Hubbry Logo
search
logo
Acid avatar
Acid
Acid
current hub
2327889

Acid

logo
Community Hub0 Subscribers
Read side by side
Acid
View on Wikipedia
from Wikipedia

Zinc, a typical metal, reacting with hydrochloric acid, a typical acid
Acids and bases
Diagrammatic representation of the dissociation of acetic acid in aqueous solution to acetate and hydronium ions.
Acid types
Base types

An acid is a molecule or ion capable of either donating a proton (i.e. hydrogen cation, H+), known as a Brønsted–Lowry acid, or forming a covalent bond with an electron pair, known as a Lewis acid.[1]

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 H3O+ and are known as Arrhenius acids. Brønsted and Lowry generalized the Arrhenius theory to include non-aqueous solvents. A Brønsted–Lowry or Arrhenius acid usually contains 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.[2] 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'.[3] An aqueous solution of an acid has 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.[1] A lower pH means a higher acidity, and thus a higher concentration of hydrogen cations 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 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[1] 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 a covalent bond with an electron pair. An example is boron trifluoride (BF3), 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 (NH3). 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 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.[4] 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 a Lewis acid explicitly as such.[4]

Definitions and concepts

[edit]
Main article: Acid–base reaction

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 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.

Arrhenius acids

[edit]
Svante Arrhenius

In 1884, Svante Arrhenius attributed the properties of acidity to hydrogen cations (H+), later described as protons or hydrons. An Arrhenius acid is a substance that, when added to water, increases the concentration of H+ ions in the water.[4][5] Chemists often write H+(aq) and refer to the hydrogen cation when describing acid–base reactions but the free hydrogen nucleus, a proton, does not exist alone in water, it exists as the hydronium ion (H3O+) or other forms (H5O2+, H9O4+). Thus, an Arrhenius acid can also be described as 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 H2O molecules:

H3O+
(aq) + OH−
(aq) ⇌ H2O(liq) + H2O(liq)

Due to this equilibrium, any increase in 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 10−7 moles per liter. Since pH is defined as the negative logarithm of the concentration of hydronium ions, acidic solutions thus have a pH of less than 7.

Brønsted–Lowry acids

[edit]
Main article: Brønsted–Lowry acid–base theory
Acetic acid, CH3COOH, is composed of a methyl group, CH3, bound chemically to a carboxylate group, COOH. The carboxylate group can lose a proton and donate it to a water molecule, H20, leaving behind an acetate anion CH3COO- and creating a hydronium cation H3O. This is an equilibrium reaction, so the reverse process can also take place.
Acetic acid, a weak acid, donates a proton (hydrogen ion, highlighted in green) to water in an equilibrium reaction to give the acetate ion and the hydronium ion. Red: oxygen, black: carbon, white: hydrogen.

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.[5] Brønsted–Lowry acid–base theory has several advantages over Arrhenius theory. Consider the following reactions of acetic acid (CH3COOH), the organic acid that gives vinegar its characteristic taste:

CH3COOH + H2O ⇌ CH3COO− + H3O+
CH3COOH + NH3 ⇌ CH3COO− + NH+4

Both theories easily describe the first reaction: CH3COOH acts as an Arrhenius acid because it acts as a source of H3O+ when dissolved in water, and it acts as a Brønsted acid by donating a proton to water. In the second example CH3COOH undergoes the same transformation, in this case donating a proton to ammonia (NH3), but does not relate to the Arrhenius definition of an acid because the reaction does not produce hydronium. Nevertheless, CH3COOH 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, NH4Cl. In aqueous solution HCl behaves as hydrochloric acid and exists as hydronium and chloride ions. The following reactions illustrate the limitations of Arrhenius's definition:

  1. H3O+
    (aq)     + Cl−
    (aq)     + NH3 → Cl−
    (aq)     + NH+
    4    (aq) + H2O
  2. HCl(benzene) + NH3(benzene) → NH4Cl(s)
  3. HCl(g) + NH3(g) → NH4Cl(s)

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 a benzene solvent and in the third gaseous HCl and NH3 combine to form the solid.

Lewis acids

[edit]
Main article: Lewis acids and bases

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.[5] 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 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 the fluoride nucleus than they are in the lone fluoride ion. BF3 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 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 H3O+ 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.[4] They dissociate in water to produce 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.

Dissociation and equilibrium

[edit]

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 is also sometimes referred to as the free acid.[6]

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 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 [H2O] means the concentration of H2O. The acid dissociation constant Ka is generally used in the context of acid–base reactions. The numerical value of Ka is equal to the product (multiplication) of 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 Ka than the weaker acid; the ratio of hydrogen cations 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 Ka spans many orders of magnitude, a more manageable constant, pKa is more frequently used, where pKa = −log10 Ka. Stronger acids have a smaller pKa than weaker acids. Experimentally determined pKa at 25 °C in aqueous solution are often quoted in textbooks and reference material.

Nomenclature

[edit]

Arrhenius acids are named according to their anions. In the classical naming system, 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:

Anion prefix Anion suffix Acid prefix Acid suffix Example
per ate per ic acid perchloric acid (HClO4)
chloric acid (HClO3)
ite ous acid chlorous acid (HClO2)
hypo ite hypo ous acid hypochlorous acid (HClO)
ide hydro ic acid hydrochloric acid (HCl)

In the IUPAC naming system, "aqueous" 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.

Acid strength

[edit]
Main article: Acid strength

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 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 (HClO4), nitric acid (HNO3) and sulfuric acid (H2SO4). In water, each of these essentially ionizes 100%. The stronger an acid is, 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, Ka and a lower pKa than weaker acids.

Sulfonic acids, which are organic oxyacids, are 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 is helium hydride ion,[7] with a proton affinity of 177.8kJ/mol.[8] Superacids can permanently protonate water to give ionic, crystalline hydronium "salts". They can also quantitatively stabilize carbocations.

While Ka 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 a simple solution of an acid compound in water is determined by the dilution of the compound and the compound's Ka.

Lewis acid strength in non-aqueous solutions

[edit]

Lewis acids have been classified in the ECW model and it has been shown that there is no one order of acid strengths.[9] The relative acceptor strength of Lewis acids toward a series of bases, versus other Lewis acids, can be illustrated by C-B plots.[10][11] It has been shown that to define 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.

Chemical characteristics

[edit]

Monoprotic acids

[edit]
See also: Acid dissociation constant § Monoprotic acids

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):

HA (aq) + H2O (l) ⇌ H3O+ (aq) + A− (aq)      Ka

Common examples of monoprotic acids in mineral acids include hydrochloric acid (HCl) and nitric acid (HNO3). On 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 (CH3COOH) and benzoic acid (C6H5COOH).

Polyprotic acids

[edit]
See also: Acid dissociation constant § Polyprotic acids

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 a very large number of acidic protons.[12]

A diprotic acid (here symbolized by H2A) can undergo one or two dissociations depending on the pH. Each dissociation has its own dissociation constant, Ka1 and Ka2.

H2A (aq) + H2O (l) ⇌ H3O+ (aq) + HA− (aq)     Ka1
HA− (aq) + H2O (l) ⇌ H3O+ (aq) + A2− (aq)       Ka2

The first dissociation constant is typically greater than the second (i.e., Ka1 > Ka2). For example, sulfuric acid (H2SO4) can donate one proton to form the bisulfate anion (HSO−
4), for which Ka1 is very large; then it can donate a second proton to form the sulfate anion (SO2−
4), wherein the Ka2 is intermediate strength. The large Ka1 for the first dissociation makes sulfuric a strong acid. In a similar manner, the weak unstable carbonic acid (H2CO3) can lose one proton to form bicarbonate anion (HCO−
3) and lose a second to form carbonate anion (CO2−
3). Both Ka values are small, but Ka1 > Ka2 .

A triprotic acid (H3A) can undergo one, two, or three dissociations and has three dissociation constants, where Ka1 > Ka2 > Ka3.

H3A (aq) + H2O (l) ⇌ H3O+ (aq) + H2A− (aq)      Ka1
H2A− (aq) + H2O (l) ⇌ H3O+ (aq) + HA2− (aq)       Ka2
HA2− (aq) + H2O (l) ⇌ H3O+ (aq) + A3− (aq)      Ka3

An inorganic example of a triprotic acid is orthophosphoric acid (H3PO4), usually just called phosphoric acid. All three protons can be successively lost to yield H2PO−
4, then HPO2−
4, and finally PO3−
4, 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 Ka 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 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: H2A, HA−, and A2−. The fractional concentrations can be calculated as below when given either the pH (which can be converted to the [H+]) or the concentrations of the acid with all its conjugate bases:

A plot of these fractional concentrations against pH, for given K1 and K2, 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 K0 = 1 and the other K-terms are the dissociation constants for the acid.

Neutralization

[edit]
Hydrochloric acid (in beaker) reacting with ammonia fumes to produce ammonium chloride (white smoke)

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:

HCl(aq) + NaOH(aq) → H2O(l) + NaCl(aq)

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 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).

Weak acid–weak base equilibrium

[edit]
Main article: Henderson–Hasselbalch equation

In order for a protonated acid to lose a proton, the pH of the system must rise above the pKa of the acid. The decreased concentration of H+ in that basic solution shifts the equilibrium towards the conjugate base form (the deprotonated form of the acid). In lower-pH (more acidic) solutions, there is a high enough H+ concentration in the solution to cause the acid to remain in its protonated form.

Solutions of weak acids and salts of their conjugate bases form buffer solutions.

Titration

[edit]
Main article: Acid–base titration

To determine the concentration of an acid in an aqueous solution, an acid–base titration is commonly performed. A strong base solution with a known concentration, usually NaOH or KOH, is added to neutralize the acid solution according to the color change of the indicator with the amount of base added.[13] The titration curve of an acid titrated by a base has two axes, with the base volume on the x-axis and the solution's pH value on the y-axis. The pH of the solution always goes up as the base is added to the solution.

Example: Diprotic acid

[edit]
This is an ideal titration curve for alanine, a diprotic amino acid.[14] Point 2 is the first equivalent point where the amount of NaOH added equals the amount of alanine in the original solution.

For each diprotic acid titration curve, from left to right, there are two midpoints, two equivalence points, and two buffer regions.[15]

Equivalence points

[edit]

Due to the successive dissociation processes, there are two equivalence points in the titration curve of a diprotic acid.[16] The first equivalence point occurs when all first protons from the first ionization are titrated.[17] In other words, the amount of OH− added equals the original amount of H2A at the first equivalence point. The second equivalence point occurs when all protons are titrated. Therefore, the amount of OH− added equals twice the amount of H2A at this time. For a weak diprotic acid titrated by a strong base, the second equivalence point must occur at pH above 7 due to the hydrolysis of the resulted salts in the solution.[17] At either equivalence point, adding a drop of base will cause the steepest rise of the pH value in the system.

Buffer regions and midpoints

[edit]

A titration curve for a diprotic acid contains two midpoints where pH=pKa. Since there are two different Ka values, the first midpoint occurs at pH=pKa1 and the second one occurs at pH=pKa2.[18] Each segment of the curve that contains a midpoint at its center is called the buffer region. Because the buffer regions consist of the acid and its conjugate base, it can resist pH changes when base is added until the next equivalent points.[5]

Applications of acids

[edit]

In industry

[edit]

Acids are fundamental reagents in treating almost all processes in modern industry. Sulfuric acid, a diprotic acid, is the most widely used acid in industry, and is also the most-produced industrial chemical in the world. It is mainly used in producing fertilizer, detergent, batteries and dyes, as well as used in processing many products such like removing impurities.[19] According to the statistics data in 2011, the annual production of sulfuric acid was around 200 million tonnes in the world.[20] For example, phosphate minerals react with sulfuric acid to produce phosphoric acid for the production of phosphate fertilizers, and zinc is produced by dissolving zinc oxide into sulfuric acid, purifying the solution and electrowinning.

In the chemical industry, acids react in neutralization reactions to produce salts. For example, nitric acid reacts with ammonia to produce ammonium nitrate, a fertilizer. Additionally, carboxylic acids can be esterified with alcohols, to produce esters.

Acids are often used to remove rust and other corrosion from metals in a process known as pickling. They may be used as an electrolyte in a wet cell battery, such as sulfuric acid in a car battery.

In food

[edit]
Carbonated water (H2CO3 aqueous solution) is commonly added to soft drinks to make them effervesce.

Tartaric acid is an important component of some commonly used foods like unripened mangoes and tamarind. Natural fruits and vegetables also contain acids. Citric acid is present in oranges, lemon and other citrus fruits. Oxalic acid is present in tomatoes, spinach, and especially in carambola and rhubarb; rhubarb leaves and unripe carambolas are toxic because of high concentrations of oxalic acid. Ascorbic acid (Vitamin C) is an essential vitamin for the human body and is present in such foods as amla (Indian gooseberry), lemon, citrus fruits, and guava.

Many acids can be found in various kinds of food as additives, as they alter their taste and serve as preservatives. Phosphoric acid, for example, is a component of cola drinks. Acetic acid is used in day-to-day life as vinegar. Citric acid is used as a preservative in sauces and pickles.

Carbonic acid is one of the most common acid additives that are widely added in soft drinks. During the manufacturing process, CO2 is usually pressurized to dissolve in these drinks to generate carbonic acid. Carbonic acid is very unstable and tends to decompose into water and CO2 at room temperature and pressure. Therefore, when bottles or cans of these kinds of soft drinks are opened, the soft drinks fizz and effervesce as CO2 bubbles come out.[21]

Certain acids are used as drugs. Acetylsalicylic acid (Aspirin) is used as a pain killer and for bringing down fevers.

In human bodies

[edit]

Acids play important roles in the human body. The hydrochloric acid present in the stomach aids digestion by breaking down large and complex food molecules. Amino acids are required for synthesis of proteins required for growth and repair of body tissues. Fatty acids are also required for growth and repair of body tissues. Nucleic acids are important for the manufacturing of DNA and RNA and transmitting of traits to offspring through genes. Carbonic acid is important for maintenance of pH equilibrium in the body.

Human bodies contain a variety of organic and inorganic compounds, among those dicarboxylic acids play an essential role in many biological behaviors. Many of those acids are amino acids, which mainly serve as materials for the synthesis of proteins.[22] Other weak acids serve as buffers with their conjugate bases to keep the body's pH from undergoing large scale changes that would be harmful to cells.[23] The rest of the dicarboxylic acids also participate in the synthesis of various biologically important compounds in human bodies.

Acid catalysis

[edit]
Main article: Acid catalysis

Acids are used as catalysts in industrial and organic chemistry; for example, sulfuric acid is used in very large quantities in the alkylation process to produce gasoline. Some acids, such as sulfuric, phosphoric, and hydrochloric acids, also effect dehydration and condensation reactions. In biochemistry, many enzymes employ acid catalysis.[24]

Biological occurrence

[edit]
Basic structure of an amino acid

Many biologically important molecules are acids. Nucleic acids, which contain acidic phosphate groups, include DNA and RNA. Nucleic acids contain the genetic code that determines many of an organism's characteristics, and is passed from parents to offspring. DNA contains the chemical blueprint for the synthesis of proteins, which are made up of amino acid subunits. Cell membranes contain fatty acid esters such as phospholipids.

An α-amino acid has a central carbon (the α or alpha carbon) that is covalently bonded to a carboxyl group (thus they are carboxylic acids), an amino group, a hydrogen atom and a variable group. The variable group, also called the R group or side chain, determines the identity and many of the properties of a specific amino acid. In glycine, the simplest amino acid, the R group is a hydrogen atom, but in all other amino acids it is contains one or more carbon atoms bonded to hydrogens, and may contain other elements such as sulfur, oxygen or nitrogen. With the exception of glycine, naturally occurring amino acids are chiral and almost invariably occur in the L-configuration. Peptidoglycan, found in some bacterial cell walls contains some D-amino acids. At physiological pH, typically around 7, free amino acids exist in a charged form, where the acidic carboxyl group (-COOH) loses a proton (-COO−) and the basic amine group (-NH2) gains a proton (-NH+
3). The entire molecule has a net neutral charge and is a zwitterion, with the exception of amino acids with basic or acidic side chains. Aspartic acid, for example, possesses one protonated amine and two deprotonated carboxyl groups, for a net charge of −1 at physiological pH.

Fatty acids and fatty acid derivatives are another group of carboxylic acids that play a significant role in biology. These contain long hydrocarbon chains and a carboxylic acid group on one end. The cell membrane of nearly all organisms is primarily made up of a phospholipid bilayer, a micelle of hydrophobic fatty acid esters with polar, hydrophilic phosphate "head" groups. Membranes contain additional components, some of which can participate in acid–base reactions.

In humans and many other animals, hydrochloric acid is a part of the gastric acid secreted within the stomach to help hydrolyze proteins and polysaccharides, as well as converting the inactive pro-enzyme, pepsinogen into the enzyme, pepsin. Some organisms produce acids for defense; for example, ants produce formic acid.

Acid–base equilibrium plays a critical role in regulating mammalian breathing. Oxygen gas (O2) drives cellular respiration, the process by which animals release the chemical potential energy stored in food, producing carbon dioxide (CO2) as a byproduct. Oxygen and carbon dioxide are exchanged in the lungs, and the body responds to changing energy demands by adjusting the rate of ventilation. For example, during periods of exertion the body rapidly breaks down stored carbohydrates and fat, releasing CO2 into the blood stream. In aqueous solutions such as blood CO2 exists in equilibrium with carbonic acid and bicarbonate ion.

CO2 + H2O ⇌ H2CO3 ⇌ H+ + HCO−3

It is the decrease in pH that signals the brain to breathe faster and deeper, expelling the excess CO2 and resupplying the cells with O2.

Aspirin (acetylsalicylic acid) is a carboxylic acid.

Cell membranes are generally impermeable to charged or large, polar molecules because of the lipophilic fatty acyl chains comprising their interior. Many biologically important molecules, including a number of pharmaceutical agents, are organic weak acids that can cross the membrane in their protonated, uncharged form but not in their charged form (i.e., as the conjugate base). For this reason the activity of many drugs can be enhanced or inhibited by the use of antacids or acidic foods. The charged form, however, is often more soluble in blood and cytosol, both aqueous environments. When the extracellular environment is more acidic than the neutral pH within the cell, certain acids will exist in their neutral form and will be membrane soluble, allowing them to cross the phospholipid bilayer. Acids that lose a proton at the intracellular pH will exist in their soluble, charged form and are thus able to diffuse through the cytosol to their target. Ibuprofen, aspirin and penicillin are examples of drugs that are weak acids.

Common acids

[edit]

Mineral acids (inorganic acids)

[edit]

Sulfonic acids

[edit]

A sulfonic acid has the general formula RS(=O)2–OH, where R is an organic radical.

Carboxylic acids

[edit]

A carboxylic acid has the general formula R-C(O)OH, where R is an organic radical. The carboxyl group -C(O)OH contains a carbonyl group, C=O, and a hydroxyl group, O-H.

Halogenated carboxylic acids

[edit]

Halogenation at alpha position increases acid strength, so that the following acids are all stronger than acetic acid.

Vinylogous carboxylic acids

[edit]

Normal carboxylic acids are the direct union of a carbonyl group and a hydroxyl group. In vinylogous carboxylic acids, a carbon-carbon double bond separates the carbonyl and hydroxyl groups.

Nucleic acids

[edit]

References

[edit]
[edit]
Revisions and contributorsEdit on WikipediaRead on Wikipedia

Acid

View on Grokipedia
from Grokipedia
In chemistry, an acid is defined as a molecular entity or chemical species capable of donating a hydron (proton, H⁺) or forming a covalent bond with an electron pair.[1] This encompasses the Brønsted-Lowry concept of proton donation and the Lewis concept of electron-pair acceptance, providing a unified framework for understanding acid behavior across various solvents and reaction conditions.[1] Acids are fundamental to numerous chemical processes, exhibiting properties such as a pH less than 7 in aqueous solutions, the ability to turn litmus paper red, and a sour taste in dilute forms. The concept of acids evolved through key theoretical advancements in the late 19th and early 20th centuries. In 1884, Svante Arrhenius proposed the first modern definition, describing acids as substances that increase the concentration of hydrogen ions (H⁺) when dissolved in water, laying the groundwork for understanding ionization in aqueous solutions. This Arrhenius model was expanded in 1923 by Johannes Brønsted and Thomas Lowry, who redefined acids as proton (H⁺) donors in any acid-base reaction, independent of the solvent and applicable to a broader range of chemical systems. Concurrently, Gilbert N. Lewis introduced a more general perspective in 1923, classifying acids as electron-pair acceptors, which extended the theory to non-protonic reactions and coordination chemistry. These definitions—Arrhenius, Brønsted-Lowry, and Lewis—remain central to contemporary acid-base chemistry, with the IUPAC Gold Book integrating them into its current nomenclature.[1] Acids are classified by their strength, source, and structure, influencing their reactivity and applications. Strong acids, such as hydrochloric acid (HCl) and sulfuric acid (H₂SO₄), fully dissociate in water to yield H⁺ ions, resulting in high conductivity and corrosive properties.[2] In contrast, weak acids like acetic acid (CH₃COOH) partially dissociate, establishing equilibrium with their conjugate bases and exhibiting milder effects.[2] Structurally, acids include binary acids (also known as hydracids; e.g., HF, HCl, HBr, HI, H₂S), which consist of hydrogen bonded to a nonmetal and contain no oxygen; oxoacids (also known as oxyacids; e.g., H₂SO₄, HNO₃, HClO₄, H₃PO₄, H₂CO₃), which contain hydrogen, oxygen, and a central nonmetal element (or certain transition metals); and organic acids (e.g., citric acid, formic acid), typically featuring a carboxyl group (-COOH) and prevalent in biological systems.[3] Common examples also encompass carbonic acid (H₂CO₃) from dissolved CO₂ and phosphoric acid (H₃PO₄) used in food additives. Beyond fundamental reactions like neutralization with bases to form salts and water, acids play pivotal roles in industry and daily life.[4] Sulfuric acid, the most industrially produced chemical worldwide, is essential for manufacturing fertilizers (e.g., phosphate-based), petroleum refining, metal extraction, and battery production, with global output of approximately 261 million metric tons annually as of 2024.[5] Hydrochloric acid is vital for steel pickling to remove rust and in pH adjustment for water treatment, while nitric acid supports explosives and fertilizer synthesis.[4] In biology and food science, organic acids like citric and lactic acid act as preservatives, flavor enhancers, and metabolic intermediates, underscoring acids' ubiquity in sustaining chemical equilibrium and enabling diverse technological advancements.[6]

Definitions

Arrhenius Acids

The Arrhenius theory of acids, developed by Swedish chemist Svante Arrhenius in his 1884 doctoral dissertation, provided the first modern definition by linking acidic properties to the electrolytic dissociation of substances in water. This groundbreaking work explained how acids behave through the production of charged particles, earning Arrhenius the Nobel Prize in Chemistry in 1903 for his contributions to understanding electrolytes. According to the Arrhenius definition, an acid is a substance that increases the concentration of hydrogen ions (H⁺, often represented as the hydronium ion H₃O⁺ in modern notation) when dissolved in water. The general dissociation reaction for an Arrhenius acid can be expressed as:
HA(aq)→H+(aq)+A−(aq) \text{HA(aq)} \rightarrow \text{H}^+(\text{aq}) + \text{A}^-(\text{aq})
This process occurs fully for strong acids and partially for weak acids, leading to observable properties like sour taste, reaction with metals, and neutralization with bases. Representative examples include hydrochloric acid (HCl), a strong acid that completely dissociates in water to produce H⁺ and Cl⁻ ions, and acetic acid (CH₃COOH), a weak acid that partially dissociates to yield H⁺ and CH₃COO⁻ ions. These dissociations directly contribute to the increased H⁺ concentration characteristic of acidic solutions. However, the Arrhenius definition is restricted to aqueous solutions and fails to account for acidic behavior in non-aqueous solvents or for substances that exhibit acidity without producing hydrogen ions, such as certain metal cations. Later theories, like Brønsted-Lowry, expanded on this by focusing on proton transfer in various media.

Brønsted–Lowry Acids

The Brønsted–Lowry theory defines an acid as a substance that donates a proton (H⁺ ion) to another substance, termed a base, which accepts the proton. This proton-transfer mechanism forms the core of acid-base reactions under this framework, expanding applicability beyond aqueous solutions to any medium where proton donation occurs. The theory was independently proposed in 1923 by Danish chemist Johannes Nicolaus Brønsted and British chemist Thomas Martin Lowry, providing a broader perspective than earlier models by emphasizing relative proton affinity rather than specific ion production.[7] In a Brønsted–Lowry acid-base reaction, the acid (HA) donates a proton to the base (B), yielding the conjugate base (A⁻) and conjugate acid (HB⁺). This process is reversible and represented by the general equilibrium:
HA+B⇌AX−+HBX+ \ce{HA + B ⇌ A^- + HB^+}
The conjugate acid-base pair consists of species differing by one proton, such as HA and A⁻, where the strength of the acid inversely relates to the strength of its conjugate base. This theory generalizes proton dissociation in water as a specific case of broader proton transfer.[7] Certain substances exhibit amphoterism, acting as both Brønsted–Lowry acids and bases depending on the reaction conditions, due to their ability to either donate or accept protons. The bicarbonate ion (HCO₃⁻) is a classic example of an amphoteric species. As a base, it accepts a proton from water:
HCOX3X−+HX2O⇌HX2COX3+OHX− \ce{HCO3^- + H2O ⇌ H2CO3 + OH^-}
As an acid, it donates a proton to water:
HCOX3X−+HX2O⇌COX3X2−+HX3OX+ \ce{HCO3^- + H2O ⇌ CO3^{2-} + H3O^+}
These reactions highlight bicarbonate's role in buffering systems, such as in biological fluids.[8] Representative examples illustrate proton donation in non-aqueous or varied contexts. The ammonium ion (NH₄⁺) functions as a Brønsted–Lowry acid by transferring a proton to the hydroxide ion:
NHX4X++OHX−⇌NHX3+HX2O \ce{NH4^+ + OH^- ⇌ NH3 + H2O}
Here, NH₄⁺ is the acid, OH⁻ is the base, NH₃ is the conjugate base, and H₂O is the conjugate acid. Similarly, the hydrogen sulfate ion (HSO₄⁻) demonstrates amphoterism: it acts as an acid by donating a proton to water to form sulfate and hydronium ions ($ \ce{HSO4^- + H2O ⇌ SO4^{2-} + H3O^+} ),orasabasebyacceptingaprotontoformsulfuricacid(), or as a base by accepting a proton to form sulfuric acid ( \ce{HSO4^- + H2O ⇌ H2SO4 + OH^-} $), though the latter is less common. The Brønsted–Lowry framework ties directly to acid strength through the acid dissociation constant ($ K_a $), which quantifies the equilibrium position of proton donation for weak acids in solution:
Ka=[A−][H+][HA] K_a = \frac{[A^-][H^+]}{[HA]}
A larger $ K_a $ indicates a stronger tendency to donate protons, reflecting greater acid strength within this theory. This equilibrium expression underpins quantitative analysis of conjugate pair behaviors.[9]

Lewis Acids

In 1923, Gilbert N. Lewis proposed a general theory of acid-base reactions that defines a Lewis acid as any species capable of accepting an electron pair from a Lewis base to form a coordinate covalent bond, broadening the scope beyond proton transfer mechanisms.[10] This definition emphasizes the role of electron deficiency in the acid, allowing it to complete its valence shell through donation from a base.[11] The general reaction can be represented as:
A (acceptor)+:B (donor)→A–B \text{A (acceptor)} + :\text{B (donor)} \rightarrow \text{A}–\text{B}
where A is the Lewis acid and :B denotes the lone pair on the base.[12] A classic example is the reaction between boron trifluoride (BF₃) and ammonia (NH₃), where the electron-deficient boron atom in BF₃ accepts the lone pair from nitrogen in NH₃ to form the adduct F₃B–NH₃.[13] Another prominent application occurs in organic synthesis, such as Friedel-Crafts alkylation reactions, where aluminum chloride (AlCl₃) acts as a Lewis acid by coordinating with the halogen of an alkyl halide to generate a carbocation electrophile.[14] Lewis acids play crucial roles in catalysis, particularly in biological systems where metal ions like Zn²⁺ function as electron-pair acceptors to activate substrates. For instance, in the enzyme carbonic anhydrase, Zn²⁺ coordinates with water to facilitate its deprotonation, enhancing the hydration of carbon dioxide.[15] This definition extends to non-protonic species, including metal cations such as Fe³⁺, which accept electron pairs from ligands due to their high charge density, and carbocations like (CH₃)₃C⁺, which seek stabilization through electron donation.[16] Protonic acids represent a subset of Lewis acids, as the H⁺ ion itself acts as an electron-pair acceptor.[17]

Properties

Dissociation and Equilibrium

In aqueous solutions, acids dissociate by ionizing to produce hydrogen ions (H⁺) and their conjugate bases, as originally conceptualized in the Arrhenius definition of acids. For a general acid HA, this process is represented as HA ⇌ H⁺ + A⁻, where the extent of ionization determines whether the acid is strong or weak. Strong acids, such as hydrochloric acid (HCl), undergo complete dissociation in water, meaning nearly 100% of the molecules ionize to form H⁺ and Cl⁻ ions, with no significant equilibrium established. In contrast, weak acids partially ionize, resulting in an equilibrium mixture of undissociated HA, H⁺, and A⁻.[18][19][20] The equilibrium for weak acid dissociation is quantified by the acid dissociation constant, $ K_a $, defined as $ K_a = \frac{[H^+][A^-]}{[HA]} $, where the concentrations are those at equilibrium and activities are approximated by concentrations in dilute solutions. This constant reflects the position of the equilibrium; a smaller $ K_a $ indicates less dissociation and a weaker tendency to produce H⁺. For example, acetic acid (CH₃COOH) has $ K_a \approx 1.8 \times 10^{-5} $ at 25°C, meaning only a small fraction ionizes in typical solutions. Pure water also exhibits a related autoionization equilibrium: $ \ce{H2O ⇌ H^+ + OH^-} $, governed by the ion product constant $ K_w = [H^+][OH^-] = 1.0 \times 10^{-14} $ at 25°C, which establishes a baseline concentration of H⁺ and OH⁻ ions even in neutral conditions.[21][22][23] The hydrogen ion concentration from weak acid dissociation can be approximated for initial calculations when the acid concentration $ C $ is much greater than the dissociated amount, yielding $ [H^+] \approx \sqrt{K_a \cdot C} $; this simplification assumes [H⁺] = [A⁻] and negligible change in [HA] from the initial value, valid for moderately dilute solutions where dissociation is less than 5%. External factors influence this equilibrium per Le Châtelier's principle: dilution decreases concentrations of all species, shifting the equilibrium toward greater dissociation to restore balance, thereby increasing the percent ionization. Temperature changes alter $ K_a $ itself, as acid dissociation is typically endothermic; higher temperatures favor the forward reaction, increasing $ K_a $ and [H⁺].[24][25][26]

Acid Strength

Acid strength quantifies the extent to which an acid donates a proton (H⁺) in solution, primarily measured by the acid dissociation constant KaK_a, defined for the equilibrium HA⇌HX++AX−\ce{HA ⇌ H+ + A-} as Ka=[HX+][AX−][HA]K_a = \frac{[\ce{H+}][\ce{A-}]}{[\ce{HA}]}. The pKa value, given by $ \mathrm{p}K_a = -\log_{10} K_a $, provides a convenient scale where a lower pKa corresponds to a stronger acid due to greater proton donation tendency.[27] In aqueous solutions, acids are classified as strong if they fully dissociate (pKa < 0), such as hydrochloric acid (HCl, pKa ≈ -7), which exists entirely as HX+\ce{H+} and ClX−\ce{Cl-}. Weak acids, with pKa > 0, partially dissociate; for example, hydrofluoric acid (HF, pKa = 3.17) ionizes only to a limited extent due to the strong H–F bond and poor stabilization of the FX−\ce{F-} conjugate base.[28][29] Several factors influence acid strength by affecting the stability of the conjugate base or the ease of proton release. Bond strength plays a key role: weaker H–A bonds favor stronger acids, as seen in the hydrogen halides where HF (strong H–F bond) is much weaker than HI (weaker H–I bond, pKa ≈ -9). Inductive effects from electron-withdrawing groups, such as halogens on a carbon chain, stabilize the negative charge on the conjugate base by withdrawing electron density, increasing acidity (e.g., chloroacetic acid is stronger than acetic acid). Resonance stabilization is particularly effective, delocalizing the conjugate base charge over multiple atoms, as in carboxylic acids where the acetate ion's charge spreads across two oxygen atoms, making them more acidic than alcohols.[30][30][30] For polyprotic acids, which can donate multiple protons, successive pKa values increase because each subsequent conjugate base is less willing to lose a proton; for sulfuric acid (H₂SO₄), pKa₁ ≈ -3 (strong first dissociation to HSOX4X−\ce{HSO4-}) while pKa₂ ≈ 2 (weaker second dissociation to SOX4X2−\ce{SO4^2-}).[31] In non-aqueous solvents, which are less basic than water, acid strengths can differ markedly due to reduced leveling effects; for instance, in acetic acid, the order reverses from aqueous behavior, with HCl weaker than HBr (and HI strongest) as the solvent's lower proton-accepting ability allows differentiation based on inherent bond polarities and conjugate base solvation.[32] Superacids, developed in the mid-20th century, exceed the strength of concentrated sulfuric acid (H₀ ≈ -12, where H₀ is the Hammett acidity function extending pH for highly acidic media); the "magic acid" system of fluorosulfuric acid (HSO₃F) with antimony pentafluoride (SbF₅) achieves H₀ < -20, enabling protonation of weak bases like hydrocarbons.[33][34]

Nomenclature

The nomenclature of acids has evolved from early trivial names based on sensory properties or origins, such as "vinegar" for acetic acid, to systematic conventions established in the late 18th and 19th centuries by chemists like Antoine Lavoisier and Jöns Jacob Berzelius, who emphasized compositional elements, with the International Union of Pure and Applied Chemistry (IUPAC) formalizing rules in the 20th century to promote precision and universality.[35][36] This shift addressed ambiguities in pre-modern naming, where acids were often described by their sources or effects rather than structure, leading to the adoption of substitutive and additive methods that reflect molecular composition.[36] For inorganic acids, binary acids—those composed of hydrogen and a single nonmetal—are named using the prefix "hydro-" followed by the stem of the nonmetal and the suffix "-ic acid," as in hydrochloric acid for HCl.[36] Oxyacids, which include oxygen, follow traditional naming based on the corresponding oxyanion: the suffix "-ic acid" denotes the anion with more oxygen or higher oxidation state (e.g., sulfuric acid for H₂SO₄, derived from sulfate), while "-ous acid" indicates fewer oxygen atoms or lower oxidation state (e.g., sulfurous acid for H₂SO₃, from sulfite); additional prefixes like "per-" (highest oxygen, as in perchloric acid, HClO₄) and "hypo-" (lowest, as in hypochlorous acid, HClO) refine these distinctions.[36] IUPAC also endorses additive nomenclature for clarity, listing ligands alphabetically around the central atom (e.g., tetraoxidosulfuric acid for H₂SO₄), though traditional names remain widely retained.[36] Organic acids employ substitutive nomenclature, prioritizing the principal functional group as the suffix. Carboxylic acids, featuring the -COOH group, are named by identifying the longest carbon chain including the carboxyl carbon and appending "-oic acid," with the chain numbered from the carboxyl group; for instance, CH₃COOH is ethanoic acid (preferred IUPAC name, or PIN), though the retained common name acetic acid is acceptable in general use.[37] Sulfonic acids, with the -SO₃H group, similarly use the suffix "-sulfonic acid" attached to the parent hydrocarbon chain or ring, such as methanesulfonic acid for CH₃SO₃H or benzenesulfonic acid for C₆H₅SO₃H.[37] Polyprotic acids, capable of donating multiple protons, extend these rules to their anions through "hydrogen" prefixes indicating remaining ionizable hydrogens, as seen in dihydrogen phosphate for H₂PO₄⁻ (from phosphoric acid, H₃PO₄) or hydrogen phosphate for HPO₄²⁻; this convention treats partial deprotonation systematically while aligning with oxyanion naming patterns.[38] Overall, IUPAC distinguishes preferred systematic names (e.g., ethanoic acid) from retained trivial ones (e.g., acetic acid) to balance innovation with established terminology, ensuring nomenclature supports both educational and practical applications without implying acid strength differences solely through naming conventions.[37][36]

Chemical Behavior

Monoprotic and Polyprotic Acids

Monoprotic acids are those capable of donating a single proton (H⁺) per molecule in aqueous solution, resulting in a single acid dissociation equilibrium characterized by one acid dissociation constant, Ka.[39] Representative examples include hydrochloric acid (HCl), a strong monoprotic acid that fully dissociates, and acetic acid (CH₃COOH), a weak monoprotic acid with Ka ≈ 1.8 × 10−5.[39] In contrast, polyprotic acids can donate more than one proton per molecule through successive dissociation steps. Diprotic acids, such as sulfuric acid (H₂SO₄), release two protons, while triprotic acids like phosphoric acid (H₃PO₄) release three.[40] For a generic diprotic acid denoted as H₂A, the stepwise dissociation equilibria are:
HX2A⇌HX++HAX−Ka1=[HX+][HAX−][HX2A] \ce{H2A ⇌ H+ + HA-} \quad K_{a1} = \frac{[\ce{H+}][\ce{HA-}]}{[\ce{H2A}]}
HAX−⇌HX++AX2−Ka2=[HX+][AX2−][HAX−] \ce{HA- ⇌ H+ + A^{2-}} \quad K_{a2} = \frac{[\ce{H+}][\ce{A^{2-}}]}{[\ce{HA-}]}
Triprotic acids follow analogous stepwise processes for each proton.[39] The acid dissociation constants for successive steps decrease markedly (Ka1 ≫ Ka2 ≫ Ka3), so pKa1 < pKa2 < pKa3; this occurs because each subsequent proton is removed from an increasingly negatively charged anion, which experiences greater electrostatic repulsion and holds the proton more tightly.[41] For sulfuric acid, for instance, Ka1 = 1.0 × 103 while Ka2 = 1.2 × 10−2, and for phosphoric acid, Ka1 = 7.1 × 10−3, Ka2 = 6.3 × 10−8, and Ka3 = 4.2 × 10−13.[40] The relative concentrations of the various species from a polyprotic acid in solution depend on the pH, with predominance shifting across the pKa values. For phosphoric acid, the dihydrogen phosphate species (H₂PO₄⁻) predominates in solutions with pH between approximately 2 and 7, the range spanning its first and second pKa values (2.1 and 7.2).[40] A key biological example is carbonic acid (H₂CO₃), a diprotic acid formed from CO₂ and H₂O in blood, where the bicarbonate ion (HCO₃⁻) is the dominant species at physiological pH (around 7.4), contributing to the bicarbonate buffer system that stabilizes blood pH between 7.35 and 7.45.[42]

Neutralization Reactions

Neutralization reactions occur when an acid reacts with a base to form a salt and water, involving the combination of hydrogen ions (H⁺) from the acid and hydroxide ions (OH⁻) from the base to produce water.[43] The general equation for such a reaction is represented as HA + BOH → BA + H₂O, where HA is the acid, BOH is the base, BA is the salt, and H₂O is water.[43] These reactions are typically exothermic, particularly for pairs of strong acids and strong bases, where the heat of neutralization is approximately -57 kJ/mol at 25°C, reflecting the formation of water from fully dissociated ions.[43] The stoichiometry of neutralization reactions depends on the number of ionizable protons in the acid and hydroxide groups in the base. For monoprotic acids, such as hydrochloric acid (HCl), the reaction follows a 1:1 molar ratio with a monohydroxy base like sodium hydroxide (NaOH): HCl + NaOH → NaCl + H₂O.[43] Polyprotic acids require multiple equivalents of base; for example, sulfuric acid (H₂SO₄), a diprotic acid, reacts with two moles of NaOH: H₂SO₄ + 2NaOH → Na₂SO₄ + 2H₂O.[43] The salts formed in neutralization reactions derive their properties from the strengths of the parent acid and base, specifically their conjugate pairs. Salts from strong acids and strong bases, such as NaCl from HCl and NaOH, are neutral with a pH of 7 in aqueous solution.[43] In contrast, salts from strong acids and weak bases, like ammonium chloride (NH₄Cl) from HCl and NH₃, are acidic (pH < 7) due to the hydrolysis of the conjugate acid of the weak base.[44] Representative examples illustrate these principles. The reaction of HCl with NaOH produces NaCl and water, a classic strong acid-strong base neutralization.[43] Historically, neutralization has been applied in soap production through saponification, where fatty acids from animal fats or vegetable oils react with lye (NaOH) to form soap salts and glycerol.[45]

Weak Acid–Weak Base Equilibria

The reaction between a weak acid (HA) and a weak base (B) proceeds incompletely, establishing an equilibrium described by the equation:
HA+B⇌A−+HB+ HA + B \rightleftharpoons A^- + HB^+
The equilibrium constant KK for this reaction is given by K=KaKbKwK = \frac{K_a K_b}{K_w}, where KaK_a is the acid dissociation constant of HA, KbK_b is the base dissociation constant of B, and KwK_w is the ion product of water. This relationship arises because the forward reaction involves proton transfer from HA to B, with the position of equilibrium favoring the side containing the weaker acid and the weaker base (i.e., the side where the pKaK_a of the acid is higher and the pKbK_b of the base is lower).[46][47] Such equilibria form the basis of buffer solutions, which are mixtures of a weak acid and its conjugate base (or a weak base and its conjugate acid) that resist changes in pH upon addition of small amounts of strong acid or base. For instance, a buffer can be prepared by mixing acetic acid (CH₃COOH) with its conjugate base acetate (CH₃COO⁻), maintaining a stable pH through the reversible proton exchange. The pH of an acidic buffer is calculated using the Henderson-Hasselbalch equation:
pH=pKa+log⁡10[A−][HA] \text{pH} = \text{p}K_a + \log_{10} \frac{[A^-]}{[HA]}
This equation, derived from the KaK_a expression, allows prediction of buffer pH based on the ratio of conjugate base to acid concentrations, assuming activity coefficients near unity.[48] A practical example is the mixture of acetic acid and ammonia (NH₃), where the equilibrium produces acetate and ammonium ions (NH₄⁺), creating a buffer system effective around neutral pH. In biological contexts, the bicarbonate buffer system—comprising carbonic acid (H₂CO₃) and bicarbonate (HCO₃⁻)—maintains blood pH near 7.4 by buffering metabolic acids and CO₂-derived protons. Buffers are most effective within approximately pK_a ± 1 unit, where the concentrations of the acid and conjugate base are within a 10:1 ratio, maximizing resistance to pH shifts. Beyond this range, buffering capacity diminishes significantly.[49]

Measurement

Titration

Titration is an analytical technique used to determine the concentration of an acid by gradually adding a base of known concentration and monitoring the pH of the solution. The procedure typically involves placing a known volume of the acid solution in an Erlenmeyer flask and titrating it with the base from a burette, recording the pH after each incremental addition using a pH meter until the equivalence point is reached. This method relies on the neutralization reaction between the acid and base, allowing for precise quantification of the acid's molarity.[50] The resulting titration curve plots pH against the volume of base added, providing a visual representation of the reaction progress. For a strong acid titrated with a strong base, the curve is sigmoidal, characterized by a low initial pH, a gradual increase, and a sharp rise near the equivalence point due to excess base. In contrast, the curve for a weak acid titrated with a strong base features gentler slopes and plateaus, reflecting the buffering capacity of the system; the acid strength influences the curve's shape, with weaker acids producing less pronounced changes in pH. Buffer regions appear midway to the equivalence point, where the pH approximates the pK_a of the acid, as the solution contains roughly equal concentrations of the acid and its conjugate base, resisting pH changes.[51] The equivalence point occurs when the moles of acid equal the moles of base added for monoprotic acids, marking complete neutralization. For strong acid-strong base titrations, this point is at pH 7, as the resulting salt solution is neutral. In weak acid-strong base titrations, the equivalence point pH exceeds 7, typically around 8-9, because the conjugate base of the weak acid hydrolyzes to produce excess OH^-. For polyprotic acids such as H_2SO_4, the curve displays two distinct equivalence points corresponding to each proton donation, with inflection breaks at a low pH around 2–3 (first, forming HSO_4^-, determined by the pK_a of HSO_4^- ≈ 2) and pH 9 (second, forming SO_4^{2-}).[51][40] To calculate the volume of base required to reach the equivalence point for a monoprotic acid, use the formula:
Veq=Cacid×VacidCbase V_{\text{eq}} = \frac{C_{\text{acid}} \times V_{\text{acid}}}{C_{\text{base}}}
where CacidC_{\text{acid}} and CbaseC_{\text{base}} are the concentrations, and VacidV_{\text{acid}} is the initial volume of acid; this assumes a 1:1 stoichiometry and follows from the equality of moles at equivalence. For polyprotic acids, the formula is adjusted by the number of equivalents, but the principle remains based on stoichiometric balance.[52]

pH and Indicators

The concept of pH was introduced in 1909 by Danish biochemist Søren Peder Lauritz Sørensen while working at the Carlsberg Laboratory, providing a practical scale to quantify the acidity of solutions based on hydrogen ion activity.[53] Sørensen's innovation addressed the need for a logarithmic measure during biochemical research on enzyme activity in brewing.[54] pH is formally defined by the International Union of Pure and Applied Chemistry (IUPAC) as the negative base-10 logarithm of the activity of hydrogen ions in solution:
pH=−log⁡10a(H+) \mathrm{pH} = -\log_{10} a(\mathrm{H^+})
where $ a(\mathrm{H^+}) $ represents the effective concentration accounting for non-ideal behavior.[55] In dilute aqueous solutions at 25°C, this approximates to $ \mathrm{pH} = -\log_{10} [\mathrm{H^+}] $, with the scale conventionally spanning 0 (highly acidic, [H⁺] = 1 M) to 14 (highly basic, [H⁺] = 10⁻¹⁴ M), and pH 7 indicating neutrality due to water's dissociation constant $ K_w = 10^{-14} $.[55][56] Values below 0 or above 14 occur in concentrated strong acids or bases, but the 0–14 range applies to most aqueous systems under standard conditions.[56] Acid-base indicators are typically weak organic acids or bases that undergo a structural change, resulting in a visible color shift near their pKₐ value, allowing qualitative pH assessment.[57] The color transition occurs over a narrow pH interval (usually 1–2 units) where the indicator's protonated and deprotonated forms coexist in comparable amounts.[57] A common example is phenolphthalein, a weak acid with pKₐ ≈ 9.3, which remains colorless below pH 8.2 in its protonated form and turns pink above pH 10.0 in its deprotonated form due to extended conjugation in the basic state.[58][58] For precise quantitative measurement, glass pH electrodes are widely used, consisting of a thin, ion-selective glass membrane that develops a potential proportional to the external [H⁺] relative to an internal reference solution.[59] This potential adheres to the Nernst equation for the hydrogen ion half-cell:
E=E0−0.059log⁡10a(H+) E = E_0 - 0.059 \log_{10} a(\mathrm{H^+})
at 25°C, corresponding to a change of 59 mV per pH unit, with the electrode potential increasing by 59 mV as the pH decreases by one unit.[59] Despite their utility, pH measurements face limitations in non-aqueous solvents, where the absence of water alters ion activity and hydration, rendering standard scales and glass electrodes unreliable without solvent-specific calibrations or alternative probes.[60] Universal indicators, blends of multiple dyes such as methyl red, bromothymol blue, and thymol blue, overcome some precision needs by displaying a continuous color gradient across pH 1–14 (red for acidic, green for neutral, purple for basic), facilitating broad-range visual approximations without equipment.[61]

Types of Acids

Mineral Acids

Mineral acids, also known as inorganic acids, are water-soluble acids derived from inorganic minerals and lack carbon in their molecular structure.[62] They encompass binary acids (also known as hydracids or hydroacids), such as hydrochloric acid ($ \ce{HCl} ),hydrobromicacid(HBr),hydroiodicacid(HI),andhydrosulfuricacid(), hydrobromic acid (HBr), hydroiodic acid (HI), and hydrosulfuric acid ( \ce{H2S} ),andoxoacids(alsoknownasoxyacids),including[sulfuricacid](/page/Sulfuricacid)(), and oxoacids (also known as oxyacids), including [sulfuric acid](/page/Sulfuric_acid) ( \ce{H2SO4} ),[nitricacid](/page/Nitricacid)(), [nitric acid](/page/Nitric_acid) ( \ce{HNO3} ),perchloricacid(), perchloric acid ( \ce{HClO4} ),[phosphoricacid](/page/Phosphoricacid)(), [phosphoric acid](/page/Phosphoric_acid) ( \ce{H3PO4} ),andcarbonicacid(), and carbonic acid ( \ce{H2CO3} $).[63] These acids are typically strong, meaning they fully dissociate in water to release hydrogen ions, contributing to their high reactivity and corrosive nature.[64] Sulfuric acid is one of the most industrially significant mineral acids, produced via the contact process, which involves the catalytic oxidation of sulfur dioxide ($ \ce{SO2} )to[sulfurtrioxide](/page/Sulfurtrioxide)() to [sulfur trioxide](/page/Sulfur_trioxide) ( \ce{SO3} ),followedbyabsorptionin[water](/page/Water).[](https://www.chemanalyst.com/Blogs/understanding−the−production−process−of−sulphuric−acid−20)Globalproductionof[sulfuricacid](/page/Sulfuricacid)reachedapproximately200millionmetrictonsperyearinthe2020s,underscoringitsroleasacornerstoneofchemical[manufacturing](/page/Manufacturing).[](https://www.emergenresearch.com/industry−report/sulfuric−acid−market)[Nitricacid](/page/Nitricacid)issynthesizedthroughthe[Ostwaldprocess](/page/Ostwaldprocess),where[ammonia](/page/Ammonia)isoxidizedoveraplatinum−rhodiumcatalystto[nitricoxide](/page/Nitricoxide)(), followed by absorption in [water](/page/Water).[](https://www.chemanalyst.com/Blogs/understanding-the-production-process-of-sulphuric-acid-20) Global production of [sulfuric acid](/page/Sulfuric_acid) reached approximately 200 million metric tons per year in the 2020s, underscoring its role as a cornerstone of chemical [manufacturing](/page/Manufacturing).[](https://www.emergenresearch.com/industry-report/sulfuric-acid-market) [Nitric acid](/page/Nitric_acid) is synthesized through the [Ostwald process](/page/Ostwald_process), where [ammonia](/page/Ammonia) is oxidized over a platinum-rhodium catalyst to [nitric oxide](/page/Nitric_oxide) ( \ce{NO} $), then further oxidized and absorbed in water to form the acid.[65] Annual global output for nitric acid was around 58 million metric tons in 2024.[66] Hydrochloric acid is commonly produced by reacting sodium chloride (salt) with sulfuric acid, generating hydrogen chloride gas that is then dissolved in water. Its global production volume stood at about 15 million metric tons in 2024.[67] These acids exhibit high strength and corrosiveness, capable of rapidly degrading metals, tissues, and materials upon contact due to their ability to donate protons and, in some cases, act as oxidizing agents.[62] For instance, sulfuric and nitric acids are among the strongest mineral acids, with pKa values indicating near-complete dissociation, while hydrochloric acid is similarly potent but non-oxidizing.[68] Phosphoric acid, however, is relatively weaker, with multiple dissociation steps yielding a lower acidity (pKa1 ≈ 2.14), making it less corrosive and suitable for applications like fertilizer production where milder reactivity is preferred.[69] A notable mixture involving mineral acids is aqua regia, composed of concentrated hydrochloric acid and nitric acid in a 3:1 ratio, which generates nascent chlorine and nitrosyl chloride to dissolve noble metals like gold that resist individual acids.[70]

Organic Acids

Organic acids are carbon-containing compounds that exhibit acidic properties, primarily through the presence of functional groups capable of donating protons. The most prevalent class is carboxylic acids, characterized by the general formula $ \ce{RCOOH} ,whereRisanorganicsubstituent,suchasin[formicacid](/page/Formicacid)(, where R is an organic substituent, such as in [formic acid](/page/Formic_acid) ( \ce{HCOOH} $), the simplest member.[71] Sulfonic acids, with the formula $ \ce{RSO3H} ,representanotherkeyclass,exemplifiedby[methanesulfonicacid](/page/Methanesulfonicacid)(, represent another key class, exemplified by [methanesulfonic acid](/page/Methanesulfonic_acid) ( \ce{CH3SO3H} );thesearenotablystrongeracidsduetotheelectron−withdrawingsulfonylgroup.Otherclassesincludeenols,whichfeatureahydroxylgroupattachedtoacarbon−carbon[doublebond](/page/Doublebond)(); these are notably stronger acids due to the electron-withdrawing sulfonyl group. Other classes include enols, which feature a hydroxyl group attached to a carbon-carbon [double bond](/page/Double_bond) ( \ce{C=C-OH} $), though they are less stable and often exist in tautomeric equilibrium with keto forms.[72] In IUPAC nomenclature, carboxylic acids are named by replacing the -e ending of the parent alkane with -oic acid, such as ethanoic acid for $ \ce{CH3COOH} .Comparedto[mineral](/page/Mineral)acids,mostorganicacidsareweaker,withcarboxylicacidstypicallyhavingpKavaluesaround4−5,indicatingpartialdissociationin[water](/page/Water);forinstance,aceticacid(. Compared to [mineral](/page/Mineral) acids, most organic acids are weaker, with carboxylic acids typically having pKa values around 4-5, indicating partial dissociation in [water](/page/Water); for instance, acetic acid ( \ce{CH3COOH} ),foundin[vinegar](/page/Vinegar),hasapKaof4.76./CarboxylicAcids/NomenclatureofCarboxylicAcids)[](https://organicchemistrydata.org/hansreich/resources/pka/)[Solubility](/page/Solubility)in[water](/page/Water)isgenerallyhighforshort−chainvariantsduetohydrogenbonding,butdecreaseswithlongerhydrophobicchains;benzoicacid(), found in [vinegar](/page/Vinegar), has a pKa of 4.76./Carboxylic_Acids/Nomenclature_of_Carboxylic_Acids)[](https://organicchemistrydata.org/hansreich/resources/pka/) [Solubility](/page/Solubility) in [water](/page/Water) is generally high for short-chain variants due to hydrogen bonding, but decreases with longer hydrophobic chains; benzoic acid ( \ce{C6H5COOH} $), used as a food preservative, has a pKa of 4.20 and limited solubility in pure water but improved in basic conditions.[73] In contrast, sulfonic acids like methanesulfonic acid exhibit pKa values around -1.9, approaching the strength of mineral acids.[74] Many organic acids play vital roles in biological systems, often derived as intermediates in the Krebs cycle (also known as the citric acid cycle), a central metabolic pathway in aerobic organisms that generates energy through the oxidation of acetyl-CoA. Key examples include citric acid, which initiates the cycle, and succinic and malic acids, which facilitate electron transfer and substrate-level phosphorylation.[75] Halogenated organic acids, such as trifluoroacetic acid ($ \ce{CF3COOH} $), deviate from the typical weakness of carboxylic acids; its pKa of 0.23 results from the high electronegativity of the three fluorine atoms, which stabilize the conjugate base by inductive withdrawal of electron density.[76][73]

Specialized Acids

Superacids represent a class of exceptionally strong acids that exceed the acidity of pure sulfuric acid, defined by a Hammett acidity function $ H_0 $ value less than -12. These media enable the protonation of notoriously unreactive hydrocarbons, such as alkanes, which was first demonstrated in the 1960s by George A. Olah using fluorosulfuric acid-antimony pentafluoride mixtures. A prominent example is magic acid, a 1:1 molar mixture of fluorosulfuric acid (FSO₃H) and antimony pentafluoride (SbF₅), which achieves an $ H_0 $ of approximately -19.2 and facilitates the formation of alkanium ions like the ethyl cation from ethane. Olah's pioneering work on these systems, including the direct observation of protonated alkanes via NMR spectroscopy, earned him the 1994 Nobel Prize in Chemistry for contributions to carbocation chemistry. Vinylogous acids feature extended conjugation through vinyl groups, allowing the acidic proton to be delocalized over a longer chain, which modulates their reactivity compared to simple analogs. Ascorbic acid (vitamin C) exemplifies this, functioning as a vinylogous carboxylic acid where the enol hydroxyl group's acidity is enhanced by resonance involving the distant carbonyl, resulting in a pKa of about 4.1 for the enolic proton. This structural motif enables ascorbic acid's role as an antioxidant, with the conjugated system facilitating electron transfer. Nucleic acids, such as DNA and RNA, incorporate phosphoric acid derivatives in their sugar-phosphate backbones, forming phosphodiester linkages that confer polyanionic character. The phosphate groups in these biopolymers exhibit two relevant pKa values: approximately 1 for the primary dissociation (yielding the monoanion) and around 6 for the secondary dissociation (to the dianion), as observed in nucleotide monophosphates like AMP. These pKa values ensure that the backbone remains negatively charged at physiological pH, stabilizing the helical structures through electrostatic repulsion and interactions with counterions. Certain acids operate exclusively under the Lewis definition, accepting electron pairs without proton donation, and boron-based compounds like boric acid, B(OH)₃, illustrate this behavior. Boric acid acts as a weak Lewis acid by coordinating to Lewis bases such as water or fluoride via its electron-deficient boron atom, forming tetrahedral adducts, though it shows minimal Brønsted acidity with a pKa exceeding 9. This property underpins its applications in complexation chemistry. In the 2020s, superacids have found emerging roles in green chemistry, particularly for degrading persistent fluorinated pollutants. For instance, a novel superacid developed in 2023 enables the conversion of non-biodegradable perfluorocarbons, akin to Teflon, into harmless fluoride ions under mild conditions, addressing environmental contamination from fluorochemicals. Similarly, halogen-substituted silicon-based Lewis superacids, reported in 2025, offer potential for sustainable catalysis by promoting selective C-H activations without hazardous solvents.

Applications and Roles

Industrial and Catalytic Uses

Sulfuric acid (H₂SO₄) is the most widely produced industrial chemical, with approximately 55% of global output used in the manufacture of phosphate fertilizers such as superphosphate and ammonium phosphate.[77] Another significant application is in lead-acid batteries, where it serves as the electrolyte to facilitate electrochemical reactions for energy storage in vehicles and backup power systems.[78] Hydrochloric acid (HCl) plays a key role in the production of polyvinyl chloride (PVC), acting as a precursor in the synthesis of vinyl chloride monomer through processes like the balanced salt process.[79] In catalysis, acids enable essential reactions in organic synthesis and petrochemical processing. The Fischer esterification, developed by Emil Fischer and Arthur Speier, involves the acid-catalyzed reaction of carboxylic acids with alcohols to form esters, a foundational method for producing biodiesel and pharmaceutical intermediates since its introduction in 1895.[80] Zeolites, as solid acid catalysts, are extensively used in fluid catalytic cracking (FCC) units in petroleum refineries to break down heavy hydrocarbons into gasoline and diesel fractions, improving yield and selectivity through their microporous structure.[81] Notable examples of acid applications include nitric acid (HNO₃) in the nitration of organic compounds to produce explosives such as TNT and nitroglycerin, where it acts as both an oxidizing and nitrating agent.[82] Phosphoric acid (H₃PO₄) is industrially employed in the beverage sector, particularly for cola production, where it provides acidity and stabilizes the formulation during large-scale manufacturing.[83] In the 2020s, there has been a notable shift toward sustainable acid production, with bio-based carboxylic acids like succinic and lactic acid gaining traction through microbial fermentation processes to replace petroleum-derived counterparts, driven by advancements in metabolic engineering for reduced carbon footprints.[84]

Food and Physiological Roles

Acids play essential roles in food and human physiology, contributing to nutrition, preservation, and digestive processes. In nutrition, citric acid, a weak organic acid, is naturally abundant in citrus fruits such as lemons and limes, where it imparts a characteristic tart flavor and contributes to the low pH of fruit juices, typically ranging from 2 to 3.[85] This acidity not only enhances taste but also aids in the bioavailability of certain minerals. Similarly, acetic acid in vinegar, at concentrations around 5%, is widely used for pickling vegetables, where it lowers the pH to create an environment that preserves flavor and texture while extending shelf life.[86] Lactic acid, produced through bacterial fermentation of sugars in foods like yogurt, sauerkraut, and kimchi, adds a tangy profile and supports probiotic content, promoting gut health.[87] In food preservation, the low pH created by these acids inhibits bacterial growth and spoilage. For instance, maintaining a pH of 4.6 or lower in acidic foods prevents the germination of harmful bacterial spores, such as those from Clostridium botulinum, ensuring safety without high-heat processing.[88] Ascorbic acid, known as vitamin C, serves as an antioxidant in fruits and vegetables, preventing oxidation that leads to browning and nutrient loss; it is particularly effective in preserving the freshness of juices and cut produce by scavenging free radicals.[89] Physiologically, hydrochloric acid (HCl) in the stomach maintains a highly acidic environment, with concentrations around 0.1 M and a pH of 1 to 2, which is crucial for digestion. This acidity activates pepsinogen into active pepsin, the primary enzyme for breaking down proteins into peptides, and kills ingested pathogens by denaturing their proteins and disrupting cellular functions when the pH drops below 3.[90][91] The sensation of sour taste arises from hydrogen ions (H⁺) stimulating specific proton-selective channels, such as OTOP1, in taste receptor cells on the tongue, signaling acidity to the brain.[92] Deficiencies in stomach acid, known as hypochlorhydria, impair these functions and are linked to malabsorption of nutrients, including proteins, vitamins (such as B12), and minerals like iron and calcium, potentially leading to deficiencies and digestive disorders, as well as an increased risk of gastrointestinal infections due to impaired killing of ingested pathogens.[93][91]

Biological and Environmental Significance

In biological systems, acids play essential roles in protein structure and function. Amino acids such as aspartic acid, which contains an acidic side chain, contribute to the overall charge and folding of proteins, often acting as proton donors in enzyme active sites to facilitate catalysis.[94] In uricotelic non-mammalian vertebrates such as birds and reptiles, uric acid serves as the primary nitrogenous waste product, allowing birds to excrete ammonia-derived waste in a semi-solid form that conserves water and minimizes toxicity during flight and arid adaptations.[95] Environmentally, acid rain—precipitation with a pH below 5.6 resulting from atmospheric reactions involving sulfur dioxide (SO₂) and nitrogen oxides (NOₓ)—has significantly impacted ecosystems since gaining widespread scientific awareness in the 1970s.[96][97] These acids leach essential nutrients like calcium and magnesium from forest soils while mobilizing toxic aluminum, leading to reduced tree growth, defoliation, and biodiversity loss in affected regions such as the northeastern United States and Europe.[98][99] Ocean acidification represents another critical environmental threat, where increased atmospheric CO₂ dissolves in seawater to form carbonic acid (H₂CO₃), causing surface ocean pH to drop by approximately 0.1 units since the Industrial Revolution.[100] This shift, with research intensifying in the post-2000s era, impairs shell formation in marine organisms like shellfish and corals by reducing carbonate ion availability, disrupting food webs and coastal economies.[101] The evolutionary origins of life may trace back to acidic conditions in a primordial soup, where prebiotic chemistry in warm, acidic pools facilitated the synthesis of organic molecules like amino acids, as hypothesized in early experiments simulating Earth's early atmosphere.[102] Remediation efforts for acidified ecosystems often involve applying limestone (calcium carbonate) to neutralize acidity in soils and waters, restoring pH and nutrient balance in forests and lakes affected by acid rain. Nucleic acids, such as DNA and RNA, exemplify specialized biological acids that store and transmit genetic information essential for all life forms.[103]

Safety and Health Impacts

Household and Laboratory Handling

Acids commonly found in households include sulfuric acid (H₂SO₄) used in battery fluid and some drain cleaners, and hydrochloric acid (HCl) present in toilet bowl cleaners and pool maintenance products.[104][105] In household settings, acids should be stored in original containers or compatible plastic or glass bottles, kept in a cool, dry, well-ventilated area away from bases, oxidizers, and foodstuffs, with secondary containment like trays to catch spills.[106][107] Before disposal, dilute acids with water and neutralize if possible, following local regulations for hazardous waste. When handling household acids, wear chemical-resistant gloves, safety goggles, and protective clothing to prevent skin and eye contact, as acids can cause corrosive burns due to their proton-donating properties.[108] For spills, immediately neutralize with a mild base like baking soda (sodium bicarbonate) to form a salt and water, then absorb and dispose properly while ventilating the area. In laboratories, acids are stored in dedicated corrosive-resistant cabinets that are ventilated and equipped with spill containment trays, using glass or polyethylene containers compatible with specific acids to prevent reactions.[109][110] Laboratory handling requires personal protective equipment including nitrile or neoprene gloves, chemical splash goggles, face shields, lab coats, and closed-toe shoes; respiratory protection may be needed for volatile acids.[108][111] Procedures for volatile acids like nitric acid (HNO₃) mandate use within a chemical fume hood to contain fumes and vapors, ensuring the sash is lowered and airflow is verified before starting work.[112] When diluting concentrated acids, always add the acid to water slowly while stirring, never the reverse, to avoid exothermic splattering and potential explosions.[113] For first aid in both settings, flush skin or eyes exposed to acids with copious lukewarm water for at least 15-20 minutes, removing contaminated clothing, and seek immediate medical attention.[114][115] In cases of acid ingestion, do not induce vomiting; instead, dilute by giving small amounts of water or milk if the person is conscious and able to swallow, then contact poison control or emergency services immediately.[116]

Acidity in Human Physiology

In human physiology, acidity plays a critical role in maintaining the body's pH balance, particularly in the blood, where the normal range is 7.35 to 7.45. This narrow range is primarily regulated by the bicarbonate buffer system, which involves the equilibrium between carbonic acid (H₂CO₃) and bicarbonate ions (HCO₃⁻), helping to neutralize excess acids or bases produced during metabolism.[49][42] Deviations below pH 7.35 result in acidosis, a condition that can impair enzyme function and cellular processes; for instance, ketoacidosis, often seen in uncontrolled diabetes, arises from the accumulation of ketone bodies that lower blood pH.[117][118] The lungs and kidneys are essential organs for pH regulation, controlling the levels of carbon dioxide (CO₂) and bicarbonate (HCO₃⁻). The respiratory system adjusts pH rapidly by altering ventilation rates to expel CO₂, a volatile acid formed from carbonic acid dissociation, while the kidneys provide longer-term control by reabsorbing or excreting HCO₃⁻ and excreting hydrogen ions over hours to days.[42][119] Imbalances in these mechanisms contribute to acid-base disorders; for example, in metabolic acidosis like ketoacidosis, compensatory hyperventilation reduces CO₂ to mitigate the pH drop.[117] Acids also have direct medical applications in physiology. Aspirin, or acetylsalicylic acid, serves as an analgesic by inhibiting cyclooxygenase enzymes, reducing prostaglandin synthesis that sensitizes pain receptors, thereby alleviating pain and inflammation.[120][121] In cases of achlorhydria, a condition characterized by insufficient hydrochloric acid (HCl) secretion in the stomach, HCl therapy can be administered to restore gastric acidity, aiding digestion and pathogen defense.[122][123] Diseases related to acidity imbalances highlight its physiological impact. Gastroesophageal reflux disease (GERD) occurs when excess stomach acid refluxes into the esophagus, causing irritation, inflammation, and symptoms like heartburn due to the corrosive nature of gastric HCl.[124][125] Gout, conversely, stems from hyperuricemia leading to the deposition of uric acid crystals in joints, triggering acute inflammatory arthritis.[126] In the 2020s, proton pump inhibitors (PPIs), which suppress acid production, have seen widespread use for managing conditions like GERD.

Environmental Effects

Acid rain primarily forms when emissions of sulfur dioxide (SO₂) and nitrogen oxides (NOₓ, including NO₂) from sources such as power plants and vehicles react with water vapor, oxygen, and other chemicals in the atmosphere to produce sulfuric and nitric acids, which then fall as precipitation.[97] These acids, derived from mineral sources, lower the pH of rain to levels as low as 4.2-5.0, far below the typical 5.6 of unpolluted rain.[97] In soils, this acidity leaches essential minerals and nutrients while mobilizing toxic aluminum, which binds to roots and impairs plant growth; in water bodies, it increases aluminum concentrations that disrupt fish gills and eggs, leading to die-offs in lakes and streams where pH drops below 5.0.[127][128] Aquatic ecosystems suffer cascading effects, with sensitive species like salmon and trout populations declining sharply in acidified waters.[129] Regulatory efforts have significantly curbed these impacts. The U.S. Clean Air Act, initially enacted in 1970 with key amendments in 1990 targeting acid rain, has reduced SO₂ emissions from power plants by over 90% since 1990 through cap-and-trade programs and flue gas desulfurization technologies like scrubbers, which capture up to 95% of SO₂ before release.[130] Internationally, the 1999 Gothenburg Protocol under the UNECE Convention on Long-Range Transboundary Air Pollution mandates emission reductions for SO₂ (up to 63% by 2010 from 1990 levels in participating countries), NOₓ (41%), and other pollutants to combat acidification across Europe and North America.[131] These measures, including widespread adoption of scrubbers in coal-fired plants, have led to measurable recovery in affected ecosystems.[132] A notable case is Germany's Black Forest, where acid rain in the 1980s damaged nearly 50% of trees through needle loss, soil degradation, and bark erosion, sparking widespread "Waldsterben" (forest dieback) alarm and prompting stricter emission controls.[133] Post-regulation, SO₂ emissions in central Europe fell by about 90% from 1980s levels, enabling forest recovery; by the 2000s, tree health improved significantly, with reduced crown damage and soil pH stabilization observed in monitoring plots.[134] Beyond atmospheric acids, rising atmospheric CO₂ drives ocean acidification by forming carbonic acid upon dissolution in seawater, with models projecting a global surface pH decline of 0.3 to 0.4 units by 2100 under high-emission scenarios, threatening shellfish and coral calcification.[135][136] As of 2025, concerns have arisen over a potential return of acid rain in the U.S. due to rollbacks in emission regulations, though overall reductions remain significant.[137][138]

References

  1. IUPAC - acid (A00071)
  2. [PDF] Three Types of Acids
  3. Sulfur Statistics and Information | U.S. Geological Survey - USGS.gov
  4. Understanding the Uses of Acids in Food Processing - Bell Chem
  5. Brønsted-Lowry Acids and Bases – Chemistry and the Environment ...
  6. https://chem.libretexts.org/Bookshelves/General_Chemistry/Map%3A_General_Chemistry_(Petrucci_et_al.
  7. [PDF] Lawrence Berkeley National Laboratory - College of Chemistry
  8. The Lewis Definitions of Acids and Bases
  9. [PDF] INTRODUCTION TO LEWIS ACID-BASE CHEMISTRY
  10. 3.2: Acids and Bases - The Lewis Definition - Chemistry LibreTexts
  11. 4.1.4: Alkylation and Acylation of Aromatic Rings - The Friedel-Crafts ...
  12. 12.5: Zinc as Lewis Acid and Template - Chemistry LibreTexts
  13. Lewis Acids/Bases; HSAB Theory
  14. 3.2: Brønsted and Lewis Acids and Bases - Chemistry LibreTexts
  15. 6.3: Strength of acids and bases - Chemistry LibreTexts
  16. Strong Acids and Water Dissociation
  17. Weak Acid Equilibrium
  18. Weak acid-base equilibria (article) - Khan Academy
  19. Ka Table - Acid
  20. The Auto-Ionization of Water, K w
  21. 13.3: Finding the pH of weak Acids, Bases, and Salts
  22. https://scienceready.com.au/pages/pka
  23. Le Chatelier's Principle Fundamentals - Chemistry LibreTexts
  24. pKa and Dissociation Equilibrium - Shimadzu Scientific Instruments
  25. [PDF] D.H. Ripin, D.A. Evans pKa's of Inorganic and Oxo-Acids Chem 206
  26. Table of Acid and Base Strength
  27. 1.4.1: Factors That Determine Acid Strength
  28. Appendix C: Dissociation Constants and pKa Values for Acids at 25°C
  29. 14.3 Relative Strengths of Acids and Bases - Chemistry 2e | OpenStax
  30. What's the strongest acid in the world? | Article - RSC Education
  31. [PDF] Fluorinated superacidic systems - Société Chimique de France
  32. [PDF] A Basic History of Acid— From Aristotle to Arnold
  33. [PDF] Nomenclature of Inorganic Chemistry | IUPAC
  34. [PDF] Brief Guide to the Nomenclature of Organic Chemistry - IUPAC
  35. Brief Guide to Nomenclature of Inorganic Chemistry
  36. 15.5 Polyprotic Acids – Chemistry Fundamentals - UCF Pressbooks
  37. Diprotic and Triprotic Acids and Bases
  38. [PDF] POLYPROTIC ACIDS AND BASES: Very important!
  39. Physiology, Acid Base Balance - StatPearls - NCBI Bookshelf - NIH
  40. Neutralization - Chemistry LibreTexts
  41. Enthalpy Change of Neutralization - Chemistry LibreTexts
  42. What are salts? - ChemTeam
  43. The Chemistry of Soap Making
  44. 8.4 Acid-Base Reactions and Buffers - AP Chem - Fiveable
  45. https://chem.libretexts.org/Bookshelves/General_Chemistry/Map%3A_Chemistry_-_The_Central_Science_(Brown_et_al.
  46. [PDF] Experiment # 9: The Henderson-Hasselbalch Equation - ULM
  47. Acid–base balance: a review of normal physiology - PMC - NIH
  48. [PDF] Experiment 6 Titration II – Acid Dissociation Constant
  49. https://chem.libretexts.org/Bookshelves/Physical_and_Theoretical_Chemistry_Textbook_Maps/Supplemental_Modules_(Physical_and_Theoretical_Chemistry
  50. [PDF] Chapter 9
  51. One-Hundred Years of pH | Journal of Chemical Education
  52. Søren Sørensen - Science History Institute
  53. pH (P04524) - IUPAC Gold Book
  54. pH | US EPA
  55. Phenolphthalein | C20H14O4 | CID 4764 - PubChem - NIH
  56. Glass Electrode - an overview | ScienceDirect Topics
  57. Determination of pH in Non-Aqueous Solutions - HORIBA
  58. Mineral Acids - Division of Research Safety
  59. Basic Chemistry Terms - CHE 121: General Chemistry - Inorganic
  60. Mineral Acid Definition and List - Science Notes
  61. Understanding the Production Process of Sulphuric Acid
  62. https://www.emergenresearch.com/industry-report/sulfuric-acid-market
  63. The Principle and Mechanism behind Ostwald's Process - BYJU'S
  64. Nitric Acid Market Size, Share, Growth and Forecast 2035
  65. Hydrochloric Acid Market Size, Share, Growth & Forecast 2035
  66. Strong Acids-Mineral Acids, Strong Acids and Their Varieties ...
  67. Phosphoric acid - The Essential Chemical Industry
  68. Aqua Regia - - Division of Research Safety | Illinois
  69. https://chem.libretexts.org/Bookshelves/Organic_Chemistry/Organic_Chemistry_(Morsch_et_al.
  70. [PDF] Carboxylic Acid Structure and Chemistry: Part 2
  71. Bordwell pKa Table - Organic Chemistry Data
  72. [PDF] Conductivity, Viscosity, Spectroscopic Properties of Organic Sulfonic ...
  73. Physiology, Krebs Cycle - StatPearls - NCBI Bookshelf
  74. Sulfuric Acid Market Size, Share And Growth Report, 2030
  75. What Is Battery Acid? Sulfuric Acid Facts - ThoughtCo
  76. Uses Of Hydrochloric Acid - BYJU'S
  77. Fischer Esterification - Chemistry LibreTexts
  78. (PDF) Application of Zeolites in Petroleum Refining - ResearchGate
  79. Nitric acid - The Essential Chemical Industry
  80. Phosphoric Acid Ingredients: Food & Tech Grade - Innophos
  81. unlocking the potential of key organic acids for sustainable industrial ...
  82. Global Inorganic Chemicals Market Size, Trends, Share 2033
  83. Quantitative Assessment of Citric Acid in Lemon Juice, Lime ... - NIH
  84. Pickling basics | UMN Extension
  85. Lactic Acid Fermentations - NCBI
  86. [PDF] Food Safety A to Z Reference Guide - FDA
  87. Vitamin C - Health Professional Fact Sheet
  88. Physiology, Pepsin - StatPearls - NCBI Bookshelf - NIH
  89. Influence of Gastric Acid on Susceptibility to Infection with Ingested ...
  90. The Cellular and Molecular Basis of Sour Taste - PubMed Central
  91. Hypochlorhydria (Low Stomach Acid): Symptoms, Tests, Treatment
  92. Amino Acids - Aspartic Acid - The Biology Project
  93. Nitrogenous Waste in Birds and Reptiles: Uric Acid - OpenEd CUNY
  94. The Acid Rain Problem | Penn State University
  95. What is Acid Rain? | US EPA
  96. How to Assess Air Pollution's Impacts on Forests
  97. [PDF] Acid deposition effects on soil chemistry and forest growth on the ...
  98. Ocean acidification | National Oceanic and Atmospheric Administration
  99. Ocean Acidification | Smithsonian Ocean
  100. Primordial soup or vinaigrette: did the RNA world evolve at acidic pH?
  101. Biochemistry, Essential Amino Acids - StatPearls - NCBI Bookshelf
  102. 15.2: Common Acids and Their Uses - Chemistry LibreTexts
  103. 10 most common acids and their uses - Monarch Chemicals
  104. [PDF] Household Hazardous Waste Storage Recommendations | EPA
  105. How to Store Acids Safely: A Comprehensive Guide
  106. [PDF] Personal Protective Equipment - OSHA
  107. [PDF] Chemical Safety Guide, 5th Ed - ORS
  108. Chemical Storage - Environmental Health and Safety
  109. Engineering Controls Database - Museums (Acids and Alkalis) - CDC
  110. [DOC] Nitric Acid Standard Operating Procedure Template
  111. Safely Diluting Acids (and Bases) - AL Research Support
  112. First aid for acid burns | British Red Cross
  113. First Aid for Chemical Exposures - CCOHS
  114. FIRST AID MEASURES FOR PQISGNINe
  115. Anion Gap and Non-Anion Gap Metabolic Acidosis - StatPearls - NCBI
  116. [PDF] Metabolic Acidosis: Got Bicarbonate?
  117. Acid-Base Homeostasis - PMC - PubMed Central - NIH
  118. The mechanism of action of aspirin - PubMed
  119. Aspirin | C9H8O4 | CID 2244 - PubChem - NIH
  120. Achlorhydria - StatPearls - NCBI Bookshelf - NIH
  121. Management of upper gastrointestinal symptoms in patients ... - NIH
  122. Gastroesophageal reflux disease (GERD) - Symptoms and causes
  123. Gastroesophageal Reflux Disease (GERD) - StatPearls - NCBI - NIH
  124. Gout - StatPearls - NCBI Bookshelf
  125. Acid Rain and Water | U.S. Geological Survey - USGS.gov
  126. Acid rain causes, effects, and solutions | National Geographic
  127. Acid Rain - Lake Scientist
  128. Progress Cleaning the Air and Improving People's Health | US EPA
  129. Protocol to Abate Acidification, Eutrophication and Ground-level ...
  130. [PDF] Review of the Gothenburg Protocol - EMEP
  131. Acid Rain: Lessons from Germany's Black Forest
  132. Recovery from acidification in central Europe—observed and ...
  133. Global Surface Ocean Acidification Indicators From 1750 to 2100
  134. Projected climate change impact on oceanic acidification - PMC
User Avatar
No comments yet.