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73-691: The Jumbles Reservoir is a heavily modified, high alkalinity , shallow reservoir in North West England . It lies in Jumbles Country Park , in the valley of Bradshaw Brook , partly in the Metropolitan Borough of Bolton , Greater Manchester , and partly in Blackburn with Darwen , Lancashire . It was opened on 11 March 1971 by Queen Elizabeth II for the then Bolton Corporation Waterworks (since privatisation

146-470: A covalent bond with an electron pair , known as a Lewis acid . The first category of acids are the proton donors, or Brønsted–Lowry acids . In the special case of aqueous solutions , proton donors form the hydronium ion H 3 O and are known as Arrhenius acids . Brønsted and Lowry generalized the Arrhenius theory to include non-aqueous solvents . A Brønsted or Arrhenius acid usually contains

219-437: A pH less than 7 and is colloquially also referred to as "acid" (as in "dissolved in acid"), while the strict definition refers only to the solute . A lower pH means a higher acidity , and thus a higher concentration of positive hydrogen ions in the solution. Chemicals or substances having the property of an acid are said to be acidic . Common aqueous acids include hydrochloric acid (a solution of hydrogen chloride that

292-414: A spans many orders of magnitude, a more manageable constant, p K a is more frequently used, where p K a = −log 10 K a . Stronger acids have a smaller p K a than weaker acids. Experimentally determined p K a at 25 °C in aqueous solution are often quoted in textbooks and reference material. Arrhenius acids are named according to their anions . In the classical naming system,

365-450: A values are small, but K a1 > K a2 . A triprotic acid (H 3 A) can undergo one, two, or three dissociations and has three dissociation constants, where K a1 > K a2 > K a3 . An inorganic example of a triprotic acid is orthophosphoric acid (H 3 PO 4 ), usually just called phosphoric acid . All three protons can be successively lost to yield H 2 PO 4 , then HPO 4 , and finally PO 4 ,

438-436: A Lewis acid explicitly as such. Modern definitions are concerned with the fundamental chemical reactions common to all acids. Most acids encountered in everyday life are aqueous solutions , or can be dissolved in water, so the Arrhenius and Brønsted–Lowry definitions are the most relevant. The Brønsted–Lowry definition is the most widely used definition; unless otherwise specified, acid–base reactions are assumed to involve

511-447: A Lewis acid, H , but at the same time, they also yield an equal amount of a Lewis base (acetate, citrate, or oxalate, respectively, for the acids mentioned). This article deals mostly with Brønsted acids rather than Lewis acids. Reactions of acids are often generalized in the form HA ⇌ H + A , where HA represents the acid and A is the conjugate base . This reaction is referred to as protolysis . The protonated form (HA) of an acid

584-638: A Lewis base and transfers a lone pair of electrons to form a bond with a hydrogen ion. The species that gains the electron pair is the Lewis acid; for example, the oxygen atom in H 3 O gains a pair of electrons when one of the H—O bonds is broken and the electrons shared in the bond become localized on oxygen. Depending on the context, a Lewis acid may also be described as an oxidizer or an electrophile . Organic Brønsted acids, such as acetic, citric, or oxalic acid, are not Lewis acids. They dissociate in water to produce

657-580: A benzene solvent and in the third gaseous HCl and NH 3 combine to form the solid. A third, only marginally related concept was proposed in 1923 by Gilbert N. Lewis , which includes reactions with acid–base characteristics that do not involve a proton transfer. A Lewis acid is a species that accepts a pair of electrons from another species; in other words, it is an electron pair acceptor. Brønsted acid–base reactions are proton transfer reactions while Lewis acid–base reactions are electron pair transfers. Many Lewis acids are not Brønsted–Lowry acids. Contrast how

730-472: A byproduct of oxidation reactions. The ocean's alkalinity varies over time, most significantly over geologic timescales (millennia). Changes in the balance between terrestrial weathering and sedimentation of carbonate minerals (for example, as a function of ocean acidification) are the primary long-term drivers of alkalinity in the ocean. Over human timescales, mean ocean alkalinity is relatively stable. Seasonal and annual variability of mean ocean alkalinity

803-433: A class of strong acids. A common example is toluenesulfonic acid (tosylic acid). Unlike sulfuric acid itself, sulfonic acids can be solids. In fact, polystyrene functionalized into polystyrene sulfonate is a solid strongly acidic plastic that is filterable. Superacids are acids stronger than 100% sulfuric acid. Examples of superacids are fluoroantimonic acid , magic acid and perchloric acid . The strongest known acid

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876-511: A covalent bond with an electron pair. An example is boron trifluoride (BF 3 ), whose boron atom has a vacant orbital that can form a covalent bond by sharing a lone pair of electrons on an atom in a base, for example the nitrogen atom in ammonia (NH 3 ). Lewis considered this as a generalization of the Brønsted definition, so that an acid is a chemical species that accepts electron pairs either directly or by releasing protons (H ) into

949-521: A hydrogen atom bonded to a chemical structure that is still energetically favorable after loss of H . Aqueous Arrhenius acids have characteristic properties that provide a practical description of an acid. Acids form aqueous solutions with a sour taste, can turn blue litmus red, and react with bases and certain metals (like calcium ) to form salts . The word acid is derived from the Latin acidus , meaning 'sour'. An aqueous solution of an acid has

1022-416: A much greater impact on oceanic alkalinity on short (minutes to centuries) timescales. Denitrification and sulfate reduction occur in oxygen-limited environments. Both of these processes consume hydrogen ions (thus increasing alkalinity) and release gases (N 2 or H 2 S), which eventually escape into the atmosphere. Nitrification and sulfide oxidation both decrease alkalinity by releasing protons as

1095-445: A neutral species, thus increasing alkalinity by one per equivalent. CO 3 however, will consume two protons before becoming a zero-level species (CO 2 ), thus it increases alkalinity by two per mole of CO 3 . [H ] and [ HSO 4 ] decrease alkalinity, as they act as sources of protons. They are often represented collectively as [H ] T . Alkalinity is typically reported as mg/L as CaCO 3 . (The conjunction "as"

1168-538: A pH of less than 7. While the Arrhenius concept is useful for describing many reactions, it is also quite limited in its scope. In 1923, chemists Johannes Nicolaus Brønsted and Thomas Martin Lowry independently recognized that acid–base reactions involve the transfer of a proton. A Brønsted–Lowry acid (or simply Brønsted acid) is a species that donates a proton to a Brønsted–Lowry base. Brønsted–Lowry acid–base theory has several advantages over Arrhenius theory. Consider

1241-590: A proton to ammonia (NH 3 ), but does not relate to the Arrhenius definition of an acid because the reaction does not produce hydronium. Nevertheless, CH 3 COOH is both an Arrhenius and a Brønsted–Lowry acid. Brønsted–Lowry theory can be used to describe reactions of molecular compounds in nonaqueous solution or the gas phase. Hydrogen chloride (HCl) and ammonia combine under several different conditions to form ammonium chloride , NH 4 Cl. In aqueous solution HCl behaves as hydrochloric acid and exists as hydronium and chloride ions. The following reactions illustrate

1314-464: A simple solution of an acid compound in water is determined by the dilution of the compound and the compound's K a . Lewis acids have been classified in the ECW model and it has been shown that there is no one order of acid strengths. The relative acceptor strength of Lewis acids toward a series of bases, versus other Lewis acids, can be illustrated by C-B plots . It has been shown that to define

1387-498: A solution with pH 7.0, which is only the case with similar acid and base strengths during a reaction. Neutralization with a base weaker than the acid results in a weakly acidic salt. An example is the weakly acidic ammonium chloride , which is produced from the strong acid hydrogen chloride and the weak base ammonia . Conversely, neutralizing a weak acid with a strong base gives a weakly basic salt (e.g., sodium fluoride from hydrogen fluoride and sodium hydroxide ). In order for

1460-444: A substance that increases the concentration of hydronium ions when added to water. Examples include molecular substances such as hydrogen chloride and acetic acid. An Arrhenius base , on the other hand, is a substance that increases the concentration of hydroxide (OH ) ions when dissolved in water. This decreases the concentration of hydronium because the ions react to form H 2 O molecules: Due to this equilibrium, any increase in

1533-455: A very large number of acidic protons. A diprotic acid (here symbolized by H 2 A) can undergo one or two dissociations depending on the pH. Each dissociation has its own dissociation constant, K a1 and K a2 . The first dissociation constant is typically greater than the second (i.e., K a1 > K a2 ). For example, sulfuric acid (H 2 SO 4 ) can donate one proton to form the bisulfate anion (HSO 4 ), for which K a1

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1606-436: A virtually limitless number of species that contribute to alkalinity. Alkalinity is frequently given as molar equivalents per liter of solution or per kilogram of solvent. In commercial (e.g. the swimming pool industry) and regulatory contexts, alkalinity might also be given in parts per million of equivalent calcium carbonate (ppm CaCO 3 ) . Alkalinity is sometimes incorrectly used interchangeably with basicity . For example,

1679-418: Is helium hydride ion , with a proton affinity of 177.8kJ/mol. Superacids can permanently protonate water to give ionic, crystalline hydronium "salts". They can also quantitatively stabilize carbocations . While K a measures the strength of an acid compound, the strength of an aqueous acid solution is measured by pH, which is an indication of the concentration of hydronium in the solution. The pH of

1752-539: Is 2 molar equivalents because twice as many H ions would be necessary to balance the charge. The total charge of a solution always equals zero. This leads to a parallel definition of alkalinity that is based upon the charge balance of ions in a solution. Certain ions, including Na , K , Ca , Mg , Cl , SO 4 , and NO 3 are " conservative " such that they are unaffected by changes in temperature, pressure or pH. Others such as HCO 3 are affected by changes in pH, temperature, and pressure. By isolating

1825-408: Is a substance that, when added to water, increases the concentration of H ions in the water. Chemists often write H ( aq ) and refer to the hydrogen ion when describing acid–base reactions but the free hydrogen nucleus, a proton , does not exist alone in water, it exists as the hydronium ion (H 3 O ) or other forms (H 5 O 2 , H 9 O 4 ). Thus, an Arrhenius acid can also be described as

1898-426: Is also sometimes referred to as the free acid . Acid–base conjugate pairs differ by one proton, and can be interconverted by the addition or removal of a proton ( protonation and deprotonation , respectively). The acid can be the charged species and the conjugate base can be neutral in which case the generalized reaction scheme could be written as HA ⇌ H + A . In solution there exists an equilibrium between

1971-544: Is appropriate in this case because the alkalinity results from a mixture of ions but is reported "as if" all of this is due to CaCO 3 .) This can be converted into milliequivalents per Liter (meq/L) by dividing by 50 (the approximate MW of CaCO 3 divided by 2). Addition (or removal) of CO 2 to a solution does not change its alkalinity, since the net reaction produces the same number of equivalents of positively contributing species (H ) as negative contributing species ( HCO 3 and/or CO 3 ). Adding CO 2 to

2044-400: Is completely dominated by carbonate and bicarbonate plus a small contribution from borate . Although alkalinity is primarily a term used by limnologists and oceanographers , it is also used by hydrologists to describe temporary hardness . Moreover, measuring alkalinity is important in determining a stream's ability to neutralize acidic pollution from rainfall or wastewater . It

2117-449: Is composed of CaCO 3 and its dissociation will add Ca and CO 3 into solution. Ca will not influence alkalinity, but CO 3 will increase alkalinity by 2 units. Increased dissolution of carbonate rock by acidification from acid rain and mining has contributed to increased alkalinity concentrations in some major rivers throughout the eastern U.S. The following reaction shows how acid rain, containing sulfuric acid, can have

2190-468: Is equal to [ HCO 3 ] + 2[ CO 3 ] is also approximately equal to the total alkalinity in surface water. Alkalinity measures the ability of a solution to neutralize acids to the equivalence point of carbonate or bicarbonate, defined as pH 4.5 for many oceanographic/limnological studies. The alkalinity is equal to the stoichiometric sum of the bases in solution. In most Earth surface waters carbonate alkalinity tends to make up most of

2263-605: Is found in gastric acid in the stomach and activates digestive enzymes ), acetic acid (vinegar is a dilute aqueous solution of this liquid), sulfuric acid (used in car batteries ), and citric acid (found in citrus fruits). As these examples show, acids (in the colloquial sense) can be solutions or pure substances, and can be derived from acids (in the strict sense) that are solids, liquids, or gases. Strong acids and some concentrated weak acids are corrosive , but there are exceptions such as carboranes and boric acid . The second category of acids are Lewis acids , which form

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2336-400: Is low. If this alkaline groundwater later comes into contact with the atmosphere, it can lose CO 2 , precipitate carbonate, and thereby become less alkaline again. When carbonate minerals, water, and the atmosphere are all in equilibrium, the reversible reaction shows that pH will be related to calcium ion concentration, with lower pH going with higher calcium ion concentration. In this case,

2409-518: Is often inversely proportional to sea surface temperature (SST). Therefore, it generally increases with high latitudes and depths. As a result, upwelling areas (where water from the deep ocean is pushed to the surface) also have higher alkalinity values. There are many programs to measure, record, and study oceanic alkalinity, together with many of the other characteristics of seawater, like temperature and salinity. These include: GEOSECS (Geochemical Ocean Sections Study), TTO/NAS (Transient Tracers in

2482-464: Is one of the best measures of the sensitivity of the stream to acid inputs. There can be long-term changes in the alkalinity of streams and rivers in response to human disturbances such as acid rain generated by SO x and NO x emissions. In 1884, Professor Wilhelm (William) Dittmar of Anderson College, now the University of Strathclyde , analysed 77 pristine seawater samples from around

2555-476: Is simply added to the name of the ionic compound. Thus, for hydrogen chloride, as an acid solution, the IUPAC name is aqueous hydrogen chloride. The strength of an acid refers to its ability or tendency to lose a proton. A strong acid is one that completely dissociates in water; in other words, one mole of a strong acid HA dissolves in water yielding one mole of H and one mole of the conjugate base, A , and none of

2628-422: Is very large; then it can donate a second proton to form the sulfate anion (SO 4 ), wherein the K a2 is intermediate strength. The large K a1 for the first dissociation makes sulfuric a strong acid. In a similar manner, the weak unstable carbonic acid (H 2 CO 3 ) can lose one proton to form bicarbonate anion (HCO 3 ) and lose a second to form carbonate anion (CO 3 ). Both K

2701-529: Is very low. Alkalinity varies by location depending on evaporation/precipitation, advection of water, biological processes, and geochemical processes. River dominated mixing also occurs close to the shore; it is strongest close to the mouth of a large river. Here, the rivers can act as either a source or a sink of alkalinity. A T follows the outflow of the river and has a linear relationship with salinity. Oceanic alkalinity also follows general trends based on latitude and depth. It has been shown that A T

2774-425: The citrate ion. Although the subsequent loss of each hydrogen ion is less favorable, all of the conjugate bases are present in solution. The fractional concentration, α (alpha), for each species can be calculated. For example, a generic diprotic acid will generate 3 species in solution: H 2 A, HA , and A . The fractional concentrations can be calculated as below when given either the pH (which can be converted to

2847-408: The ocean . Perhaps the most well known is the dissolution of calcium carbonate to form Ca and CO 3 (carbonate). The carbonate ion has the potential to absorb two hydrogen ions. Therefore, it causes a net increase in ocean alkalinity. Calcium carbonate dissolution occurs in regions of the ocean which are undersaturated with respect to calcium carbonate. The increasing carbon dioxide level in

2920-520: The 19th century, it is a variation of dumbles ; a northern term for a ravine like valley with wooded sides down which tumbles a fast flowing stream. This reservoir is also fed from the Wayoh and the Turton and Entwistle reservoirs. Alkalinity In freshwater , particularly those on non- limestone terrains, alkalinities are low and involve a lot of ions. In the ocean, on the other hand, alkalinity

2993-522: The CO 2 equivalence point where the major component in water is dissolved CO 2 which is converted to H 2 CO 3 in an aqueous solution. There are no strong acids or bases at this point. Therefore, the alkalinity is modeled and quantified with respect to the CO 2 equivalence point. Because the alkalinity is measured with respect to the CO 2 equivalence point, the dissolution of CO 2 , although it adds acid and dissolved inorganic carbon, does not change

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3066-573: The Ocean/North Atlantic Study), JGOFS (Joint Global Ocean Flux Study), WOCE (World Ocean Circulation Experiment), CARINA (Carbon dioxide in the Atlantic Ocean). The following packages calculate the state of the carbonate system in seawater (including pH): Acid An acid is a molecule or ion capable of either donating a proton (i.e. hydrogen ion, H ), known as a Brønsted–Lowry acid , or forming

3139-406: The [H ]) or the concentrations of the acid with all its conjugate bases: A plot of these fractional concentrations against pH, for given K 1 and K 2 , is known as a Bjerrum plot . A pattern is observed in the above equations and can be expanded to the general n -protic acid that has been deprotonated i -times: where K 0 = 1 and the other K-terms are the dissociation constants for

3212-424: The acid and its conjugate base. The equilibrium constant K is an expression of the equilibrium concentrations of the molecules or the ions in solution. Brackets indicate concentration, such that [H 2 O] means the concentration of H 2 O . The acid dissociation constant K a is generally used in the context of acid–base reactions. The numerical value of K a is equal to the product (multiplication) of

3285-435: The acid. Neutralization is the reaction between an acid and a base, producing a salt and neutralized base; for example, hydrochloric acid and sodium hydroxide form sodium chloride and water: Neutralization is the basis of titration , where a pH indicator shows equivalence point when the equivalent number of moles of a base have been added to an acid. It is often wrongly assumed that neutralization should result in

3358-439: The addition of CO 2 lowers the pH of a solution, thus reducing basicity while alkalinity remains unchanged ( see example below ). A variety of titrants , endpoints, and indicators are specified for various alkalinity measurement methods. Hydrochloric and sulfuric acids are common acid titrants, while phenolpthalein , methyl red , and bromocresol green are common indicators. In typical groundwater or seawater ,

3431-412: The alkalinity. In natural conditions, the dissolution of basic rocks and addition of ammonia [NH 3 ] or organic amines leads to the addition of base to natural waters at the CO 2 equivalence point. The dissolved base in water increases the pH and titrates an equivalent amount of CO 2 to bicarbonate ion and carbonate ion. At equilibrium, the water contains a certain amount of alkalinity contributed by

3504-424: The atmosphere , due to carbon dioxide emissions , results in increasing absorption of CO 2 from the atmosphere into the oceans. This does not affect the ocean's alkalinity but it does result in a reduction in pH value (called ocean acidification ). Ocean alkalinity enhancement has been proposed as one option to add alkalinity to the ocean and therefore buffer against pH changes. Biological processes have

3577-438: The concentration of hydronium is accompanied by a decrease in the concentration of hydroxide. Thus, an Arrhenius acid could also be said to be one that decreases hydroxide concentration, while an Arrhenius base increases it. In an acidic solution, the concentration of hydronium ions is greater than 10 moles per liter. Since pH is defined as the negative logarithm of the concentration of hydronium ions, acidic solutions thus have

3650-412: The concentration of weak acid anions. Conversely, the addition of acid converts weak acid anions to CO 2 and continuous addition of strong acids can cause the alkalinity to become less than zero. For example, the following reactions take place during the addition of acid to a typical seawater solution: It can be seen from the above protonation reactions that most bases consume one proton (H ) to become

3723-423: The concentrations of the products divided by the concentration of the reactants, where the reactant is the acid (HA) and the products are the conjugate base and H . The stronger of two acids will have a higher K a than the weaker acid; the ratio of hydrogen ions to acid will be higher for the stronger acid as the stronger acid has a greater tendency to lose its proton. Because the range of possible values for K

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3796-669: The conservative ions on one side of this charge balance equation, the nonconservative ions which accept or donate protons and thus define alkalinity are clustered on the other side of the equation. This combined charge balance and proton balance is called total alkalinity . Total alkalinity is not (much) affected by temperature, pressure, or pH, and is thus itself a conservative measurement, which increases its usefulness in aquatic systems. All anions except HCO 3 and CO 3 have low concentrations in Earth's surface water (streams, rivers, and lakes). Thus carbonate alkalinity , which

3869-409: The effect of increasing river alkalinity by increasing the amount of bicarbonate ion: Another way of writing this is: The lower the pH, the higher the concentration of bicarbonate will be. This shows how a lower pH can lead to higher alkalinity if the amount of bicarbonate produced is greater than the amount of H remaining after the reaction. This is the case since the amount of acid in the rainwater

3942-600: The existence of ions in solution, and defined acids as hydronium ion donors and bases as hydroxide ion donors. For that work, he received the Nobel Prize in Chemistry in 1903. See also Svante Arrhenius#Ionic disassociation . Alkalinity roughly refers to the molar amount of bases in a solution that can be converted to uncharged species by a strong acid. For example, 1 mole of HCO 3 in solution represents 1 molar equivalent, while 1 mole of CO 3

4015-413: The fluoride nucleus than they are in the lone fluoride ion. BF 3 is a Lewis acid because it accepts the electron pair from fluoride. This reaction cannot be described in terms of Brønsted theory because there is no proton transfer. The second reaction can be described using either theory. A proton is transferred from an unspecified Brønsted acid to ammonia, a Brønsted base; alternatively, ammonia acts as

4088-449: The following reactions are described in terms of acid–base chemistry: In the first reaction a fluoride ion , F , gives up an electron pair to boron trifluoride to form the product tetrafluoroborate . Fluoride "loses" a pair of valence electrons because the electrons shared in the B—F bond are located in the region of space between the two atomic nuclei and are therefore more distant from

4161-427: The following reactions of acetic acid (CH 3 COOH), the organic acid that gives vinegar its characteristic taste: Both theories easily describe the first reaction: CH 3 COOH acts as an Arrhenius acid because it acts as a source of H 3 O when dissolved in water, and it acts as a Brønsted acid by donating a proton to water. In the second example CH 3 COOH undergoes the same transformation, in this case donating

4234-428: The higher the pH, the more bicarbonate and carbonate ion there will be, in contrast to the paradoxical situation described above, where one does not have equilibrium with the atmosphere. In the ocean, alkalinity is completely dominated by carbonate and bicarbonate plus a small contribution from borate . Thus the chemical equation for alkalinity in seawater is: There are many methods of alkalinity generation in

4307-454: The ionic suffix is dropped and replaced with a new suffix, according to the table following. The prefix "hydro-" is used when the acid is made up of just hydrogen and one other element. For example, HCl has chloride as its anion, so the hydro- prefix is used, and the -ide suffix makes the name take the form hydrochloric acid . Classical naming system: In the IUPAC naming system, "aqueous"

4380-467: The limitations of Arrhenius's definition: As with the acetic acid reactions, both definitions work for the first example, where water is the solvent and hydronium ion is formed by the HCl solute. The next two reactions do not involve the formation of ions but are still proton-transfer reactions. In the second reaction hydrogen chloride and ammonia (dissolved in benzene ) react to form solid ammonium chloride in

4453-418: The measured total alkalinity is set equal to: (Subscript T indicates the total concentration of the species in the solution as measured. This is opposed to the free concentration, which takes into account the significant amount of ion pair interactions that occur in seawater.) Alkalinity can be measured by titrating a sample with a strong acid until all the buffering capacity of the aforementioned ions above

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4526-503: The more easily it loses a proton, H . Two key factors that contribute to the ease of deprotonation are the polarity of the H—A bond and the size of atom A, which determines the strength of the H—A bond. Acid strengths are also often discussed in terms of the stability of the conjugate base. Stronger acids have a larger acid dissociation constant , K a and a lower p K a than weaker acids. Sulfonic acids , which are organic oxyacids, are

4599-600: The order of Lewis acid strength at least two properties must be considered. For Pearson's qualitative HSAB theory the two properties are hardness and strength while for Drago's quantitative ECW model the two properties are electrostatic and covalent. Monoprotic acids, also known as monobasic acids, are those acids that are able to donate one proton per molecule during the process of dissociation (sometimes called ionization) as shown below (symbolized by HA): Common examples of monoprotic acids in mineral acids include hydrochloric acid (HCl) and nitric acid (HNO 3 ). On

4672-423: The orthophosphate ion, usually just called phosphate . Even though the positions of the three protons on the original phosphoric acid molecule are equivalent, the successive K a values differ since it is energetically less favorable to lose a proton if the conjugate base is more negatively charged. An organic example of a triprotic acid is citric acid , which can successively lose three protons to finally form

4745-740: The other hand, for organic acids the term mainly indicates the presence of one carboxylic acid group and sometimes these acids are known as monocarboxylic acid. Examples in organic acids include formic acid (HCOOH), acetic acid (CH 3 COOH) and benzoic acid (C 6 H 5 COOH). Polyprotic acids, also known as polybasic acids, are able to donate more than one proton per acid molecule, in contrast to monoprotic acids that only donate one proton per molecule. Specific types of polyprotic acids have more specific names, such as diprotic (or dibasic) acid (two potential protons to donate), and triprotic (or tribasic) acid (three potential protons to donate). Some macromolecules such as proteins and nucleic acids can have

4818-413: The pH of bicarbonate or carbonate is consumed. This point is functionally set to pH 4.5. At this point, all the bases of interest have been protonated to the zero level species, hence they no longer cause alkalinity. In the carbonate system the bicarbonate ions [ HCO 3 ] and the carbonate ions [ CO 3 ] have become converted to carbonic acid [H 2 CO 3 ] at this pH. This pH is also called

4891-423: The protonated acid HA. In contrast, a weak acid only partially dissociates and at equilibrium both the acid and the conjugate base are in solution. Examples of strong acids are hydrochloric acid (HCl), hydroiodic acid (HI), hydrobromic acid (HBr), perchloric acid (HClO 4 ), nitric acid (HNO 3 ) and sulfuric acid (H 2 SO 4 ). In water each of these essentially ionizes 100%. The stronger an acid is,

4964-480: The reservoir is now owned by United Utilities ). The reservoir's original purpose was to guarantee water for the Croal - Irwell river system and the associated industries. At the most northernly point of the reservoir is an old disused quarry, (Jumbles Quarry) it is now flooded with water due to the presence of the reservoir therefore making it impossible to tell that it is even there! The name Jumbles appeared during

5037-413: The solution lowers its pH, but does not affect alkalinity. At all pH values: Only at high (basic) pH values: Addition of CO 2 to a solution in contact with a solid can (over time) affect the alkalinity, especially for carbonate minerals in contact with groundwater or seawater. The dissolution (or precipitation) of carbonate rock has a strong influence on the alkalinity. This is because carbonate rock

5110-406: The solution, which then accept electron pairs. Hydrogen chloride, acetic acid, and most other Brønsted–Lowry acids cannot form a covalent bond with an electron pair, however, and are therefore not Lewis acids. Conversely, many Lewis acids are not Arrhenius or Brønsted–Lowry acids. In modern terminology, an acid is implicitly a Brønsted acid and not a Lewis acid, since chemists almost always refer to

5183-415: The total alkalinity due to the common occurrence and dissolution of carbonate rocks and other geological weathering processes that produce carbonate anions. Other common natural components that can contribute to alkalinity include borate , hydroxide , phosphate , silicate , dissolved ammonia , and the conjugate bases of organic acids (e.g., acetate ). Solutions produced in a laboratory may contain

5256-439: The transfer of a proton (H ) from an acid to a base. Hydronium ions are acids according to all three definitions. Although alcohols and amines can be Brønsted–Lowry acids, they can also function as Lewis bases due to the lone pairs of electrons on their oxygen and nitrogen atoms. In 1884, Svante Arrhenius attributed the properties of acidity to hydrogen ions (H ), later described as protons or hydrons . An Arrhenius acid

5329-590: The world brought back by the Challenger expedition . He found that in seawater the major ions were in a fixed ratio, confirming the hypothesis of Johan Georg Forchhammer , that is now known as the Principle of Constant Proportions. However, there was one exception. Dittmar found that the concentration of calcium was slightly greater in the deep ocean, and named this increase alkalinity. Also in 1884, Svante Arrhenius submitted his PhD theses in which he advocated

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