The Origins of a Universal Measure: Søren Sørensen and the pH Scale

Every chemistry student learns the pH scale: a simple numeric range from 0 to 14 that tells us whether a solution is acidic, neutral, or alkaline. But the scale's elegant simplicity belies a fascinating history of scientific problem-solving, industrial necessity, and one man's determination to bring order to a chaotic field. The pH scale stands as one of the most enduring inventions of early 20th-century chemistry, and its creator, Danish chemist Søren Peter Lauritz Sørensen, remains a figure of profound influence in analytical science. His work transformed the way scientists describe acidity, enabling breakthroughs in biochemistry, medicine, and countless industrial processes.

Chemistry Before the pH Scale: A Problem of Standardization

In the late 19th and early 20th centuries, chemists had a variety of methods to measure acidity, but none were standardized or easily comparable. They used litmus paper, which turned red in acid and blue in base, but it gave only a rough qualitative indication. Other approaches involved titration with indicators or measuring electrical conductivity, but these were laborious and required careful calibration. Scientists working in fields such as biochemistry, fermentation, and medicine urgently needed a reliable, quantitative way to express the concentration of hydrogen ions in a solution.

Acidity was understood to be related to the presence of hydrogen ions (H⁺), but expressing that concentration in molarity produced unwieldy numbers, especially for very dilute solutions. For example, a solution with a hydrogen ion concentration of 0.0000001 moles per liter was difficult to conceptualize and compare with others. The scientific community recognized the need for a logarithmic compression that would make such numbers human-readable. That challenge fell to Sørensen, who was perfectly positioned to solve it.

Furthermore, the lack of a universal pH standard hindered reproducibility across laboratories. Researchers publishing results on enzyme activity or bacterial growth could not reliably compare their acidity conditions. The stage was set for a transformative innovation.

Who Was Søren Sørensen?

Early Life and Education

Søren Sørensen was born in 1868 in the small Danish town of Havrebjerg. He studied chemistry at the University of Copenhagen, earning his doctorate in 1899. His early research focused on inorganic chemistry and the composition of minerals, but his career path shifted dramatically when he joined the Carlsberg Laboratory in 1901. The Carlsberg Laboratory was the research arm of the Carlsberg Brewery, established by brewer J.C. Jacobsen, who believed that scientific inquiry was essential to improving beer production and, by extension, the broader field of biotechnology.

Jacobsen founded the laboratory in 1876 with a clear mission: to advance the science of brewing through chemical and physiological research. By the time Sørensen arrived, the lab was already renowned for its work on yeast, enzymes, and protein chemistry. This environment fostered interdisciplinary thinking and rigorous experimentation.

The Carlsberg Laboratory: A Cradle of Innovation

At Carlsberg, Sørensen specialized in the chemistry of proteins and fermentation. It was in this context that he encountered a persistent problem: how to precisely control the acidity of enzyme reactions. Enzymes are notoriously sensitive to pH, and even small variations could alter their activity. Sørensen realized that the existing methods of measuring acidity were inadequate for the precision required in biochemical research. He set out to create a universally applicable scale.

The Carlsberg Laboratory gave Sørensen access to state-of-the-art equipment, including the hydrogen electrode, which could measure the electromotive force of a solution with great accuracy. This technology was essential to his eventual breakthrough.

The Breakthrough Concept: Negative Logarithm

In 1909, Sørensen published a landmark paper titled "Études Enzymatiques II: Sur la mésure et l'importance de la concentration des ions hydrogène dans les réactions enzymatiques" (Enzymatic Studies II: On the measurement and importance of hydrogen ion concentration in enzymatic reactions). In it, he proposed a deceptively simple definition: pH = −log₁₀[H⁺], where [H⁺] is the concentration of hydrogen ions expressed in moles per liter.

This logarithmic transformation meant that a tenfold change in hydrogen ion concentration corresponded to a change of exactly one pH unit. A solution with [H⁺] = 10⁻⁷ M (0.0000001 mol/L) had a pH of 7, which Sørensen designated as neutral. Solutions with [H⁺] > 10⁻⁷ M (i.e., pH < 7) were acidic, and those with [H⁺] < 10⁻⁷ M (pH > 7) were basic or alkaline. The range from 0 to 14 conveniently covered most aqueous solutions encountered in the laboratory and in nature. The notation "pH" stood for "potential of hydrogen," though Sørensen himself used the French term puissance d'hydrogène.

Why Logarithmic?

The choice of a logarithmic scale was no accident. Hydrogen ion concentrations vary over many orders of magnitude—from as low as 10⁻¹⁴ M in strong bases to as high as 10 M in concentrated strong acids. A linear scale would be impossible to read for such a wide range. By using base-10 logarithms, Sørensen compressed this vast span into a handy 0–14 range, where each integer step represents a tenfold change. This clever mathematical trick made the scale both intuitive and practical for everyday use.

The First pH Scale: A Simple but Powerful Tool

Sørensen's initial scale was based on measurements obtained using a hydrogen electrode and a reference electrode, which allowed precise voltage measurements that directly correlated with hydrogen ion activity. He published tables of pH values for common solutions, including buffers used in enzyme studies. The scale was immediately adopted by his colleagues at Carlsberg and gradually spread through the world of biochemistry and industrial chemistry.

The beauty of the pH scale lay in its integration of thermodynamic theory with practical utility. It gave researchers a single number to describe a solution's acid-base character, making it possible to compare results across experiments, laboratories, and even disciplines. Sørensen's invention transformed the nebulous concept of "acidity" into a reproducible, quantifiable property.

Refinements and Wider Adoption

From Concentration to Activity

Although Sørensen's original definition used hydrogen ion concentration, later scientists realized that what actually matters in chemical reactions is the activity of hydrogen ions, not just their stoichiometric concentration. In 1924, the Dutch-American chemists Johannes Nicolaus Brønsted and Thomas Lowry independently refined the definition to pH = −log₁₀(aH⁺), where aH⁺ represents the hydrogen ion activity. However, for most practical purposes, the original concentration-based definition remains accurate enough for everyday use, especially in dilute solutions.

The Development of pH Meters

Early adopters included the brewing industry (naturally), the emerging pharmaceutical sector, and medical laboratories. By the 1930s, pH meters had been developed that simplified measurement even further—no more cumbersome hydrogen electrodes. The first commercially available pH meter was produced by the Dutch company Radiometer in 1936, and later by Beckman Instruments in the United States. The invention of the glass electrode—which measures pH by detecting voltage changes across a thin glass membrane—made portable, routine pH measurement possible. This glass electrode, invented by Fritz Haber and Zygmunt Klemensiewicz in 1909, was improved significantly by Arnold O. Beckman in 1935, leading to the first practical pH meter.

Today, the pH scale is universally taught and used, from high school chemistry labs to advanced research in genomics and materials science. It underpins everything from soil testing and water treatment to the formulation of cosmetics and pharmaceuticals. Chemistry World provides an excellent historical overview of the pH scale's development.

Applications of the pH Scale Across Science and Industry

The influence of Sørensen's invention extends far beyond the laboratory bench. The pH scale is now an indispensable tool in countless fields:

Environmental Science

Monitoring pH is critical for assessing water quality in rivers, lakes, and oceans. Acid rain, caused by industrial emissions of sulfur dioxide and nitrogen oxides, can lower the pH of natural waters, harming aquatic life. Soil pH affects plant growth, nutrient availability, and the activity of microbes—farmers routinely test soil pH to optimize crop yields. The U.S. Environmental Protection Agency provides extensive data on acid rain pH monitoring, showcasing Sørensen's scale in real-world environmental policy.

Medical Diagnostics

Human blood has a tightly regulated pH of about 7.4, and deviations can indicate serious health conditions. Physicians measure blood pH to diagnose acidosis (pH < 7.35) or alkalosis (pH > 7.45), which can result from kidney dysfunction, lung disease, or metabolic disorders. Urinalysis and gastric fluid analysis also rely on pH measurements. The development of the pH scale directly enabled the modern understanding of acid-base balance in the human body, including the role of the bicarbonate buffering system.

Food and Beverage Industry

From brewery to dairy, pH controls flavor, texture, and safety. Beer fermentation proceeds optimally at a specific pH range; cheese-making depends on pH to curdle milk properly; and canned foods must be sufficiently acidic to prevent the growth of dangerous bacteria like Clostridium botulinum. The Carlsberg Laboratory itself continues to explore the role of pH in brewing, carrying on Sørensen's legacy. The pH of wine, for example, influences its color, taste, and aging potential.

Industrial Chemistry

Chemical manufacturing processes, from producing fertilizers to synthesizing polymers, require precise pH control to maximize yield and minimize side reactions. Wastewater treatment plants use pH adjustment to remove heavy metals and neutralize corrosive effluents before discharge. The scale provides a universal language for engineers and chemists worldwide. In the pharmaceutical industry, pH affects the solubility and stability of active ingredients, making it a critical parameter in drug formulation.

Agriculture and Soil Science

Soil pH is a master variable that determines the availability of essential nutrients like nitrogen, phosphorus, and potassium. Most crops thrive in a pH range of 6.0 to 7.5. Lime is applied to raise pH in acidic soils, while sulfur is used to lower pH in alkaline soils. The ability to measure and adjust soil pH has dramatically improved crop yields worldwide.

Modern Extensions and Future Directions

Modern research has extended the concept beyond the original 0–14 range. Ultra-strong acids, such as fluoroantimonic acid, have pH values far below 0, while superbasic solutions of lithium diisopropylamide can have pH values above 14. These extensions rely on the same logarithmic principle Sørensen introduced. Additionally, pH measurement is now possible at microscopic scales using fluorescent dyes and microelectrodes, enabling real-time monitoring inside living cells.

Non-aqueous pH scales have been developed for solvents like ethanol and acetonitrile, using the same negative-logarithm concept but with different reference standards. The IUPAC has defined a unified acidity scale (pHabs) that can compare acidities across different solvents. This represents the latest chapter in the ongoing refinement of Sørensen's original idea.

Legacy and Continuing Relevance

Honors and Recognition

Sørensen received numerous honors during his lifetime, including membership in the Royal Danish Academy of Sciences and Letters. In 2009, on the centenary of his pH proposal, the International Union of Pure and Applied Chemistry (IUPAC) recognized the scale's global significance. The Carlsberg Laboratory maintains a small museum dedicated to his work, and his portrait hangs in chemistry departments around the world.

The American Chemical Society's National Historic Chemical Landmark article on the pH scale provides an excellent overview. Another valuable resource is the original 1909 paper available through the Royal Society of Chemistry's historical archives. Encyclopedia Britannica also offers a concise history of the pH scale.

Conclusion: A Simple Idea That Changed Science

Sørensen's pH scale exemplifies how a single, elegant concept can revolutionize an entire field. By replacing a confusing jumble of measurement methods with a straightforward numeric scale, he made acidity a universal language for scientists, engineers, and technicians. The scale's endurance—over 115 years and counting—speaks to its fundamental importance. Today, whenever a student reads a pH meter or a farmer tests soil, they are using a tool that sprang from the work of a Danish chemist who simply wanted to measure fermentation reactions more accurately. The history of the pH scale is not just a story of scientific progress; it is a reminder that the most impactful inventions often emerge from solving a specific, practical problem with clarity and ingenuity.