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57-407: KHP dissociates completely in water, giving the potassium cation (K) and hydrogen phthalate anion (HP or Hphthalate) and then, acting as a weak acid , hydrogen phthalate reacts reversibly with water to give hydronium (H 3 O) and phthalate ions. KHP can be used as a buffering agent in combination with hydrochloric acid (HCl) or sodium hydroxide (NaOH). The buffering region is dependent upon

114-404: A {\displaystyle K_{{\ce {a}}}} is known it can be used to determine the extent of dissociation in a solution with a given concentration of the acid, T H {\displaystyle T_{H}} , by applying the law of conservation of mass . where T H {\displaystyle T_{H}} is the value of the analytical concentration of

171-448: A {\displaystyle K_{{\ce {a}}}} value. The strength of a weak organic acid may depend on substituent effects. The strength of an inorganic acid is dependent on the oxidation state for the atom to which the proton may be attached. Acid strength is solvent-dependent. For example, hydrogen chloride is a strong acid in aqueous solution, but is a weak acid when dissolved in glacial acetic acid . The usual measure of

228-552: A {\displaystyle \mathrm {p} K_{{\ce {a}}}} = 3.2) or DMSO ( p K a {\displaystyle \mathrm {p} K_{{\ce {a}}}} = 15), has p K a {\displaystyle \mathrm {p} K_{{\ce {a}}}} values indicating that it undergoes incomplete dissociation in these solvents, making it a weak acid. However, as the rigorously dried, neat acidic medium, hydrogen fluoride has an H 0 {\displaystyle H_{0}} value of –15, making it

285-506: A {\displaystyle \mathrm {p} K_{{\ce {a}}}} , cannot be measured experimentally. The values in the following table are average values from as many as 8 different theoretical calculations. Also, in water The following can be used as protonators in organic chemistry Sulfonic acids , such as p-toluenesulfonic acid (tosylic acid) are a class of strong organic oxyacids . Some sulfonic acids can be isolated as solids. Polystyrene functionalized into polystyrene sulfonate

342-414: A {\displaystyle \mathrm {p} K_{{\ce {a}}}=-\log K_{\text{a}}} ) than weaker acids. The stronger an acid is, the more easily it loses a proton, H + {\displaystyle {\ce {H+}}} . Two key factors that contribute to the ease of deprotonation are the polarity of the H − A {\displaystyle {\ce {H-A}}} bond and

399-455: A differentiating solvent for the three acids, while water is not. An important example of a solvent which is more basic than water is dimethyl sulfoxide , DMSO, ( CH 3 ) 2 SO {\displaystyle {\ce {(CH3)2SO}}} . A compound which is a weak acid in water may become a strong acid in DMSO. Acetic acid is an example of such

456-434: A hydrogen-based economy . It is an alternative to electrolysis , and does not require hydrocarbons like current methods of steam reforming . But note that all of the available energy in the hydrogen so produced is supplied by the heat used to make it. Sulfuric acid is rarely encountered naturally on Earth in anhydrous form, due to its great affinity for water . Dilute sulfuric acid is a constituent of acid rain , which

513-610: A base and can be protonated, forming the [H 3 SO 4 ] ion. Salts of [H 3 SO 4 ] have been prepared (e.g. trihydroxyoxosulfonium hexafluoroantimonate(V) [H 3 SO 4 ] [SbF 6 ] ) using the following reaction in liquid HF : The above reaction is thermodynamically favored due to the high bond enthalpy of the Si–F bond in the side product. Protonation using simply fluoroantimonic acid , however, has met with failure, as pure sulfuric acid undergoes self-ionization to give [H 3 O] ions: which prevents

570-408: A burnt appearance in which the carbon appears much like soot that results from fire. Although less dramatic, the action of the acid on cotton , even in diluted form, destroys the fabric. The reaction with copper(II) sulfate can also demonstrate the dehydration property of sulfuric acid. The blue crystals change into white powder as water is removed. Sulfuric acid reacts with most bases to give

627-417: A considerable amount of heat is released; thus, the reverse procedure of adding water to the acid is generally avoided since the heat released may boil the solution, spraying droplets of hot acid during the process. Upon contact with body tissue, sulfuric acid can cause severe acidic chemical burns and secondary thermal burns due to dehydration. Dilute sulfuric acid is substantially less hazardous without

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684-411: A dehydrating agent, forming the nitronium ion NO + 2 , which is important in nitration reactions involving electrophilic aromatic substitution . This type of reaction, where protonation occurs on an oxygen atom, is important in many organic chemistry reactions, such as Fischer esterification and dehydration of alcohols. When allowed to react with superacids , sulfuric acid can act as

741-472: A more strongly protonating medium than 100% sulfuric acid and thus, by definition, a superacid . (To prevent ambiguity, in the rest of this article, "strong acid" will, unless otherwise stated, refer to an acid that is strong as measured by its p K a {\displaystyle \mathrm {p} K_{{\ce {a}}}} value ( p K a {\displaystyle \mathrm {p} K_{{\ce {a}}}} < –1.74). This usage

798-441: A simple solution of an acid in water is determined by both K a {\displaystyle K_{{\ce {a}}}} and the acid concentration. For weak acid solutions, it depends on the degree of dissociation , which may be determined by an equilibrium calculation. For concentrated solutions of acids, especially strong acids for which pH < 0, the H 0 {\displaystyle H_{0}} value

855-487: A slower rate, so that the acid neutralizing capacity (ANC) of the aquifer can neutralize the produced acid. In such cases, the total dissolved solids (TDS) concentration of the water can be increased from the dissolution of minerals from the acid-neutralization reaction with the minerals. Sulfuric acid is used as a defense by certain marine species, for example, the phaeophyte alga Desmarestia munda (order Desmarestiales ) concentrates sulfuric acid in cell vacuoles. In

912-462: A strong acid is perchloric acid . Any acid with a p K a {\displaystyle \mathrm {p} K_{{\ce {a}}}} value which is less than about -2 is classed as a strong acid. This results from the very high buffer capacity of solutions with a pH value of 1 or less and is known as the leveling effect . The following are strong acids in aqueous and dimethyl sulfoxide solution. The values of p K

969-402: A strong base. The conjugate of a weak acid is often a weak base and vice versa . The strength of an acid varies from solvent to solvent. An acid which is strong in water may be weak in a less basic solvent, and an acid which is weak in water may be strong in a more basic solvent. According to Brønsted–Lowry acid–base theory , the solvent S can accept a proton. For example, hydrochloric acid

1026-662: A substance. An extensive bibliography of p K a {\displaystyle \mathrm {p} K_{{\ce {a}}}} values in solution in DMSO and other solvents can be found at Acidity–Basicity Data in Nonaqueous Solvents . Superacids are strong acids even in solvents of low dielectric constant. Examples of superacids are fluoroantimonic acid and magic acid . Some superacids can be crystallised. They can also quantitatively stabilize carbocations . Lewis acids reacting with Lewis bases in gas phase and non-aqueous solvents have been classified in

1083-510: A vapor pressure of <1 mmHg at 40 °C. In the solid state, sulfuric acid is a molecular solid that forms monoclinic crystals with nearly trigonal lattice parameters. The structure consists of layers parallel to the (010) plane, in which each molecule is connected by hydrogen bonds to two others. Hydrates H 2 SO 4 · n H 2 O are known for n = 1, 2, 3, 4, 6.5, and 8, although most intermediate hydrates are stable against disproportionation . Anhydrous H 2 SO 4

1140-428: Is hygroscopic and readily absorbs water vapor from the air . Concentrated sulfuric acid is a strong oxidant with powerful dehydrating properties, making it highly corrosive towards other materials, from rocks to metals. Phosphorus pentoxide is a notable exception in that it is not dehydrated by sulfuric acid but, to the contrary, dehydrates sulfuric acid to sulfur trioxide . Upon addition of sulfuric acid to water,

1197-467: Is a better measure of acidity than the pH. A strong acid is an acid that dissociates according to the reaction where S represents a solvent molecule, such as a molecule of water or dimethyl sulfoxide (DMSO), to such an extent that the concentration of the undissociated species HA {\displaystyle {\ce {HA}}} is too low to be measured. For practical purposes a strong acid can be said to be completely dissociated. An example of

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1254-442: Is a common laboratory demonstration. The sugar darkens as carbon is formed, and a rigid column of black, porous carbon called a carbon snake may emerge. Similarly, mixing starch into concentrated sulfuric acid gives elemental carbon and water. The effect of this can also be seen when concentrated sulfuric acid is spilled on paper. Paper is composed of cellulose , a polysaccharide related to starch. The cellulose reacts to give

1311-484: Is a strong acid: The product of this ionization is HSO − 4 , the bisulfate anion. Bisulfate is a far weaker acid: The product of this second dissociation is SO 2− 4 , the sulfate anion. Concentrated sulfuric acid has a powerful dehydrating property, removing water ( H 2 O ) from other chemical compounds such as table sugar ( sucrose ) and other carbohydrates , to produce carbon , steam , and heat. Dehydration of table sugar (sucrose)

1368-429: Is a very polar liquid, having a dielectric constant of around 100. It has a high electrical conductivity , a consequence of autoprotolysis , i.e. self- protonation  : The equilibrium constant for autoprotolysis (25 °C) is: The corresponding equilibrium constant for water , K w is 10 , a factor of 10 (10 billion) smaller. In spite of the viscosity of the acid, the effective conductivities of

1425-404: Is a weak acid in solution in pure acetic acid , HO 2 CCH 3 {\displaystyle {\ce {HO2CCH3}}} , which is more acidic than water. The extent of ionization of the hydrohalic acids decreases in the order HI > HBr > HCl {\displaystyle {\ce {HI > HBr > HCl}}} . Acetic acid is said to be

1482-400: Is an example of a substance that is a solid strong acid. A weak acid is a substance that partially dissociates or partly ionizes when it is dissolved in a solvent. In solution, there is an equilibrium between the acid, HA {\displaystyle {\ce {HA}}} , and the products of dissociation. The solvent (e.g. water) is omitted from this expression when its concentration

1539-502: Is called acid mine drainage (AMD) or acid rock drainage (ARD). The Fe can be further oxidized to Fe : The Fe produced can be precipitated as the hydroxide or hydrous iron oxide : The iron(III) ion ("ferric iron") can also oxidize pyrite: When iron(III) oxidation of pyrite occurs, the process can become rapid. pH values below zero have been measured in ARD produced by this process. ARD can also produce sulfuric acid at

1596-403: Is consistent with the common parlance of most practicing chemists .) When the acidic medium in question is a dilute aqueous solution, the H 0 {\displaystyle H_{0}} is approximately equal to the pH value, which is a negative logarithm of the concentration of aqueous H + {\displaystyle {\ce {H+}}} in solution. The pH of

1653-408: Is effectively unchanged by the process of acid dissociation. The strength of a weak acid can be quantified in terms of a dissociation constant , K a {\displaystyle K_{a}} , defined as follows, where [ H ] {\displaystyle {\ce {[H]}}} signifies the concentration of a chemical moiety, X. When a numerical value of K

1710-402: Is formed by atmospheric oxidation of sulfur dioxide in the presence of water – i.e. oxidation of sulfurous acid . When sulfur-containing fuels such as coal or oil are burned, sulfur dioxide is the main byproduct (besides the chief products carbon oxides and water). Sulfuric acid is formed naturally by the oxidation of sulfide minerals, such as pyrite : The resulting highly acidic water

1767-432: Is from the carboxylate group, as illustrated by the following series of halogenated butanoic acids . In a set of oxoacids of an element, p K a {\displaystyle \mathrm {p} K_{{\ce {a}}}} values decrease with the oxidation state of the element. The oxoacids of chlorine illustrate this trend. † theoretical Sulfuric acid Sulfuric acid ( American spelling and

Potassium hydrogen phthalate - Misplaced Pages Continue

1824-434: Is fully protonated. The solution is then titrated with a strong base until only the deprotonated species, A − {\displaystyle {\ce {A-}}} , remains in solution. At each point in the titration pH is measured using a glass electrode and a pH meter . The equilibrium constant is found by fitting calculated pH values to the observed values, using the method of least squares . It

1881-419: Is only partially dissociated, or is partly ionized in water with both the undissociated acid and its dissociation products being present, in solution, in equilibrium with each other. Acetic acid ( CH 3 COOH {\displaystyle {\ce {CH3COOH}}} ) is an example of a weak acid. The strength of a weak acid is quantified by its acid dissociation constant , K

1938-493: Is prepared by slowly adding 98% sulfuric acid to an equal volume of water, with good stirring: the temperature of the mixture can rise to 80 °C (176 °F) or higher. Sulfuric acid contains not only H 2 SO 4 molecules, but is actually an equilibrium of many other chemical species, as it is shown in the table below. Sulfuric acid is a colorless oily liquid, and has a vapor pressure of <0.001 mmHg at 25 °C and 1 mmHg at 145.8 °C, and 98% sulfuric acid has

1995-429: Is said to be dibasic because it can lose two protons and react with two molecules of a simple base. Phosphoric acid ( H 3 PO 4 {\displaystyle {\ce {H3PO4}}} ) is tribasic. For a more rigorous treatment of acid strength see acid dissociation constant . This includes acids such as the dibasic acid succinic acid , for which the simple method of calculating

2052-416: Is sometimes stated that "the conjugate of a weak acid is a strong base". Such a statement is incorrect. For example, acetic acid is a weak acid which has a K a {\displaystyle K_{{\ce {a}}}} = 1.75 x 10 . Its conjugate base is the acetate ion with K b = 10 / K a = 5.7 x 10 (from the relationship K a × K b = 10 ), which certainly does not correspond to

2109-471: The H 3 SO + 4 and HSO − 4 ions are high due to an intramolecular proton-switch mechanism (analogous to the Grotthuss mechanism in water), making sulfuric acid a good conductor of electricity. It is also an excellent solvent for many reactions. The hydration reaction of sulfuric acid is highly exothermic , dilution. As indicated by its acid dissociation constant , sulfuric acid

2166-469: 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 the order of Lewis acid strength at least two properties must be considered. For the qualitative HSAB theory the two properties are hardness and strength while for

2223-936: The chemical formula HA {\displaystyle {\ce {HA}}} , to dissociate into a proton , H + {\displaystyle {\ce {H+}}} , and an anion , A − {\displaystyle {\ce {A-}}} . The dissociation or ionization of a strong acid in solution is effectively complete, except in its most concentrated solutions. Examples of strong acids are hydrochloric acid ( HCl ) {\displaystyle {\ce {(HCl)}}} , perchloric acid ( HClO 4 ) {\displaystyle {\ce {(HClO4)}}} , nitric acid ( HNO 3 ) {\displaystyle {\ce {(HNO3)}}} and sulfuric acid ( H 2 SO 4 ) {\displaystyle {\ce {(H2SO4)}}} . A weak acid

2280-520: The chemical industry . It is most commonly used in fertilizer manufacture but is also important in mineral processing , oil refining , wastewater processing , and chemical synthesis . It has a wide range of end applications, including in domestic acidic drain cleaners , as an electrolyte in lead-acid batteries , as a dehydrating compound, and in various cleaning agents . Sulfuric acid can be obtained by dissolving sulfur trioxide in water. Although nearly 100% sulfuric acid solutions can be made,

2337-418: The preferred IUPAC name ) or sulphuric acid ( Commonwealth spelling ), known in antiquity as oil of vitriol , is a mineral acid composed of the elements sulfur , oxygen , and hydrogen , with the molecular formula H 2 SO 4 . It is a colorless, odorless, and viscous liquid that is soluble with water. Pure sulfuric acid does not occur naturally due to its strong affinity to water vapor ; it

Potassium hydrogen phthalate - Misplaced Pages Continue

2394-448: The stratosphere , the atmosphere's second layer that is generally between 10 and 50 km above Earth's surface, sulfuric acid is formed by the oxidation of volcanic sulfur dioxide by the hydroxyl radical : Because sulfuric acid reaches supersaturation in the stratosphere, it can nucleate aerosol particles and provide a surface for aerosol growth via condensation and coagulation with other water-sulfuric acid aerosols. This results in

2451-418: The stratospheric aerosol layer . The permanent Venusian clouds produce a concentrated acid rain, as the clouds in the atmosphere of Earth produce water rain. Jupiter 's moon Europa is also thought to have an atmosphere containing sulfuric acid hydrates. Sulfuric acid is produced from sulfur , oxygen and water via the conventional contact process (DCDA) or the wet sulfuric acid process (WSA). In

2508-502: The acid produced in the lead chamber itself (<70% to avoid contamination with nitrosylsulfuric acid ) and tower acid being the acid recovered from the bottom of the Glover tower. They are now obsolete as commercial concentrations of sulfuric acid, although they may be prepared in the laboratory from concentrated sulfuric acid if needed. In particular, "10 M" sulfuric acid (the modern equivalent of chamber acid, used in many titrations ),

2565-661: The acid. When all the quantities in this equation are treated as numbers, ionic charges are not shown and this becomes a quadratic equation in the value of the hydrogen ion concentration value, [ H ] {\displaystyle {\ce {[H]}}} . This equation shows that the pH of a solution of a weak acid depends on both its K a {\displaystyle K_{{\ce {a}}}} value and its concentration. Typical examples of weak acids include acetic acid and phosphorous acid . An acid such as oxalic acid ( HOOC − COOH {\displaystyle {\ce {HOOC-COOH}}} )

2622-842: The conversion of H 2 SO 4 to [H 3 SO 4 ] by the HF/ SbF 5 system. Even dilute sulfuric acid reacts with many metals via a single displacement reaction, like other typical acids , producing hydrogen gas and salts (the metal sulfate). It attacks reactive metals (metals at positions above copper in the reactivity series ) such as iron , aluminium , zinc , manganese , magnesium , and nickel . Concentrated sulfuric acid can serve as an oxidizing agent , releasing sulfur dioxide: Lead and tungsten , however, are resistant to sulfuric acid. Hot concentrated sulfuric acid oxidizes carbon (as bituminous coal ) and sulfur : Benzene and many derivatives undergo electrophilic aromatic substitution with sulfuric acid to give

2679-412: The corresponding sulfonic acids : Sulfuric acid can be used to produce hydrogen from water : The compounds of sulfur and iodine are recovered and reused, hence the process is called the sulfur–iodine cycle . This process is endothermic and must occur at high temperatures, so energy in the form of heat has to be supplied. The sulfur–iodine cycle has been proposed as a way to supply hydrogen for

2736-606: The corresponding sulfate or bisulfate. Sulfuric acid reacts with sodium chloride , and gives hydrogen chloride gas and sodium bisulfate : Aluminium sulfate , also known as paper maker's alum, is made by treating bauxite with sulfuric acid: Sulfuric acid can also be used to displace weaker acids from their salts. Reaction with sodium acetate , for example, displaces acetic acid , CH 3 COOH , and forms sodium bisulfate : Similarly, treating potassium nitrate with sulfuric acid produces nitric acid . When combined with nitric acid , sulfuric acid acts both as an acid and

2793-399: The oxidation of organics to carbon dioxide and water, with subsequent quantitation of the carbon dioxide. Many TOC analysts suggest testing their instruments with two standards: one typically easy for the instrument to oxidize (KHP), and one more difficult to oxidize. For the latter, benzoquinone is suggested. Weak acid Acid strength is the tendency of an acid , symbolised by

2850-415: The oxidative and dehydrating properties; though, it is handled with care for its acidity. Sulfuric acid is a very important commodity chemical; a country's sulfuric acid production is a good indicator of its industrial strength. Many methods for its production are known, including the contact process , the wet sulfuric acid process , and the lead chamber process . Sulfuric acid is also a key substance in

2907-405: The pH of a solution, shown above, cannot be used. The experimental determination of a p K a {\displaystyle \mathrm {p} K_{{\ce {a}}}} value is commonly performed by means of a titration . A typical procedure would be as follows. A quantity of strong acid is added to a solution containing the acid or a salt of the acid, to the point where the compound

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2964-469: The pKa, and is typically +/- 1.0 pH units of the pKa. The pKa of KHP is 5.4, so its pH buffering range would be 4.4 to 6.4; however, due to the presence of the second acidic group that bears the potassium ion, the first pKa also contributes to the buffering range well below pH 4.0, which is why KHP is a good choice for use as a reference standard for pH 4.00. KHP is also a useful standard for total organic carbon (TOC) testing. Most TOC analyzers are based on

3021-410: The quantitative ECW model the two properties are electrostatic and covalent. In organic carboxylic acids, an electronegative substituent can pull electron density out of an acidic bond through the inductive effect , resulting in a smaller p K a {\displaystyle \mathrm {p} K_{{\ce {a}}}} value. The effect decreases, the further the electronegative element

3078-442: The size of atom A, which determine the strength of the H − A {\displaystyle {\ce {H-A}}} bond. Acid strengths also depend on the stability of the conjugate base. While the p K a {\displaystyle \mathrm {p} K_{{\ce {a}}}} value measures the tendency of an acidic solute to transfer a proton to a standard solvent (most commonly water or DMSO ),

3135-426: The strength of an acid is its acid dissociation constant ( K a {\displaystyle K_{{\ce {a}}}} ), which can be determined experimentally by titration methods. Stronger acids have a larger K a {\displaystyle K_{{\ce {a}}}} and a smaller logarithmic constant ( p K a = − log ⁡ K

3192-445: The subsequent loss of SO 3 at the boiling point brings the concentration to 98.3% acid. The 98.3% grade, which is more stable in storage, is the usual form of what is described as "concentrated sulfuric acid". Other concentrations are used for different purposes. Some common concentrations are: "Chamber acid" and "tower acid" were the two concentrations of sulfuric acid produced by the lead chamber process , chamber acid being

3249-683: The tendency of an acidic solvent to transfer a proton to a reference solute (most commonly a weak aniline base) is measured by its Hammett acidity function , the H 0 {\displaystyle H_{0}} value. Although these two concepts of acid strength often amount to the same general tendency of a substance to donate a proton, the p K a {\displaystyle \mathrm {p} K_{{\ce {a}}}} and H 0 {\displaystyle H_{0}} values are measures of distinct properties and may occasionally diverge. For instance, hydrogen fluoride, whether dissolved in water ( p K

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