pH

From Wikipedia, the free encyclopedia

Lemon juice tastes sour because it contains 5% to 6% citric acid, which has a pH of 2.2.

In chemistry, pH (/piːˈeɪtʃ/) is a numeric scale used to specify the acidity or alkalinity of an aqueous solution. It is the negative of the logarithm to base 10 of the activity of the hydrogen ion. Solutions with a pH less than 7 are acidic and solutions with a pH greater than 7 are alkaline or basic. Pure water is neutral, being neither an acid nor a base. Contrary to popular belief, the pH value can be less than 0 or greater than 14 for very strong acids and bases respectively.^[1]
pH measurements are important in medicine, biology, chemistry, agriculture, forestry, food science, environmental science, oceanography, civil engineering, chemical engineering, nutrition, water treatment & water purification, and many other applications.
The pH scale is traceable to a set of standard solutions whose pH is established by international agreement.^[2] Primary pH standard values are determined using a concentration cell with transference, by measuring the potential difference between a hydrogen electrode and a standard electrode such as the silver chloride electrode. The pH of aqueous solutions can be measured with a glass electrode and a pH meter, or indicator.
pH is the negative of the logarithm to base 10 of the activity of the (solvated) hydronium ion, more often (albeit somewhat inaccurately) expressed as the measure of the hydronium ion concentration.^[3]
The rest of this article uses the technically correct word "base" and its inflections in place of "alkaline", which specifically refers to a base dissolved in water^{[citation needed]}, and its inflections.

History

The concept of p[H] was first introduced by Danish chemist Søren Peder Lauritz Sørensen at the Carlsberg Laboratory in 1909^[4] and revised to the modern pH in 1924 to accommodate definitions and measurements in terms of electrochemical cells. In the first papers, the notation had the "H" as a subscript to the lowercase "p", as so: p_H.

The exact meaning of the "p" in "pH" is disputed, but according to the Carlsberg Foundation pH stands for "power of hydrogen".^[5] It has also been suggested that the "p" stands for the German Potenz (meaning "power"), others refer to French puissance (also meaning "power", based on the fact that the Carlsberg Laboratory was French-speaking). Another suggestion is that the "p" stands for the Latin terms pondus hydrogenii (engl. quantity of hydrogen), potentia hydrogenii (engl. capacity of hydrogen), or potential hydrogen. It is also suggested that Sørensen used the letters "p" and "q" (commonly paired letters in mathematics) simply to label the test solution (p) and the reference solution (q).^[6] Current use in chemistry is that p stands for "decimal cologarithm of", as also in the term pK_a, used for acid dissociation constants.^[7]

Definition and measurement

pH

pH is defined as the decimal logarithm of the reciprocal of the hydrogen ion activity, a_H+, in a solution.^[2]

This definition was adopted because ion-selective electrodes, which are used to measure pH, respond to activity. Ideally, electrode potential, E, follows the Nernst equation, which, for the hydrogen ion can be written 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⁺ number of electrons transferred is one. It follows that electrode potential is proportional to pH when pH is defined in terms of activity. Precise measurement of pH is presented in International Standard ISO 31-8 as follows:^[8] 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. The hydrogen-ion selective electrode is a standard hydrogen electrode.

Reference electrode | concentrated solution of KCl || test solution | H₂ | Pt^{[clarification needed]}

Firstly, the cell is filled with a solution of known hydrogen ion activity and the emf, E_S, is measured. Then the emf, E_X, of the same cell containing the solution of unknown pH is measured.

The difference between the two measured emf 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".

To apply this process in practice, a glass electrode is used rather than the cumbersome hydrogen electrode. A combined glass electrode has an in-built reference electrode. It is calibrated against buffer solutions of known hydrogen ion activity. IUPAC has proposed the use of a set of buffer solutions of known H⁺ activity.^[2] Two or more buffer solutions are used in order to accommodate the fact that the "slope" may differ slightly from ideal. To implement this approach to calibration, the electrode is first immersed in a standard solution and the reading on a pH meter is adjusted to be equal to the standard buffer's value. The reading from a second standard buffer solution is then adjusted, using the "slope" control, to be equal to the pH for that solution. Further details, are given in the IUPAC recommendations.^[2] When more than two buffer solutions are used the electrode is calibrated by fitting observed pH values to a straight line with respect to standard buffer values. Commercial standard buffer solutions usually come with information on the value at 25 °C and a correction factor to be applied for other temperatures.

The pH scale is logarithmic and therefore pH is a dimensionless quantity.

p[H]

This was the original definition of Sørensen,^[5] which was superseded in favor of pH in 1924. 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 alkaline in the presence of a relatively high concentration of background electrolyte. Since the concentrations of acid and alkaline 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.^[9] The calibration yields 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, f, 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 glass electrode (and other ion selective electrodes) 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.

The difference between p[H] and pH is quite small. It has been stated^[10] that pH = p[H] + 0.04. It is common practice to use the term "pH" for both types of measurement.

pH indicators

Chart showing the variation of color of universal indicator paper with pH

Indicators may be used to measure pH, by making use of the fact that their color changes with pH. Visual comparison of the color of a test solution with a standard color chart provides a means to measure pH accurate to the nearest whole number. More precise measurements are possible if the color is measured spectrophotometrically, using a colorimeter of spectrophotometer. Universal indicator consists of a mixture of indicators such that there is a continuous color change from about pH 2 to pH 10. Universal indicator paper is made from absorbent paper that has been impregnated with universal indicator.

pOH

Relation between p[OH] and p[H] (red = acid region, blue = basic region)

pOH is sometimes used as a measure of the concentration of hydroxide ions, OH⁻, or alkalinity. pOH values are derived from pH measurements. The concentration of hydroxide ions in water is related to the concentration of hydrogen ions by

where K_W is the self-ionisation 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 soil alkalinity.

Extremes of pH

Measurement of pH below about 2.5 (ca. 0.003 mol dm⁻³ acid) and above about 10.5 (ca. 0.0003 mol dm⁻³ alkaline) requires special procedures because, when using the glass electrode, the Nernst law breaks down under those conditions. Various factors contribute to this. It cannot be assumed that liquid junction potentials are independent of pH.^[11] Also, extreme pH implies that the solution is concentrated, so electrode potentials are affected by ionic strength variation. 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.^[12] Specially constructed electrodes are available which partly overcome these problems.

Runoff from mines or mine tailings can produce some very low pH values.^[13]

Non-aqueous solutions

Hydrogen ion concentrations (activities) can be measured in non-aqueous solvents. pH values based on these measurements belong to a different scale from aqueous pH values, because activities relate to different standard states. Hydrogen ion activity, a_H⁺, can be defined^[14][15] as:

where μ_H⁺ is the chemical potential of the hydrogen ion, μ^o_H⁺ 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, requiring an intersolvent scale which involves the transfer activity coefficient of hydrolyonium ion.

pH is an example of an acidity function. Other acidity functions can be defined. For example, the Hammett acidity function, H₀, has been developed in connection with superacids.

The concept of "Unified pH scale" has been developed on the basis of the absolute chemical potential of the proton. This scale applies to liquids, gases and even solids.^[16]

Applications

pH values of some common substances

Pure water is neutral. When an acid is dissolved in water, the pH will be less than 7. When a base, or alkaline, 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 dm⁻³ has a pH of 0. A solution of a strong alkaline, such as sodium hydroxide, at concentration 1 mol dm⁻³, has a pH of 14. Thus, measured pH values will lie mostly in the range 0 to 14, though negative pH values and values above 14 are entirely possible. Since pH is a logarithmic scale, a difference of one pH unit is equivalent to a tenfold difference in hydrogen ion concentration. The pH of an aqueous solution of a salt such as sodium chloride is slightly different from that of a neutral solution, even though the salt is neither acidic nor basic. This is because the hydrogen and hydroxide ions' activity is dependent on ionic strength, so K_w varies with ionic strength.

Pure water does not contain any ions and therefore cannot have a pH value (log(0) is infinity). however, if 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 carbonate and hydrogen ions (essentially creating carbonic acid).

pH in nature

pH-dependent plant pigments that can be used as pH indicators occur in many plants, including hibiscus, red cabbage (anthocyanin) and red wine. The juice of citrus fruits is acidic mainly because it contains citric acid. Other carboxylic acids occur in many living systems. For example, lactic acid is produced by muscle activity. The state of protonation of phosphate derivatives, such as ATP, is pH-dependent. The functioning of the oxygen-transport enzyme hemoglobin is affected by pH in a process known as the Root effect.

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.^[17] However, pH measurement is complicated by the chemical properties of seawater, and several distinct pH scales exist in chemical oceanography.^[18]

As part of its operational definition of the pH scale, the IUPAC defines 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, as a consequence, are not recommended for use in characterizing 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.^[19] 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 pH_T. 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 'pH_F', 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,^[20] therefore [H⁺]_F must be estimated using the [SO₄²⁻] and the stability constant of HSO₄⁻, K_S^*:

[H⁺]_F = [H⁺]_T − [HSO₄⁻] = [H⁺]_T ( 1 + [SO₄²⁻] / 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 'pH_SWS', 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. As a consequence, 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:

pH_F = − log [H⁺]_F

pH_T = − log ( [H⁺]_F + [HSO₄⁻] ) = − log [H⁺]_T

pH_SWS = − 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, in particular, in relation to the ocean's carbonate system.^[18] 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^[21]

Compartment	pH
Gastric acid	1
Lysosomes	4.5
Granules of chromaffin cells	5.5
Human skin	5.5
Urine	6.0
Cytosol	7.2
Cerebrospinal fluid (CSF)	7.5
Blood	7.34–7.45
Mitochondrial matrix	7.5
Pancreas secretions	8.1

The pH of different cellular compartments, body fluids, and organs is usually tightly regulated in a process called acid-base homeostasis. 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. Alkalosis is the opposite condition, with blood pH being excessively high.

The pH of blood is usually slightly basic with a value of pH 7.365. 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 demineralization. Enzymes and other proteins have an optimum pH range and can become inactivated or denatured outside this range.

Calculations of pH

The calculation of the pH of a solution containing acids and/or bases is an example of a chemical speciation calculation, that is, a mathematical procedure for calculating the concentrations of all chemical species that are present in the solution. The complexity of the procedure depends on the nature of the solution. For strong acids and bases no calculations are necessary except in extreme situations. The pH of a solution containing a weak acid requires the solution of a quadratic equation. The pH of a solution containing a weak base may require the solution of a cubic equation. The general case requires the solution of a set of non-linear simultaneous equations.

A complicating factor is that water itself is a weak acid and a weak base. It dissociates according to the equilibrium

with a dissociation constant, K_w defined as

where [H⁺] stands for the concentration of the aquated hydronium ion and [OH⁻] represents the concentration of the hydroxide ion. This equilibrium needs to be taken into account at high pH and when the solute concentration is extremely low.

Strong acids and bases

Strong acids and bases are compounds that, for practical purposes, are completely dissociated in water. Under normal circumstances this means that the concentration of hydrogen ions in acidic solution can be taken to be equal to the concentration of the acid. The pH is then equal to minus the logarithm of the concentration value. Hydrochloric acid (HCl) is an example of a strong acid. The pH of a 0.01M solution of HCl is equal to −log₁₀(0.01), that is, pH = 2. Sodium hydroxide, NaOH, is an example of a strong base. The p[OH] value of a 0.01M solution of NaOH is equal to −log₁₀(0.01), that is, p[OH] = 2. From the definition of p[OH] above, this means that the pH is equal to about 12. For solutions of sodium hydroxide at higher concentrations the self-ionization equilibrium must be taken into account.

Self-ionization must also be considered when concentrations are extremely low. Consider, for example, a solution of hydrochloric acid at a concentration of 5×10⁻⁸M. The simple procedure given above would suggest that it has a pH of 7.3. This is clearly wrong as an acid solution should have a pH of less than 7. Treating the system as a mixture of hydrochloric acid and the amphoteric substance water, a pH of 6.89 results.^[22]

Weak acids and bases

A weak acid or the conjugate acid of a weak base can be treated using the same formalism.

Acid:

Base:

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 mass conservation in terms of the two "reagents" H and A.

C stands for analytical concentration. In some texts one mass balance equation is replaced by an equation of charge balance. This is satisfactory for simple cases like this one, but is more difficult to apply to more complicated cases as those below. Together with the equation defining K_a, there are now three equations in three unknowns. When an acid is dissolved in water C_A = C_H = C_a, the concentration of the acid, so [A] = [H]. After some further algebraic manipulation an equation in the hydrogen ion concentration may be obtained.

Solution of this quadratic equation gives the hydrogen ion concentration and hence p[H] or, more loosely, pH. This procedure is illustrated in an ICE table which can also be used to calculate the pH when some additional (strong) acid or alkaline has been added to the system, that is, when C_A ≠ C_H.
For example, what is the pH of a 0.01M solution of benzoic acid, pK_a = 4.19?
Step 1:
Step 2: Set up the quadratic equation.
Step 3: Solve the quadratic equation.  ; pH = 3.11

For alkaline solutions an additional term is added to the mass-balance equation for hydrogen. Since addition of hydroxide reduces the hydrogen ion concentration, and the hydroxide ion concentration is constrained by the self-ionization equilibrium to be equal to

In this case the resulting equation in [H] is a cubic equation.

General method

Some systems, such as with polyprotic acids, are amenable to spreadsheet calculations.^[23] With three or more reagents or when many complexes are formed with general formulae such as A_pB_qH_r the following general method can be used to calculate the pH of a solution. For example, with three reagents, each equilibrium is characterized by and equilibrium constant, β.

Next, write down the mass-balance equations for each reagent:

Note that 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 3 non-linear simultaneous equations in the three unknowns, [A], [B] and [H]. Because the equations are non-linear, and because concentrations may range over many powers of 10, the solution of these equations is not straightforward. However, many computer programs are available which can be used to perform these calculations. There may be more than three reagents. The calculation of hydrogen ion concentrations, using this formalism, is a key element in the determination of equilibrium constants by potentiometric titration.

Carbonate

From Wikipedia, the free encyclopedia

Carbonate

Names
Systematic IUPAC name Carbonate
Identifiers
ChemSpider	18519
InChI[show] InChI=1S/CH2O3/c2-1(3)4/h(H2,2,3,4)/p-2 Key: BVKZGUZCCUSVTD-UHFFFAOYSA-L InChI=1/CH2O3/c2-1(3)4/h(H2,2,3,4)/p-2 Key: BVKZGUZCCUSVTD-NUQVWONBAE
Jmol-3D images	Image
PubChem	19660
SMILES[show] C(=O)([O-])[O-]
Properties
Chemical formula	CO2− 3
Molar mass	60.01 g·mol−1
Except where otherwise noted, data are given for materials in their standard state (at 25 °C [77 °F], 100 kPa).
Infobox references

In chemistry, a carbonate is a salt of carbonic acid, characterized by the presence of the carbonate ion, CO2−
3. The name may also mean an ester of carbonic acid, an organic compound containing the carbonate group C(=O)(O–)₂.

The term is also used as a verb, to describe carbonation: the process of raising the concentrations of carbonate and bicarbonate ions in water to produce carbonated water and other carbonated beverages — either by the addition of carbon dioxide gas under pressure, or by dissolving carbonate or bicarbonate salts into the water.

In geology and mineralogy, the term "carbonate" can refer both to carbonate minerals and carbonate rock (which is made of chiefly carbonate minerals), and both are dominated by the carbonate ion, CO2−
3. Carbonate minerals are extremely varied and ubiquitous in chemically precipitated sedimentary rock. The most common are calcite or calcium carbonate, CaCO₃, the chief constituent of limestone (as well as the main component of mollusc shells and coral skeletons); dolomite, a calcium-magnesium carbonate CaMg(CO₃)₂; and siderite, or iron(II) carbonate, FeCO₃, an important iron ore. Sodium carbonate ("soda" or "natron") and potassium carbonate ("potash") have been used since antiquity for cleaning and preservation, as well as for the manufacture of glass. Carbonates are widely used in industry, e.g. in iron smelting, as a raw material for Portland cement and lime manufacture, in the composition of ceramic glazes, and more.

Structure and bonding

The carbonate ion is the simplest oxocarbon anion. It consists of one carbon atom surrounded by three oxygen atoms, in a trigonal planar arrangement, with D_3h molecular symmetry. It has a molecular mass of 60.01 g/mol and carries a negative two formal charge. It is the conjugate base of the hydrogen carbonate (bicarbonate) ion, HCO₃⁻, which is the conjugate base of H₂CO₃, carbonic acid.

The Lewis structure of the carbonate ion has two (long) single bonds to negative oxygen atoms, and one short double bond to a neutral oxygen.




This structure is incompatible with the observed symmetry of the ion, which implies that the three bonds are equally long and that the three oxygen atoms are equivalent. As in the case of the isoelectronic nitrate ion, the symmetry can be achieved by a resonance between three structures:

This resonance can be summarized by a model with fractional bonds and delocalized charges:

Chemical properties

Metal carbonates generally decompose on heating, liberating carbon dioxide from the long term carbon cycle to the short term carbon cycle and leaving behind an oxide of the metal. This process is called calcination, after calx, the Latin name of quicklime or calcium oxide, CaO, which is obtained by roasting limestone in a lime kiln.

A carbonate salt forms when a positively charged ion, M+, M2+, or M3+, attaches to the negatively charged oxygen atoms of the ion, forming an ionic compound:

2 M+ + CO2−
3 → M
2CO
3

M2+ + CO2−
3 → MCO
3

2 M3+ + 3 CO2−
3 → M
2(CO
3)
3

Most carbonate salts are insoluble in water at standard temperature and pressure, with solubility constants of less than 1×10⁻⁸. Exceptions include lithium, sodium, potassium and ammonium carbonates, as well as many uranium carbonates.

In aqueous solution, carbonate, bicarbonate, carbon dioxide, and carbonic acid exist together in a dynamic equilibrium. In strongly basic conditions, the carbonate ion predominates, while in weakly basic conditions, the bicarbonate ion is prevalent. In more acid conditions, aqueous carbon dioxide, CO
2(aq), is the main form, which, with water, H
2O, is in equilibrium with carbonic acid - the equilibrium lies strongly towards carbon dioxide. Thus sodium carbonate is basic, sodium bicarbonate is weakly basic, while carbon dioxide itself is a weak acid.

Carbonated water is formed by dissolving CO
2 in water under pressure. When the partial pressure of CO
2 is reduced, for example when a can of soda is opened, the equilibrium for each of the forms of carbonate (carbonate, bicarbonate, carbon dioxide, and carbonic acid) shifts until the concentration of CO
2 in the solution is equal to the solubility of CO₂ at that temperature and pressure. In living systems an enzyme, carbonic anhydrase, speeds the interconversion of CO₂ and carbonic acid.

Although the carbonate salts of most metals are insoluble in water, the same is not true of the bicarbonate salts. In solution this equilibrium between carbonate, bicarbonate, carbon dioxide and carbonic acid changes consonant to changing temperature and pressure conditions. In the case of metal ions with insoluble carbonates, e.g. CaCO
3, formation of insoluble compounds results. This is an explanation for the buildup of scale inside pipes caused by hard water.

Organic carbonates

In organic chemistry a carbonate can also refer to a functional group within a larger molecule that contains a carbon atom bound to three oxygen atoms, one of which is double bonded. These compounds are also known as organocarbonates or carbonate esters, and have the general formula ROCOOR′, or RR′CO₃. Important organocarbonates include dimethyl carbonate, the cyclic compounds ethylene carbonate and propylene carbonate, and the phosgene replacement, triphosgene.

Biological significance

It works as a buffer in the blood as follows: when pH is too low, the concentration of hydrogen ions is too high, so one exhales CO₂. This will cause the equation to shift left, essentially decreasing the concentration of H⁺ ions, causing a more basic pH.

When pH is too high, the concentration of hydrogen ions in the blood is too low, so the kidneys excrete bicarbonate (HCO₃⁻). This causes the equation to shift right, essentially increasing the concentration of hydrogen ions, causing a more acidic pH.

There are 3 important reversible reactions that control the above pH balance.^[1]

1. H₂CO₃(aq) H⁺(aq) + HCO₃⁻(aq)
2. H₂CO₃(aq) CO₂(aq) + H₂O(l)
3. CO₂(aq) CO₂(g)

Exhaled CO₂(g) depletes CO₂(aq) which in turn consumes H₂CO₃ causing the aforementioned shift left in the first reaction by Le Chatelier's principle. By the same principle when the pH is too high, the kidneys excrete bicarbonate (HCO₃⁻) into urine as urea via the Urea Cycle (a.k.a. the Krebs-Henseleit Ornithine Cycle). By removing the bicarbonate more H⁺ is generated from carbonic acid (H₂CO₃) which come from CO₂(g) produced by cellular respiration.

Crucially this same buffer operates in the oceans. It is a major factor in climate change and the long term carbon cycle. This is due to the large number of marine organisms (especially coral) which are formed of calcium carbonate. Increased solubility of carbonate through increased temperatures results in lower production of marine calcite and increased concentration of atmospheric carbon dioxide. This in turn increases Earth temperature and is a part of the carbon cycle largely ignored by the global news media. The tonnage of CO₃^2- is on a geological scale and may all be redissolved into the sea and released to the atmosphere, increasing CO₂ levels even more.

Carbonate salts

• Carbonate overview:

Carbonates

H2CO3

He

LiCO3

BeCO3

B

C

(NH4)2CO3, NH4HCO3

O

F

Ne

Na2CO3, NaHCO3, Na3H(CO3)2

MgCO3, Mg(HCO3)2

Al2(CO3)3

Si

P

S

Cl

Ar

K2CO3, KHCO3

CaCO3, Ca(HCO3)2

Sc

Ti

V

Cr

MnCO3

FeCO3

CoCO3

NiCO3

CuCO3

ZnCO3

Ga

Ge

As

Se

Br

Kr

Rb2CO3

SrCO3

Y

Zr

Nb

Mo

Tc

Ru

Rh

Pd

Ag2CO3

CdCO3

In

Sn

Sb

Te

I

Xe

Cs2CO3, CsHCO3

BaCO3

Hf

Ta

W

Re

Os

Ir

Pt

Au

Hg

Tl2CO3

PbCO3

(BiO)2CO3

Po

At

Rn

Fr

Ra

Rf

Db

Sg

Bh

Hs

Mt

Ds

Rg

Cn

Uut

Fl

Uup

Lv

Uus

Uuo

↓

La2(CO3)3

Ce

Pr

Nd

Pm

Sm

Eu

Gd

Tb

Dy

Ho

Er

Tm

Yb

Lu

Ac

Th

Pa

UO2CO3

Np

Pu

Am

Cm

Bk

Cf

Es

Fm

Md

No

Lr

Presence outside Earth

It is generally thought that the presence of carbonates in rock is strong evidence for the presence of liquid water. Recent observations of the Planetary nebula NGC 6302 shows evidence for carbonates in space,^[2] where aqueous alteration similar to that on Earth is unlikely. Other minerals have been proposed which would fit the observations.

Until recently carbonate deposits have not been found on Mars via remote sensing or in situ missions, even though Martian meteorites contain small amounts. Groundwater may have existed at both Gusev^[3] and Meridiani Planum.^[4]

Carbonic acid

From Wikipedia, the free encyclopedia

Carbonic acid

Names
IUPAC name Carbonic acid
Other names Carbon dioxide solution; Dihydrogen carbonate; acid of air; Aerial acid; Hydroxymethanoic acid
Identifiers
CAS Registry Number	463-79-6 Y
ChEBI	CHEBI:28976 Y
ChEMBL	ChEMBL1161632 Y
ChemSpider	747 Y
InChI[show] InChI=1S/CH2O3/c2-1(3)4/h(H2,2,3,4) Y Key: BVKZGUZCCUSVTD-UHFFFAOYSA-N Y InChI=1/CH2O3/c2-1(3)4/h(H2,2,3,4) Key: BVKZGUZCCUSVTD-UHFFFAOYAU
Jmol-3D images	Image
KEGG	C01353 Y
SMILES[show] O=C(O)O
Properties
Chemical formula	H2CO3
Molar mass	62.03 g/mol
Density	1.668 g/cm3
Solubility in water	Exists only in solution
Acidity (pKa)	3.6 (pKa1 for H2CO3 only), 6.3 (pKa1 including CO2(aq)), 10.32 (pKa2)
Except where otherwise noted, data are given for materials in their standard state (at 25 °C [77 °F], 100 kPa).
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Carbonic acid is a chemical compound with the chemical formula H₂CO₃ (equivalently OC(OH)₂). It is also a name sometimes given to solutions of carbon dioxide in water (carbonated water), because such solutions contain small amounts of H₂CO₃. In physiology, carbonic acid is described as volatile acid or respiratory acid, because it is the only acid excreted as a gas by the lungs.^[1]
Carbonic acid, which is a weak acid, forms two kinds of salts, the carbonates and the bicarbonates. In geology, carbonic acid causes limestone to dissolve producing calcium bicarbonate which leads to many limestone features such as stalactites and stalagmites.
The long-held belief that carbonic acid could not exist as a pure compound has reportedly been recently disproved by the preparation of the pure substance in both solid and gas form by University of Innsbruck researchers.^[2]

Chemical equilibrium

When carbon dioxide dissolves in water it exists in chemical equilibrium producing carbonic acid:^[3]

CO₂ + H₂O H₂CO₃

The hydration equilibrium constant at 25 °C is called K_h, which in the case of carbonic acid is [H₂CO₃]/[CO₂] ≈ 1.7×10⁻³ in pure water^[4] and ≈ 1.2×10⁻³ in seawater.^[5] Hence, the majority of the carbon dioxide is not converted into carbonic acid, remaining as CO₂ molecules. In the absence of a catalyst, the equilibrium is reached quite slowly. The rate constants are 0.039 s⁻¹ for the forward reaction (CO₂ + H₂O → H₂CO₃) and 23 s⁻¹ for the reverse reaction (H₂CO₃ → CO₂ + H₂O). Carbonic acid is used in the making of soft drinks, inexpensive and artificially carbonated sparkling wines, and other bubbly drinks. The addition of two molecules of water to CO₂ would give orthocarbonic acid, C(OH)₄, which exists only in minute amounts in aqueous solution.

Addition of base to an excess of carbonic acid gives bicarbonate (hydrogen carbonate). With excess base, carbonic acid reacts to give carbonate salts.

Role of carbonic acid in blood

Carbonic acid is an intermediate step in the transport of CO₂ out of the body via respiratory gas exchange. The hydration reaction of CO₂ is generally very slow in the absence of a catalyst, but red blood cells contain carbonic anhydrase, which both increases the reaction rate and dissociates a hydrogen ion (H⁺) from the resulting carbonic acid, leaving bicarbonate (HCO₃⁻) dissolved in the blood plasma. This catalysed reaction is reversed in the lungs, where it converts the bicarbonate back into CO₂ and allows it to be expelled. This equilibration plays an important role as a buffer in mammalian blood.^[6]

Role of carbonic acid in ocean chemistry

The oceans of the world have absorbed almost half of the CO₂ emitted by humans from the burning of fossil fuels.^[7]  The extra dissolved carbon dioxide has caused the ocean's average surface pH to shift by about 0.1 unit from pre-industrial levels.^[8] This process is known as ocean acidification.^[9]

Acidity of carbonic acid

Carbonic acid is one of the polyprotic acids: It is diprotic - it has two protons, which may dissociate from the parent molecule. Thus, there are two dissociation constants, the first one for the dissociation into the bicarbonate (also called hydrogen carbonate) ion HCO₃⁻:

H₂CO₃ HCO₃⁻ + H⁺

K_a1 = 2.5×10⁻⁴;^[3] pK_a1 = 3.6 at 25 °C.

Care must be taken when quoting and using the first dissociation constant of carbonic acid. In aqueous solution, carbonic acid exists in equilibrium with carbon dioxide, and the concentration of H₂CO₃ is much lower than the concentration of CO₂. In many analyses, H₂CO₃ includes dissolved CO₂ (referred to as CO₂(aq)), H₂CO₃* is used to represent the two species when writing the aqueous chemical equilibrium equation. The equation may be rewritten as follows:^[3]

H₂CO₃* HCO₃⁻ + H⁺

K_a(app) = 4.6×10⁻⁷; pK(app) = 6.3 at 25 °C and ionic strength = 0.0

Whereas this apparent pK_a is quoted as the dissociation constant of carbonic acid, it is ambiguous: it might better be referred to as the acidity constant of dissolved carbon dioxide, as it is particularly useful for calculating the pH of CO₂-containing solutions. A similar situation applies to sulfurous acid (H₂SO₃), which exists in equilibrium with substantial amounts of unhydrated sulfur dioxide.
The second constant is for the dissociation of the bicarbonate ion into the carbonate ion CO₃²⁻:

HCO₃⁻ CO₃²⁻ + H⁺

K_a2 = 4.69×10⁻¹¹; pK_a2 = 10.329 at 25 °C and ionic strength = 0.0

The three acidity constants are defined as follows:

pH and composition of carbonic acid solutions

At a given temperature, the composition of a pure carbonic acid solution (or of a pure CO₂ solution) is completely determined by the partial pressure of carbon dioxide above the solution. To calculate this composition, account must be taken of the above equilibria between the three different carbonate forms (H₂CO₃, HCO₃⁻ and CO₃²⁻) as well as of the hydration equilibrium between dissolved CO₂ and H₂CO₃ with constant (see above) and of the following equilibrium between the dissolved CO₂ and the gaseous CO₂ above the solution:

CO₂(gas) CO₂(dissolved) with where k_H=29.76 atm/(mol/L) at 25 °C (Henry constant)

The corresponding equilibrium equations together with the relation and the charge neutrality condition result in six equations for the six unknowns [CO₂], [H₂CO₃], [H⁺], [OH⁻], [HCO₃⁻] and [CO₃²⁻], showing that the composition of the solution is fully determined by . The equation obtained for [H⁺] is a cubic whose numerical solution yields the following values for the pH and the different species concentrations:

(atm)	pH	[CO2] (mol/L)	[H2CO3] (mol/L)	[HCO3−] (mol/L)	[CO32−] (mol/L)
1.0 × 10−8	7.00	3.36 × 10−10	5.71 × 10−13	1.42 × 10−09	7.90 × 10−13
1.0 × 10−7	6.94	3.36 × 10−09	5.71 × 10−12	5.90 × 10−09	1.90 × 10−12
1.0 × 10−6	6.81	3.36 × 10−08	5.71 × 10−11	9.16 × 10−08	3.30 × 10−11
1.0 × 10−5	6.42	3.36 × 10−07	5.71 × 10−10	3.78 × 10−07	4.53 × 10−11
1.0 × 10−4	5.92	3.36 × 10−06	5.71 × 10−09	1.19 × 10−06	5.57 × 10−11
3.5 × 10−4	5.65	1.18 × 10−05	2.00 × 10−08	2.23 × 10−06	5.60 × 10−11
1.0 × 10−3	5.42	3.36 × 10−05	5.71 × 10−08	3.78 × 10−06	5.61 × 10−11
1.0 × 10−2	4.92	3.36 × 10−04	5.71 × 10−07	1.19 × 10−05	5.61 × 10−11
1.0 × 10−1	4.42	3.36 × 10−03	5.71 × 10−06	3.78 × 10−05	5.61 × 10−11
1.0 × 10+0	3.92	3.36 × 10−02	5.71 × 10−05	1.20 × 10−04	5.61 × 10−11
2.5 × 10+0	3.72	8.40 × 10−02	1.43 × 10−04	1.89 × 10−04	5.61 × 10−11
1.0 × 10+1	3.42	3.36 × 10−01	5.71 × 10−04	3.78 × 10−04	5.61 × 10−11

• We see that in the total range of pressure, the pH is always largely lower than pKa₂ so that the CO₃²⁻ concentration is always negligible with respect to HCO₃⁻ concentration. In fact CO₃²⁻ plays no quantitative role in the present calculation (see remark below).
• For vanishing , the pH is close to the one of pure water (pH = 7) and the dissolved carbon is essentially in the HCO₃⁻ form.
• For normal atmospheric conditions ( atm), we get a slightly acid solution (pH = 5.7) and the dissolved carbon is now essentially in the CO₂ form. From this pressure on, [OH⁻] becomes also negligible so that the ionized part of the solution is now an equimolar mixture of H⁺ and HCO₃⁻.
• For a CO₂ pressure typical of the one in soda drink bottles ( ~ 2.5 atm), we get a relatively acid medium (pH = 3.7) with a high concentration of dissolved CO₂. These features contribute to the sour and sparkling taste of these drinks.
• Between 2.5 and 10 atm, the pH crosses the pKa₁ value (3.60) giving a dominant H₂CO₃ concentration (with respect to HCO₃⁻) at high pressures.
• A plot of the equilibrium concentrations of these different forms of dissolved inorganic carbon (and which species is dominant), as a function of the pH of the solution, is known as a Bjerrum plot.

Remark

As noted above, [CO₃²⁻] may be neglected for this specific problem, resulting in the following very precise analytical expression for [H⁺]:

Spectroscopic studies of carbonic acid

Theoretical calculations show that the presence of even a single molecule of water causes carbonic acid to revert to carbon dioxide and water. In the absence of water, the dissociation of gaseous carbonic acid is predicted to be very slow, with a half-life of 180,000 years.^[10] This may only apply to isolated carbonic acid molecules however, as it has been predicted to catalyze its own decomposition^[11]

It has long been recognized that pure carbonic acid cannot be obtained at room temperatures (about 20 °C or about 70 °F). It can be generated by exposing a frozen mixture of water and carbon dioxide to high-energy radiation, and then warming to remove the excess water. The carbonic acid that remained was characterized by infrared spectroscopy. The fact that the carbonic acid was prepared by irradiating a solid H₂O + CO₂ mixture may suggest that H₂CO₃ might be found in outer space, where frozen ices of H₂O and CO₂ are common, as are cosmic rays and ultraviolet light, to help them react.^[10] The same carbonic acid polymorph (denoted beta-carbonic acid) was prepared by heating alternating layers of glassy aqueous solutions of bicarbonate and acid in vacuo, which causes protonation of bicarbonate, followed by removal of the solvent. The previously suggested alpha-carbonic acid, which was prepared by the same technique using methanol rather than water as a solvent was shown to be a monomethyl ester CH₃OCOOH.^[12]