In chemistry, pH is a logarithmic scale used to determine the acidity of a medium. This is approximately the negative base 10 logarithm of the molar concentration, measured in units of moles per liter of hydrogen ions. It can also be called an indicator of the acidity of the environment. More precisely, it is the negative base 10 logarithm of hydrogen ion activity. At 25°C, solutions with a pH less than 7 are acidic, and solutions with a pH greater than 7 are basic. The neutral pH value is temperature dependent and is less than 7 as the temperature rises. Pure water is neutral, pH=7 (at 25°C), neither acidic nor alkaline. Contrary to popular belief, the pH value can be less than 0 or greater than 14 for very strong acids and bases, respectively.
Measurements of pH are important in agronomy, medicine, chemistry, water treatment and many other areas.
The pH scale is relevant for a set of standard solutions, the acidity of which is established by the internationalagreement. Primary pH standards are determined using a transfer concentration cell by measuring the potential difference between a hydrogen electrode and a standard electrode such as silver chloride. The pH of aqueous solutions can be measured with a glass electrode and a pH meter or indicator.
The pH concept was first introduced by the Danish chemist Søren Peter Laurits Sørensen at the Carlsberg laboratory in 1909 and revised to the current pH level in 1924 to accommodate definitions and measurements in terms of electrochemical cells. In the early works, the notation had the letter H in lowercase p, which means: pH.
Origin of the name
The exact meaning of the p is disputed, but according to the Carlsberg Foundation, pH means "the power of hydrogen." It has also been suggested that the p stands for the German word potenz ("power"), others refer to the French puisance (also meaning "power", based on the fact that the Carlsberg laboratory was French). Another suggestion is that p refers to the Latin term pondus hydroii (amount of hydrogen), potentio hydroii (capacity of hydrogen), or potential hydroli (hydrogen potential). It is also suggested that Sørensen used the letters p and q (usually conjugate letters in mathematics) simply to denote test solution (p) and reference solution (q). Currently, in chemistry, p stands for the decimal logarithm, and is also used in the term pKa, used for dissociation constants of the acidity of a medium.
Bacteriologist Alice Evans, known for the influence of her work on dairy products and food safety, credited William Mansfield Clark and his colleagues for developing methods for measuring pH in the 1910s, which subsequently had a wide impact on laboratory and industrial use. In her memoirs, she does not mention how much or how little Clarke and his colleagues knew of Sorensen's work in the years prior. Already at that time, scientists were actively studying the issue of acidity / alkalinity of the environment.
Influence of acid
Dr. Clark's attention was directed to the effect of acid on bacterial growth. And thanks to this, he supplemented the idea of the then science of the hydrogen index of the acidity of the environment. He found that it was the intensity of the acid in terms of the concentration of hydrogen ions that affected their growth. But existing methods for measuring the acidity of a medium determined the amount, not the intensity of the acid. Then, with his colleagues, Dr. Clark developed precise methods for measuring the concentration of hydrogen ions. These methods have replaced the imprecise titration method for acid determination in biological laboratories around the world. It has also been found that they can be used in many industrial and other processes in which they are widely used.
The first electronic pH measurement method was invented by Arnold Orville Beckman, a professor at the California Institute of Technology, in 1934. It was at this point that the local citrus growerSunkist wanted a better method for quickly testing the pH of the lemons they harvested from nearby orchards. The influence of the acidity of the medium was always taken into account.
For example, for a solution with a hydrogen ion activity of 5 × 10–6 (at this level, this is, in fact, 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 5.3. It is believed 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, the amount of pure 107 moles (pH 7) of water contains about 1 g of dissociated hydrogen ions (or, more precisely, 19 g of H3O + hydronium ions) and 17 g hydroxide ions.
The role of temperature
Note that pH is temperature dependent. For example, at 0 °C the pH of pure water is 7.47. At 25 °C it is 7, and at 100 °C it is 6.14.
Electrode potential is proportional to pH when pH is defined in terms of activity. Accurate pH measurement is presented in the international standard ISO 31-8.
A galvanic cell is configured to measure the electromotive force (EMF) between the reference electrode and the hydrogen ion activity sensing electrode when both are immersed in the same aqueous solution. The reference electrode may be a silver chloride object or a calomel electrode. A hydrogen ion selective electrode is standard for these applications.
To put this process into practice, a glass electrode is used instead of a bulky hydrogen electrode. Hehas a built-in reference electrode. It is also calibrated against buffer solutions with known hydrogen ion activity. IUPAC suggested using a set of buffer solutions with known H+ activity. Two or more buffer solutions are used to account for the fact that the slope may be slightly less than ideal. To implement this calibration approach, the electrode is first immersed in a standard solution and the pH meter reading is set to the value of the standard buffer.
The reading from the second standard buffer solution is then corrected using slope control to be equal to the pH level for that solution. When more than two buffer solutions are used, the electrode is calibrated by fitting the observed pH values to a straight line against standard buffer values. Commercial standard buffer solutions are usually supplied with information about the value at 25 °C and the correction factor to be applied for other temperatures.
The pH scale is logarithmic and, therefore, pH is a dimensionless quantity, often used, among other things, to measure the acidity of the internal environment of the cell. This was Sorensen's original definition, which was replaced in 1909.
However, it is possible to directly measure the hydrogen ion concentration if the electrode is calibrated in terms of hydrogen ion concentrations. One way to do this, which has been widely used, is to titrate a solution of known concentrationstrong acid with a solution of a known concentration of a strong alkali in the presence of a relatively high concentration of a supporting electrolyte. Since the acid and alkali concentrations are known, it is easy to calculate the hydrogen ion concentration so that the potential can be related to the measured value.
Indicators can be used to measure pH using the fact that their color changes. Visual comparison of the color of the test solution with a standard color scale allows pH to be measured with integer accuracy. More accurate measurements are possible if the color is measured spectrophotometrically using a colorimeter or spectrophotometer. The universal indicator is made up of a mixture of indicators so that there is a permanent color change from about pH 2 to pH 10. Universal indicator paper is made from absorbent paper that has been impregnated with a universal indicator. Another method for measuring pH is to use an electronic pH meter.
Measuring pH below about 2.5 (about 0.003 moles of acid) and above about 10.5 (about 0.0003 moles of alkali) requires special procedures because Nernst's law is violated at such values when using a glass electrode. Various factors contribute to this. It cannot be assumed that liquid transition potentials are independent of pH. Also, extreme pH means that the solution is concentrated, so the electrode potentials are affected by the change in ionic strength. At high pH, the glass electrode may besubject to alkaline error as the electrode becomes sensitive to the concentration of cations such as Na+ and K+ in solution. Specially designed electrodes are available that partially overcome these problems.
Runoff from mines or mine waste can result in very low pH values.
Pure water is neutral. It is not acidic. When the acid dissolves in water, the pH will be below 7 (25°C). When an alkali dissolves in water, the pH will be greater than 7. A 1 mol solution of a strong acid such as hydrochloric acid has a pH of zero. A solution of a strong alkali such as sodium hydroxide at a concentration of 1 mol has a pH of 14. Thus, measured pH values will generally lie in the range of 0 to 14, although negative pH values and values above 14 are quite possible.
Much depends on the acidity of the solution medium. Because pH is a logarithmic scale, a difference of one pH unit is equivalent to ten times the difference in hydrogen ion concentration. Neutrality PH does not quite reach 7 (at 25 °C), although in most cases this is a good approximation. Neutrality is defined as the condition in which [H+]=[OH-]. Since the self-ionization of water keeps the product of these concentrations [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, so the ph value also matters, which should be neutrallevel. Pure water and a solution of NaCl in pure water are neutral because the dissociation of water produces the same amount of both ions. However, the pH of a neutral NaCl solution will be slightly different from the pH of neutral pure water, since the activity of hydrogen and hydroxide ions depends on ionic strength, so Kw varies with ionic strength.
Dependent plant pigments that can be used as pH indicators are found in many plants, including hibiscus, red cabbage (anthocyanin), and red wine. Citrus juice is acidic because it contains citric acid. Other carboxylic acids are found in many living systems. For example, lactic acid is produced by muscle activity. The state of protonation of phosphate derivatives, such as ATP, depends on the acidity of the pH medium. The functioning of the hemoglobin oxygen transfer enzyme is affected by pH in a process known as the root effect.
In seawater, pH is typically limited to between 7.5 and 8.4. It plays an important role in the carbon cycle in the ocean, and there is evidence of ongoing ocean acidification caused by carbon dioxide emissions. However, measuring pH is complicated by the chemical properties of sea water, and there are several different pH scales in chemical oceanography.
As part of the operational definition of the acidity (pH) scale, IUPAC defines a series of buffer solutions in the pH range (often referred to asNBS or NIST). These solutions have a relatively low ionic strength (≈0.1) compared to seawater (≈0.7) and as a result are not recommended for use in seawater pH characterization because differences in ionic strength cause changes in electrode potential. To solve this problem, an alternative series of buffers based on artificial sea water has been developed.
This new series solves the problem of ionic strength differences between samples and buffers, and the new pH scale for medium acidity is called the common scale, often referred to as pH. The overall scale was determined using a medium containing sulfate ions. These ions experience protonation, H+ + SO2-4 ⇌ HSO-4, so the total scale includes the influence of both protons (free hydrogen ions) and hydrogen sulfide ions:
[H+] T=[H+] F + [HSO-4].
The alternative free scale, often referred to as the pHF, omits this consideration and focuses exclusively on [H+]F, making it in principle a simpler representation of hydrogen ion concentration. Only [H+] T can be determined, so [H+] F should be estimated using [SO2-4] and the stability constant HSO-4, KS:
[H +] F=[H+] T - [HSO-4]=[H+] T (1 + [SO2-4] / K S) -1.
However, it is difficult to estimate KS in sea water, limiting the usefulness of a simpler free scale.
Another scale, known as the seawater scale, often referred to as pHSWS, takes into account the further proton bonding between hydrogen ions and fluoride ions, H+ + F- ⇌HF. The result is the following expression for [H+] SWS:
[H+] SWS=[H+] F + [HSO-4] + [HF]
However, the benefit of considering this additional complexity depends on the fluorine content of the medium. For example, in sea water, sulfate ions are found in much higher concentrations (> 400 times) than the concentrations of fluorine. As a consequence, for most practical purposes, the difference between the common scale and the seawater scale is very small.
The following three equations summarize the three pH scales:
pHF=- log [H+] FpHT=- log ([H+] F + [HSO-4])=- log [H+] TpHSWS=- log ([H+] F + [HSO-4] + [HF])=- log [H+]
From a practical point of view, the three pH scales of an acidic environment (or seawater) differ in their values up to 0.12 pH units, and the differences are much larger than is usually required for accurate pH measurements, in particular in relation to the carbonate system ocean.