pH
In chemistry, pH (/piːˈeɪtʃ/ pee-AYCH), also referred to as acidity or basicity, historically denotes "potential of hydrogen" (or "power of hydrogen").[1] It is a logarithmic scale used to specify the acidity or basicity of aqueous solutions. Acidic solutions (solutions with higher concentrations of hydrogen (H+) ions) are measured to have lower pH values than basic or alkaline solutions.
The pH scale is logarithmic and inversely indicates the activity of hydrogen ions in the solution
where [H+] is the
The pH scale is
History
In 1909, the
In modern chemistry, the p stands for "the negative decimal logarithm of", and is used in the term pKa for acid dissociation constants,[10] so pH is "the negative decimal logarithm of H+ ion concentration", while pOH is "the negative decimal logarithm of OH- ion concentration".
Bacteriologist
Definition
pH
The pH of a solution is defined as the decimal
For example, for a solution with a hydrogen ion activity of 5×10−6 (i.e., the concentration of hydrogen ions in moles per litre), the pH of the solution can be calculated as follows:
The concept of pH was developed because ion-selective electrodes, which are used to measure pH, respond to activity. The electrode potential, E, follows the Nernst equation for the hydrogen ion, which can be expressed as:
where E is a measured potential, E0 is the standard electrode potential, R is the gas constant, T is the temperature in Kelvin, F is the Faraday constant. For H+, the number of electrons transferred is one. The electrode potential is proportional to pH when pH is defined in terms of activity.
The precise measurement of pH is presented in International Standard ISO 31-8 as follows:[14] 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, and 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 electromotive force, ES, is measured. Then the electromotive force, EX, of the same cell containing the solution of unknown pH is measured.
The difference between the two measured electromotive force 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".
In practice, a
The pH scale is logarithmic and therefore pH is a dimensionless quantity.[16]
p[H]
This was the original definition of Sørensen in 1909,[17] which was superseded in favor of pH in 1924. [H] is the concentration of hydrogen ions, denoted [H+] in modern chemistry. 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.
It is possible to measure the concentration of hydrogen ions directly using an electrode calibrated in terms of hydrogen ion concentrations. One common method is to titrate a solution of known concentration of a strong acid with a solution of known concentration of strong base in the presence of a relatively high concentration of background electrolyte. By knowing the concentrations of the acid and base, the concentration of hydrogen ions can be calculated and the measured potential can be correlated with concentrations. The calibration is usually carried out using a Gran plot.[18] This procedure makes the activity of hydrogen ions equal to the numerical value of concentration.
The glass electrode (and other
The difference between p[H] and pH is quite small, and it has been stated that pH = p[H] + 0.04.[19] However, 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−. 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-ionization 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
Measurement
pH Indicators
Average pH of common solutions | ||
---|---|---|
Substance | pH range | Type |
Battery acid | < 1 | Acid |
Gastric acid | 1.0 – 1.5 | |
Orange juice | 3.3 – 4.2 | |
Vinegar | 4-5 | |
Black coffee | 5 – 5.03 | |
Milk | 6.5 – 6.8 | |
Pure water at 25 °C
|
7 | Neutral |
Sea water
|
7.5 – 8.4 | Base |
Ammonia | 11.0 – 11.5 | |
Bleach | 12.5 | |
1 M NaOH | 14 |
pH can be measured using indicators, which change color depending on the pH of the solution they are in. By comparing the color of a test solution to a standard color chart, the pH can be estimated to the nearest whole number. For more precise measurements, the color can be measured using a
Non-aqueous solutions
pH values can be measured in non-aqueous solutions, but they are based on a different scale from aqueous pH values, because the
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 because of differences in the solvated proton ions, such as lyonium ions, which require an intersolvent scale which involves the transfer activity coefficient of hydronium/lyonium ion.
pH is an example of an acidity function, but there are others that can be defined. For example, the Hammett acidity function, H0, has been developed in connection with Superacids.
Unified absolute pH scale
In 2010, a new approach to measuring pH was proposed, called the "unified absolute pH scale". This approach allows for a common reference standard to be used across different solutions, regardless of their pH range. The unified absolute pH scale is based on the absolute chemical potential of the proton, as defined by the Lewis acid–base theory. This scale is applicable to liquids, gases, and even solids.[22] The advantages of the unified absolute pH scale include consistency, accuracy, and applicability to a wide range of sample types. It is precise and versatile because it serves as a common reference standard for pH measurements. However, implementation efforts, compatibility with existing data, complexity, and potential costs are some challenges.
Extremes of pH measurements
The measurement of pH can become difficult at extremely acidic or alkaline conditions, such as below pH 2.5 (ca. 0.003 mol/dm3 acid) or above pH 10.5 (above ca. 0.0003 mol/dm3 alkaline). This is due to the breakdown of the Nernst equation in such conditions when using a glass electrode. There are several factors that contribute to this problem. Firstly, liquid junction potentials may not be independent of pH.[23] Secondly, the high ionic strength of concentrated solutions can affect the electrode potentials. 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.[24] To overcome these problems, specially constructed electrodes are available.
Runoff from mines or mine tailings can produce some extremely low pH values.[25]
Applications
Pure water has a pH of 7 at 25°C, meaning it is neutral. When an acid is dissolved in water, the pH will be less than 7, while a base, or alkali, will have a pH greater than 7. A strong acid, such as hydrochloric acid, at concentration 1 mol dm−3 has a pH of 0, while a strong alkali like sodium hydroxide, at the same concentration, has a pH of 14. Since pH is a logarithmic scale, a difference of one in pH is equivalent to a tenfold difference in hydrogen ion concentration.
Neutrality is not exactly 7 at 25°C, but 7 serves as a good approximation in most cases. Neutrality occurs when the concentration of hydrogen ions ([H+]) equals the concentration of hydroxide ions ([OH−]), or when their 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.
When 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).
- CO
2+ H
2O ⇌ HCO−
3+ H+
pH in soil
The United States Department of Agriculture Natural Resources Conservation Service, formerly Soil Conservation Service classifies soil pH ranges as follows:[26]
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 | 9.0–10.5 |
Hyper alkaline | > 10.5 |
In Europe, topsoil pH is influenced by soil parent material, erosional effects, climate and vegetation. A recent map[27] of topsoil pH in Europe shows the alkaline soils in Mediterranean, Hungary, East Romania, North France. Scandinavian countries, Portugal, Poland and North Germany have more acid soils.
pH in plants
Plants contain pH-dependent
pH in the ocean
The pH of
3) and a hydrogen ion
Three pH scales in oceanography
The measurement of pH in seawater is complicated by the
As part of its
4↔ HSO−
4, such that the total scale includes the effect of both protons
- [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,[31] 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 summarize the three scales of pH:
- pHF = −log10[H+]F
- pHT = −log10([H+]F + [HSO−
4]) = −log10[H+]T - pHSWS = −log10(H+]F + [HSO−
4] + [HF]) = −log10[v]SWS
pH of various body fluids
pH of various body fluids[32] Compartment pH Gastric acid 1.5–3.5[33][34] Lysosomes 4.5[32] Human skin 4.7[35] Granules of chromaffin cells 5.5 Urine 6.0 Breast milk 7.0-7.45[36] Cytosol 7.2 Blood (natural pH) 7.34–7.45[32] Cerebrospinal fluid (CSF) 7.5 Mitochondrial matrix 7.5 Pancreas secretions 8.1
In living organisms, the pH of various Body fluids, cellular compartments, and organs is tightly regulated to maintain a state of acid-base balance known as acid–base homeostasis. Acidosis, defined by blood pH below 7.35, is the most common disorder of acid–base homeostasis and occurs when there is an excess of acid in the body. In contrast, alkalosis is characterized by excessively high blood pH.
Blood pH is usually slightly basic, with a pH of 7.365, referred to as physiological pH in biology and medicine. Plaque formation in teeth can create a local acidic environment that results in tooth decay through demineralization. Enzymes and other Proteins have an optimal pH range for function and can become inactivated or denatured outside this range.
pH calculations
When calculating the pH of a solution containing acids and/or bases, a
Water itself is a weak acid and a weak base, so its dissociation must be taken into account at high pH and low solute concentration (see amphoterism). It dissociates according to the equilibrium
- 2 H2O ⇌ H3O+ (aq) + OH− (aq)
with a dissociation constant, Kw defined as
where [H+] stands for the concentration of the aqueous
Strong acids and bases
However, self-ionization of water must also be considered when concentrations of a strong acid or base is very low or high. For instance, a 5×10−8M solution of HCl would be expected to have a pH of 7.3 based on the above procedure, which is incorrect as it is acidic and should have a pH of less than 7. In such cases, the system can be treated as a mixture of the acid or base and water, which is an
Weak acids and bases
A
- Acid HA: HA ⇌ H+ + A−
- Base A: HA+ ⇌ H+ + A
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
C stands for
Solution of this
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 the addition of hydroxide reduces the hydrogen ion concentration, and the hydroxide ion concentration is constrained by the self-ionization equilibrium to be equal to , the resulting equation is:
General method
Some systems, such as with
Next, write down the mass-balance equations for each reagent:
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 three
See also
- pH indicator
- Arterial blood gas
- Chemical equilibrium
- pCO2
- pKa
References
- doi:10.1021/ed081p21. Archived(PDF) from the original on 14 December 2019. Retrieved 15 July 2020.
- .
- ^ (PDF) from the original on 24 September 2007.
- ^ a b Sørensen, S. P. L. (1909). "Über die Messung und die Bedeutung der Wasserstoffionenkonzentration bei enzymatischen Prozessen" (PDF). Biochem. Z. 21: 131–304. Archived (PDF) from the original on 15 April 2021. Retrieved 22 March 2021.
Original German: Für die Zahl p schlage ich den Namen Wasserstoffionenexponent und die Schreibweise pH• vor. Unter dem Wasserstoffionexponenten (pH•) einer Lösungwird dann der Briggsche Logarithmus des reziproken Wertes des auf Wasserstoffionenbezagenen Normalitäts faktors de Lösungverstanden.
Two other publications appeared in 1909, one in French and one in Danish. - from the original on 6 August 2020. Retrieved 21 July 2019.
- .
- S2CID 110716297.
- S2CID 104410918. Retrieved 16 June 2022.
- ^ Bradley, David (21 February 2018). "When it comes to caustic wit and an acid tongue, mind your Ps and Qs". Materials Today. Retrieved 16 June 2022.
- PMID 10637613.
- ^ a b Evans, Alice C. (1963). "Memoirs" (PDF). NIH Office of History. National Institutes of Health Office of History. Archived from the original (PDF) on 15 December 2017. Retrieved 27 March 2018.
- ^ "Origins: Birth of the pH Meter". Caltech Engineering & Science Magazine. Archived from the original on 6 November 2018. Retrieved 11 March 2018.
- ^ Tetrault, Sharon (June 2002). "The Beckmans". Orange Coast. Orange Coast Magazine. Archived from the original on 15 April 2021. Retrieved 11 March 2018.
- ^ Quantities and units – Part 8: Physical chemistry and molecular physics, Annex C (normative): pH. International Organization for Standardization, 1992.
- (PDF) from the original on 24 September 2007.
- .
- ^ "Carlsberg Group Company History Page". Carlsberggroup.com. Archived from the original on 18 January 2014. Retrieved 7 May 2013.
- .
- ISBN 0-582-22628-7, Section 13.23, "Determination of pH"
- ISBN 0-632-03583-8. pp. 49–50. Electronic version.
- PMID 20715223.
- .
- ISBN 0-582-22628-7, Section 13.19 The glass electrode
- from the original on 23 September 2017. Retrieved 4 November 2018.
- ^ 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.
- PMID 31798185.
- S2CID 255431338.
- ^ ISBN 0-444-50946-1
- .
- .
- ^ OCLC 1017876653. Archived from the original on 8 May 2022. Retrieved 8 May 2022.)
{{cite book}}
: CS1 maint: multiple names: authors list (link - ^ "Stomach acid test". University of California San Francisco. Retrieved 21 February 2024.
- ISBN 9780321735287. Archivedfrom the original on 8 May 2022. Retrieved 8 May 2022.
- from the original on 21 March 2022. Retrieved 8 May 2022.
- PMID 3748680.
- ^ 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.
- ISBN 978-0-470-38123-6.
External links
- pH value (P9440) (see uses)