilibrium expression.
HA A + H+
The equilibrium constant for this reaction is specified by
where [] indicates a concentration. The analytical concentration of the two reagents, CA for [A and CH for [H+] must be equal to the sum of concentrations of those species that contain the reagents. CH is the concentration of added mineral acid.
CA = [A + Ka[A[H+]
CH = [H+] + Ka[A[H+]
From the first equation
Substitution of this expression into the second equation gives
This simplifies to a quadratic equation in the hydrogen ion concentration
Solution of this equation gives [H+] and hence pH.
This method can also be used for polyprotic acids. For example, for the diprotic acid oxalic acid, writing A2 for the oxalate ion,
CA = [A2 + 1[A2[H+] + 2[A2[H+]2
CH = [H+] + 1[A2[H+] + 22[A2[H+]2
where 1 and 2 are cumulative protonation constants. Following the same procedure of substituting from the first equation into the second, a cubic equation in [H+] results. In general, the degree of the equation is one more than the number of ionisable protons. The solution of these equations can be obtained relatively easily with the aid of a spreadsheet such as EXCEL or Origin. The pH always has an amount of fractional figures equal to the amount of significant figures of the concentration.
pH in nature
Hydrangea macrophylla blossoms are either pink or blue, depending on a pH-dependent mobilization and uptake of soil aluminium into the plants.
pH-dependent plant pigments that can be used as pH indicators occur in many plants, including hibiscus, marigold, red cabbage (anthocyanin), and red wine.
Seawater
The pH of seawater plays an important role in the ocean’s carbon cycle and there is evidence of ongoing ocean acidification caused by carbon dioxide emissions. However, pH measurement is complicated by the chemical properties of seawater, and several distinct pH scales exist in chemical oceanography.
As part of its operational definition of the pH scale, the IUPAC define 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 consequently are not recommended for use in characterising 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. 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+ + SO42 HSO4, such that the total scale includes the effect of both protons (free hydrogen ions) and hydrogen sulfate ions:
[H+]T = [H+]F + [HSO4
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. Analytically, only [H+]T can be determined, therefore, [H+]F must be estimated using the [SO42 and the stability constant of HSO4, KS*:
[H+]F = [H+]T [HSO4 = [H+]T ( 1 + [SO42 / KS* )1
However, it is difficult to estimate KS* 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 + [HSO4 + [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. Consequently, 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 + [HSO4 ) = log [H+]T
pHSWS = log ( [H+]F + [HSO4 + [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, particularly in relation to the ocean’s carbonate system. 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
Compartment
pH
Gastric acid
0.7
Lysosomes
4.5
Granules of chromaffin cells
5.5
Human skin
5.5
Urine
6.0
Neutral H2O at 37 C
6.81
Cytosol
7.2
Cerebrospinal fluid (CSF)
7.3
Blood
7.347.45
Mitochondrial matrix
7.5
Pancreas secretions
8.1
General symptoms of acidosis, resulting from decrease in body pH.
The pH of different cellular compartments, body fluids, and organs is usually tightly regulated in a process called acid-base homeostasis.
The pH of blood is usually slightly basic with a value of pH 7.4. 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 demineralisation.
Enzymes and other proteins have an optimum pH range and can become inactivated or denatured outside this range.
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.
In the body, pH can be estimated from known base excess (be) and bicarbonate concentration (HCO3) by the following equation:
Reference ranges for blood tests, showing concentration of protons (purple) at left. It can be seen that the ranges are kept in a narrow range, and that free protons are among the compounds with the very smallest mass concentrations.
See also
Acidosis
Alkalosis
pH meter
pH indicator
References
^ [http://pages.pomona.edu/~wes04747/handout/equil_io.htm “General Chemistry Handouts, Pomona College: How to Solve Equilibrium Problems: An Approach Based on Stoichiometry Part II, Acid-Base Reactions”]. http://pages.pomona.edu/~wes04747/handout/equil_io.htm.
^ a b original reference requires subscription Isaac Feldman (1956). “Use and Abuse of pH measurements”. http://pubs.acs.org/doi/pdf/10.1021/ac60120a014. – cited and applied in publicly accessible “13C and 170 NMR Binding Constant Studies of Uranyl Carbonate Complexes in Near-Neutral Aqueous Solution Yucca Mountain Project Milestone Report 3351”. –
^ “The Measurement of pH – Definition, Standards and Procedures Report of the Working Party on pH, IUPAC Provisional Recommendation”]. 2001. http://www.iupac.org/reports/provisional/abstract01/rondinini_prs.pdf. A proposal to revise the current IUPAC 1985 and ISO 31-8 definition of pH.
^ 319 Sorensen, S. P. L., Enzymstudien. II, Ueber die Messung und die Bedeutung der Wasserstoffionenkonzentration bei enzymatischen Prozessen, Biochem. Zeitschr., 1909, vol. 21, pp. 131-304.
^ Two other publications appeared in 1909 one in French and one in Danisch
^ a b c Carlsberg Group Company History Page, http://www.carlsberggroup.com/Company/Research/Pages/pHValue.aspx
^ University of Waterloo – The pH Scale, http://www.science.uwaterloo.ca/~cchieh/cact/c123/ph.html
^ Nrby, Jens. 2000. The origin and the meaning of the little p in pH. Trends in the Biochemical Sciences 25:36-37., http://download.cell.com/trends/biochemical-sciences/pdf/PIIS0968000499015170.pdf
^ Fundamentals of Analytical Toxicology, http://books.google.com.br/books?id=LBag6XlAJY0C
^ One-Hundred Years of pH Rollie J. Myers Journal of Chemical Education Vol. 87 No. 1 January 2010 doi:10.1021/ed800002c
^ “pH”. IUPAC Goldbook. http://goldbook.iupac.org/P04524.html.
^ Quantities and units Part 8: Physical chemistry and molecular physics, Annex C (normative): pH. International Organization for Standardization, 1992.
^ Definitions of pH scales, standard reference values, measurement of pH, and related terminology. Pure Appl. Chem. (1985), 57, pp 531542.
^ Nordstrom, DK et al. (2000) Negative pH and extremely acidic mine waters from Iron Mountain California. Environ Sci Technol,34, 254-258.
^ Zemaitis, J.F.; Clark, D.M; Rafal, M; Scrivner, N.C. (1986). Handbook of Aqueous Electrolyte Thermodynamics: Theory & Application. Wiley. ISBN 978-0-8169-0350-4. Chapter 4
^ Rossotti, F.J.C.; Rossotti, H. (1965). “Potentiometric titrations using Gran plots: A textbook omission”. J. Chem. Ed. 42: 375378.
^ Mendham, J.; Denney, R. C.; Barnes, J. D.; Thomas, M.J.K.; Denney, R. C.; 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”
^ chemistry.about.com
^ Raven, J. A. et al. (2005). Ocean acidification due to increasing atmospheric carbon dioxide. Royal Society, London, UK.
^ 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).
^ Hansson, I. (1973) A new set of pH-scales and standard buffers for seawater. Deep Sea Research, 20: 479-491.
^ Dickson, A. G. (1984) pH scales and proton-transfer reactions in saline media such as sea water. Geochim. Cosmochim. Acta, 48: 22992308.
^ Boron, Walter, F.; Boulpaep, E.L. (2004). Medical Physiology: A Cellular And Molecular Approaoch. Elsevier/Saunders. ISBN 1-4160-2328-3.
^ Answers.com Medical Encyclopedia: Metabolic Acidosis: Causes and symptoms By Altha Roberts Edgren. Retrieved on April 13, 2009
^ Symptoms mentioned in both metabolic and respiratory acidosis from the following two references:
– Wrongdiagnosis.com > Symptoms of Metabolic Acidosis Retrieved on April 13, 2009
– Wrongdiagnosis.com > Symptoms of Respiratory acidosis Retrieved on April 13, 2009
^ Medical Calculators > Calculated Bicarbonate & Base Excess teven Pon, MD, Weill Medical College of Cornell University
External links
Online pH Calculator
Categories: Acid-base chemistry | Equilibrium chemistry | Units of measure | Water quality indicatorsHidden categories: Vague or ambiguous time
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