b. January 9, 1868, Havrebjerg (near Slagelse), Denmark
d. February 12, 1939, Charlottenlund, Denmark
|Søren Peter Lauritz Sørensen, a Danish biochemist, suggested a convenient way of expressing acidity the negative logarithm of hydrogen ion concentration (pH).|
Søren Peter Lauritz Sørensen, a Danish biochemist, is the
man who established the concept of pH, defining it as pH = -log[H+]. He was born in Havrevjerg near
Slaglese in Denmark as as son of an owner farmer (Søren's father
Sørensen and Søren's mother - Kirstine
Søren's father - Hans Sørensen
Søren's mother - Kirstine Katrine Sørensen
Søren Peter Lauritz Sørensen first began to study medicine
at the University of Copenhagen, but soon moved to chemistry. He
graduated from the University of Copenhagen in 1881, and for
following ten years he worked on inorganic syntheses under Prof.
S.M. Jorgensen (1837-1914) at the Technical University of Denmark
in Copenhagen. Sørensen obtained his Ph.D. in 1899.
wife - Anna
|Søren Peter Lauritz Sørensen was maried with Anna Louise Willumsen on 18 November 1892 and they had four children: Inge Sørensen, Viggo Sørensen, Arthur Sørensen, Annelise Sørensen.|
Sørensen's work and life were connected with Carlsberg
Laboratory in Copenhagen. J.C. Jacobsen (1811-1887), the founder
of the Carlsberg Brewery in Copenhagen, established the Carlsberg
Laboratory in 1876. Professor J. Kjeldahl (1849-1900) became the
first head of the Chemical Department of the Laboratory and
worked until his tragic death in July 1900. Professor. S.P.L. Sørensen
was invited as the successor of Prof. J. Kjeldahl in 1901. It was
at this laboratory that Prof. Sorensen carried out numerous
pioneering works by his elaborate experimental technique in the
field of biochemistry.
|While at the Carlsberg Laboratories, Sørensen started to study amino acids, proteins and enzymes. In earlier stage his researches are categorized in the following four domains: (1) electrometric method of determining hydrogen ion concentration, (2) preparation of pH buffer solutions, (3) colorimetric method of measuring pH, and (4) application of these procedures to study on enzymes and proteins. The results in these fields by his group were summarized in amino acids and protein. Their brilliant achievements were published in 17 reports from the Laboratory. Some of these became a classic work in the real sense of the word and continued materials for many papers, i.e., serving as a foundation of protein chemistry. It is noted that he studied on coagulation of proteins on heating and then succeeded in crystallizing the protein egg alubumin by ammonium sulphate precipitation (1936). The Carlsberg Laboratory became renowned as one of the world's most productive centers of study in the field of biochemistry, attracting many chemists all over the world. Sørensen subsequently became a leader in the application of thermodynamics to proteins chemistry, and in this work he was assisted by his second wife, Margrethe Høyrup Sørensen.|
|Because hydrogen ion concentration played a key role in enzymatic reactions he devised a simple way of expressing it. By taking a negative logarithm of hydrogen ion concentration a convenient scale can be established; this is the well-known pH value. Numerical values based on this unit, now universally in use, give an indication of the acidity of solutions. He also developed buffer solutions to maintain constant pH of solutions (Sørensen buffers).|
|Historical background for
In the late 1880's, Svante Arrhenius proposed that acids were substances that delivered hydrogen ion to the solution. He has also pointed out that the law of mass action could be applied to ionic reactions, such as an acid dissociating into hydrogen ion and a negatively charged anion. This idea was followed up by Wilhelm Ostwald, who calculated the dissociation constants (the modern symbol is Ka) of many weak acids. Ostwald also showed that the value of the constant is measure of an acid's strength. By 1894, the dissociation constant of water (today called Kw) was measured to the modern value of 1×10-14. In 1904, H. Friedenthal recommended that the hydrogen ion concentration be used to characterize solutions. He also pointed out that alkaline (modern word = basic) solutions could also be characterized this way since the hydroxyl concentration was always (1×10-14 / the hydrogen ion concentration). Many consider this to be the real introduction of the pH scale.
Søren Sørensen visting
Cornell University (1924)
|The context for the introduction of pH was the slow changeover from the old color-change tests for indicating the degree of acidity or basicity to electrical methods. In the latter, the current generated in an electrochemical cell by ions migrating to oppositely charged electrodes was measured, using a highly sensitive (and delicate) galvanometer. Until Sørensen developed the pH scale, there was no widely accepted way of expressing hydrogen ion concentrations. His scale removes the awkward negative power for hydrogen ion concentrations that range over orders of magnitude: from about ~12 M at the high end to ~10-15 M at the low end. Instead Sørensen suggested that the power could be represented by a pH scale in which 7 is neutral, and 1 and 14 are the extremes of acidity and alkalinity, respectively.|
|pH introduced into the
The pH scale was quickly accepted by the biochemical research community in their studies of the facinating ability of living tissues to "buffer" against excessive acidity or alkalinity. Largely due to the German medical chemist Leonor Michaelis (1875-1949), who published a book on hydrogen ion concentration in 1914, wider community of chemists finally adopted the pH scale. The use of pH became even more widespread in 1935 when Arnold Beckman (b. 1900) developed and sold a simple portable direct-reading pH meter.
The letters pH are an abbreviate for "pondus hydrogenii" (translated as potential hydrogen) meaning hydrogen power as acidity is caused by a predominance of hydrogen ions (H+). Dr. Sørensen has been credited as the founder of the modern pH concept.
In Sørensen 's original paper, pH is written as PH. According to the Compact Oxford English Dictionary, the modern notation "pH" was first adopted in 1920 by W. M. Clark (inventor of the Clark oxygen electrode) for typographical convenience. "p-Functions" have also been adopted for other concentrations and concentration-related numbers. For example, "pCa = 5.0" means a concentration of calcium ions equal to 10-5 M, and pKa = 4.0 means an acid dissociation constant equal to 10-4.
|The pH Definition Using
the Concentration of Hydrogen Ions
The pH scale is a measure of acid/base strength. It is defined on a logarithmic scale using the molar concentration of hydronium ions in solution.
Pure water autoionizes to produce equal concentrations of hydronium and hydroxide ions:
Whose equilibrium obeys the law of mass action in the form Kw = [H3O+] [OH-] = 1.0×10-14 (at 25 0C). This modern form of the equation for water autoionization recognizes that protons do not exist in solution but instead are bound to an electron lone pair in water: H+ + H2O => H3O+. The obsolete term hydrogen ion and its concentration [H+] have been replaced by "hydronium ion" and [H3O+] but we continue to use pH (and not pH3O). Since the hydronium ion concentration and the hydroxide ion concentration are equal in pure water, it follows that [H3O+] = [OH-] = 1.0 10-7. Then the pH of pure water is pH = - log(1.0×10-7) = 7.00. In dilute acid solution the hydronium ion concentration is higher; e.g., in micromolar hydrochloric acid, 10-6 M HCl, the hydronium ion concentration is [H3O+] = 10-6 mol/L so that the pH = 6.00. That is, one step lower (higher) on the pH scale represents 10 times higher (or lower) hydronium ion concentration.
In a similar way to pH, the concentration of hydroxide ion is also expressed on a logarithmic negative power scale: pOH = - log([OH-] ). Further, since water autoionization equlibrium relates [H3O+] to [OH-] , the pH and pOH are related: [H3O+] [OH-] = 1.0×10-14, so that -log([H3O+]) - log([OH-]]) = -14.00 or pH + pOH = 14.00.
This scale covers a very large range of [H+], from 0.1 to 10-14. Because the negative log of [H+] is used in the pH scale, the pH scale usually have positive values. Furthermore, the larger the pH, the smaller the [H+].When [H+] is high, we usually do not use the pH value, but simply the [H+]. For example, when [H+] = 1.0, pH = 0. Because we seldom say the pH is 0, and that is why you consider pH = 0 such an odd expression. A pH = -0.30 is equivalent to a [H+] of 2.0 M. Negative pH values are only for academic exercises. Using the concentrations directly conveys a better sense than the pH scales.
Definition that Uses the Hydrogen Ion Activity:
pH has been more accurately defined as pH = -log aH+ where aH+ is the hydrogen ion activity. In solutions that contain other ions, activity and concentration are not the same. The activity is an effective concentration of hydrogen ions, rather than the true concentration; it accounts for the fact that other ions surrounding the hydrogen ions will shield them and affect their ability to participate in chemical reactions. These other ions effectively change the hydrogen ion concentration in any process that involves H+. In practice, Sørenson's original definition can still be used, because the instrument used to make the measurement can be calibrated with solutions of known [H+], with the concentration of background ions carefully controlled.
The Experimental Definition:
IUPAC has endorsed a pH scale based on comparison with a standard buffer of known pH using electrochemical measurements. The IUPAC pH scale is very slightly different from the theoretical definition, since it considers factors that are not included in the (thermodynamic) theoretical pH.
|Comparison of acidity of
solutions in aqueous and non-aqueous solutions:
The hydrogen ion concentration in pure water around room temperature is about 1.0×10-7 M. A pH of 7 is considered "neutral", because the concentration of hydrogen ions is exactly equal to the concentration of hydroxide (OH-) ions produced by dissociation of the water. Increasing the concentration of hydrogen ions above 1.0×10-7 M produces a solution with a pH of less than 7, and the solution is considered "acidic". Decreasing the concentration below 1.0×10-7 M produces a solution with a pH above 7, and the solution is considered "alkaline" or "basic". pH is often used to compare solution acidities. For example, a solution of pH 1 is said to be 10 times as acidic as a solution of pH 2, because the hydrogen ion concentration at pH 1 is ten times the hydrogen ion concentration at pH 2. This is correct as long as the solutions being compared both use the same solvent. You can't use pH to compare the acidities in different solvents because the neutral pH is different for each solvent. For example, the concentration of hydrogen ions in pure ethanol is about 1.58×10-10 M, so ethanol is neutral at pH 9.8. A solution with a pH of 8 would be considered acidic in ethanol, but basic in water!
|Sørensen retired from the post of the director of the chemical section of Carlsberg Laboratories in 1938. Professor Sørensen, gifted with rare talent as a chemist with a fervent hope for progress of human welfare and peace, passed away peacefully on Feburary 12, 1939, Charlottenlund, Denmark.|
See related material
available in the Internet:
What is pH?
This text has been compiled from
the biographies of Sørensen
available in the Internet:
( 1, 2, 3, 4, 5, 6 ).
(updated & corrected on February 17, 2003)