De Wikipedia, la enciclopedia libre
Saltar a navegación Saltar a búsqueda

El zinc , un metal típico, reacciona con el ácido clorhídrico , un ácido típico.

Un ácido es una molécula o ión capaz de donar un protón (ion hidrógeno H + ) (un ácido de Brønsted-Lowry ) o, alternativamente, capaz de formar un enlace covalente con un par de electrones (un ácido de Lewis ). [1]

La primera categoría de ácidos son los donantes de protones o ácidos de Brønsted-Lowry . En el caso especial de las soluciones acuosas , los donantes de protones forman el ion hidronio H 3 O + y se conocen como ácidos de Arrhenius . Brønsted y Lowry generalizaron la teoría de Arrhenius para incluir disolventes no acuosos. Un ácido de Brønsted o Arrhenius generalmente contiene un átomo de hidrógeno unido a una estructura química que sigue siendo energéticamente favorable después de la pérdida de H + .

Los ácidos de Arrhenius acuosos tienen propiedades características que proporcionan una descripción práctica de un ácido. [2] Los ácidos forman soluciones acuosas con un sabor agrio, pueden volverse rojo tornasol azul y reaccionar con bases y ciertos metales (como el calcio ) para formar sales . La palabra ácido se deriva del latín acidus / acēre , que significa 'amargo'. [3] Una solución acuosa de un ácido tiene un pH menor que 7 y coloquialmente también se la conoce como "ácido" (como en "disuelto en ácido"), mientras que la definición estricta se refiere solo al soluto . [1]Un pH más bajo significa una acidez más alta y, por lo tanto, una concentración más alta de iones de hidrógeno positivos en la solución . Se dice que los productos químicos o sustancias que tienen la propiedad de un ácido son ácidos .

Los ácidos acuosos comunes incluyen ácido clorhídrico (una solución de cloruro de hidrógeno que se encuentra en el ácido gástrico en el estómago y activa las enzimas digestivas ), ácido acético (el vinagre es una solución acuosa diluida de este líquido), ácido sulfúrico (usado en baterías de automóviles ), y ácido cítrico (que se encuentra en frutas cítricas). Como muestran estos ejemplos, los ácidos (en el sentido coloquial) pueden ser soluciones o sustancias puras, y pueden derivarse de ácidos (en el sentido estricto [1] ) que son sólidos, líquidos o gases. Los ácidos fuertes y algunos ácidos débiles concentrados son corrosivos., pero hay excepciones como los carboranos y el ácido bórico .

La segunda categoría de ácidos son los ácidos de Lewis , que forman un enlace covalente con un par de electrones. Un ejemplo es el trifluoruro de boro (BF 3 ), cuyo átomo de boro tiene un orbital vacante que puede formar un enlace covalente al compartir un solo par de electrones en un átomo en una base, por ejemplo, el átomo de nitrógeno en el amoníaco (NH 3 ). Lewis consideró esto como una generalización de la definición de Brønsted, de modo que un ácido es una especie química que acepta pares de electrones directamente o liberando protones (H +) en la solución, que luego aceptan pares de electrones. Sin embargo, el cloruro de hidrógeno, el ácido acético y la mayoría de los demás ácidos de Brønsted-Lowry no pueden formar un enlace covalente con un par de electrones y, por lo tanto, no son ácidos de Lewis. [4] Por el contrario, muchos ácidos de Lewis no son ácidos de Arrhenius o Brønsted-Lowry. En la terminología moderna, un ácido es implícitamente un ácido de Brønsted y no un ácido de Lewis, ya que los químicos casi siempre se refieren a un ácido de Lewis explícitamente como un ácido de Lewis . [4]

Definiciones y conceptos [ editar ]

Las definiciones modernas se refieren a las reacciones químicas fundamentales comunes a todos los ácidos.

La mayoría de los ácidos que se encuentran en la vida cotidiana son soluciones acuosas o pueden disolverse en agua, por lo que las definiciones de Arrhenius y Brønsted-Lowry son las más relevantes.

La definición de Brønsted-Lowry es la más utilizada; A menos que se especifique lo contrario, se supone que las reacciones ácido-base implican la transferencia de un protón (H + ) de un ácido a una base.

Los iones de hidronio son ácidos según las tres definiciones. Aunque los alcoholes y las aminas pueden ser ácidos de Brønsted-Lowry, también pueden funcionar como bases de Lewis debido a los pares de electrones solitarios en sus átomos de oxígeno y nitrógeno.

Ácidos de Arrhenius [ editar ]

Svante Arrhenius

En 1884, Svante Arrhenius atribuyó las propiedades de la acidez a los iones de hidrógeno (H + ), posteriormente descritos como protones o hidrones . Un ácido de Arrhenius es una sustancia que, cuando se agrega al agua, aumenta la concentración de iones H + en el agua. [4] [5] Tenga en cuenta que los químicos a menudo escriben H + ( aq ) y se refieren al ion hidrógeno cuando describen reacciones ácido-base, pero el núcleo de hidrógeno libre, un protón , no existe solo en el agua, existe como ion hidronio (H 3 O +) u otras formas (H 5 O 2 + , H 9 O 4 + ). Por lo tanto, un ácido de Arrhenius también se puede describir como una sustancia que aumenta la concentración de iones hidronio cuando se agrega al agua. Los ejemplos incluyen sustancias moleculares como el cloruro de hidrógeno y el ácido acético.

Una base de Arrhenius , por otro lado, es una sustancia que aumenta la concentración de iones de hidróxido (OH - ) cuando se disuelve en agua. Esto disminuye la concentración de hidronio porque los iones reaccionan para formar moléculas de H 2 O:

H 3 O+
(aq)
+ OH-
(aq)
⇌ H 2 O (l) + H 2 O (l)

Debido a este equilibrio, cualquier aumento en la concentración de hidronio va acompañado de una disminución en la concentración de hidróxido. Por lo tanto, también se podría decir que un ácido de Arrhenius es uno que disminuye la concentración de hidróxido, mientras que una base de Arrhenius la aumenta.

En una solución ácida, la concentración de iones hidronio es superior a 10-7 moles por litro. Dado que el pH se define como el logaritmo negativo de la concentración de iones hidronio, las soluciones ácidas tienen un pH inferior a 7.

Ácidos de Brønsted-Lowry[ editar ]

El ácido acético , un ácido débil , dona un protón (ión hidrógeno, resaltado en verde) al agua en una reacción de equilibrio para dar el ión acetato y el ión hidronio . Rojo: oxígeno, negro: carbono, blanco: hidrógeno.

Si bien el concepto de Arrhenius es útil para describir muchas reacciones, también es bastante limitado en su alcance. En 1923, los químicos Johannes Nicolaus Brønsted y Thomas Martin Lowry reconocieron de forma independiente que las reacciones ácido-base implican la transferencia de un protón. Un ácido de Brønsted-Lowry (o simplemente ácido de Brønsted) es una especie que dona un protón a una base de Brønsted-Lowry. [5] La teoría ácido-base de Brønsted-Lowry tiene varias ventajas sobre la teoría de Arrhenius. Considere las siguientes reacciones del ácido acético (CH 3 COOH), el ácido orgánico que le da al vinagre su sabor característico:

CH
3
COOH
+ H
2
O
CH
3
ARRULLO-
+ H
3
O+
CH
3
COOH
+ NH
3
CH
3
ARRULLO-
+ NH+
4

Ambas teorías describen fácilmente la primera reacción: el CH 3 COOH actúa como un ácido de Arrhenius porque actúa como una fuente de H 3 O + cuando se disuelve en agua, y actúa como un ácido de Brønsted al donar un protón al agua. En el segundo ejemplo, el CH 3 COOH sufre la misma transformación, en este caso donando un protón al amoníaco (NH 3 ), pero no se relaciona con la definición de Arrhenius de un ácido porque la reacción no produce hidronio. Sin embargo, el CH 3 COOH es un ácido de Arrhenius y de Brønsted-Lowry.

La teoría de Brønsted-Lowry se puede utilizar para describir reacciones de compuestos moleculares en solución no acuosa o en fase gaseosa. El cloruro de hidrógeno (HCl) y el amoníaco se combinan en varias condiciones diferentes para formar cloruro de amonio , NH 4 Cl. En solución acuosa, el HCl se comporta como ácido clorhídrico y existe como iones hidronio y cloruro. Las siguientes reacciones ilustran las limitaciones de la definición de Arrhenius:

  1. H 3 O+
    (aq)
    + Cl-
    (aq)
    + NH 3 → Cl-
    (aq)
    + NH+
    4
    (aq) + H 2 O
  2. HCl (benceno) + NH 3 (benceno) → NH 4 Cl (s)
  3. HCl (g) + NH 3 (g) → NH 4 Cl (s)

Al igual que con las reacciones del ácido acético, ambas definiciones funcionan para el primer ejemplo, donde el agua es el solvente y el ion hidronio está formado por el soluto de HCl. Las siguientes dos reacciones no implican la formación de iones, pero siguen siendo reacciones de transferencia de protones. En la segunda reacción, el cloruro de hidrógeno y el amoníaco (disuelto en benceno ) reaccionan para formar cloruro de amonio sólido en un solvente de benceno y en la tercera reacción, el HCl gaseoso y el NH 3 se combinan para formar el sólido.

Ácidos de Lewis [ editar ]

Un tercer concepto, sólo marginalmente relacionado, fue propuesto en 1923 por Gilbert N. Lewis , que incluye reacciones con características ácido-base que no implican una transferencia de protones. Un ácido de Lewis es una especie que acepta un par de electrones de otra especie; en otras palabras, es un aceptor de pares de electrones. [5] Las reacciones ácido-base de Brønsted son reacciones de transferencia de protones, mientras que las reacciones ácido-base de Lewis son transferencias de pares de electrones. Muchos ácidos de Lewis no son ácidos de Brønsted-Lowry. Compare cómo se describen las siguientes reacciones en términos de química ácido-base:

En la primera reacción, un ion fluoruro , F - , cede un par de electrones al trifluoruro de boro para formar el producto tetrafluoroborato . El fluoruro "pierde" un par de electrones de valencia porque los electrones compartidos en el enlace B-F están ubicados en la región del espacio entre los dos núcleos atómicos y, por lo tanto, están más distantes del núcleo del fluoruro que en el ion fluoruro solitario. BF 3es un ácido de Lewis porque acepta el par de electrones del fluoruro. Esta reacción no se puede describir en términos de la teoría de Brønsted porque no hay transferencia de protones. La segunda reacción se puede describir usando cualquiera de las teorías. Un protón se transfiere de un ácido de Brønsted no especificado al amoníaco, una base de Brønsted; alternativamente, el amoníaco actúa como una base de Lewis y transfiere un par de electrones solitarios para formar un enlace con un ion hidrógeno. La especie que gana el par de electrones es el ácido de Lewis; por ejemplo, el átomo de oxígeno en H 3 O + gana un par de electrones cuando uno de los enlaces H — O se rompe y los electrones compartidos en el enlace se localizan en el oxígeno. Dependiendo del contexto, un ácido de Lewis también puede describirse como un oxidante o un electrófilo.. Los ácidos orgánicos de Brønsted, como el ácido acético, cítrico u oxálico, no son ácidos de Lewis. [4] Se disocian en agua para producir un ácido de Lewis, H + , pero al mismo tiempo también producen una cantidad igual de una base de Lewis (acetato, citrato u oxalato, respectivamente, para los ácidos mencionados). Este artículo trata principalmente de ácidos de Brønsted en lugar de ácidos de Lewis.

Disociación y equilibrio [ editar ]

Las reacciones de los ácidos a menudo se generalizan en la forma HA ⇌ H + + A - , donde HA representa el ácido y A - es la base conjugada . Esta reacción se conoce como protólisis . La forma protonada (HA) de un ácido también se denomina a veces ácido libre . [6]

Los pares de conjugado ácido-base difieren en un protón y pueden interconvertirse mediante la adición o eliminación de un protón ( protonación y desprotonación , respectivamente). Tenga en cuenta que el ácido puede ser la especie cargada y la base conjugada puede ser neutra, en cuyo caso el esquema de reacción generalizado podría escribirse como HA + ⇌ H + + A. En solución existe un equilibrio entre el ácido y su base conjugada. La constante de equilibrio K es una expresión de las concentraciones de equilibrio de las moléculas o los iones en solución. Los paréntesis indican concentración, de modo que [H 2 O] significa la concentración de H 2 O. 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 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 ions to acid will be higher for the stronger acid as the stronger acid has a greater tendency to lose its proton. Because the range of possible values for 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:

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]

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 more negative 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. 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.[7] The relative acceptor strength of Lewis acids toward a series of bases, versus other Lewis acids, can be illustrated by C-B plots.[8][9] 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]

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]

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).

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]

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]

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.[10] 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.[11] 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.[12]

Equivalence points[edit]

Due to the successive dissociation processes, there are two equivalence points in the titration curve of a diprotic acid.[13] The first equivalence point occurs when all first hydrogen ions from the first ionization are titrated.[14] 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 hydrogen ions 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.[14] 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.[15] Each segment of the curve which 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]

Acids exist universally in our life. There are both numerous kinds of natural acid compounds with biological functions and massive synthesized acids which are used in many ways.

In industry[edit]

Acids are fundamental reagents in treating almost all processes in today's industry. Sulfuric acid, a diprotic acid, is the most widely used acid in industry, which 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.[16] According to the statistics data in 2011, the annual production of sulfuric acid was around 200 million tonnes in the world.[17] 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.[18]

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.[19] Other weak acids serve as buffers with their conjugate bases to keep the body's pH from undergoing large scale changes which would be harmful to cells.[20] The rest of the dicarboxylic acids also participate in the synthesis of various biologically important compounds in human bodies.

Acid catalysis[edit]

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.[21]

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) which 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 which 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]

  • Hydrogen halides and their solutions: hydrofluoric acid (HF), hydrochloric acid (HCl), hydrobromic acid (HBr), hydroiodic acid (HI)
  • Halogen oxoacids: hypochlorous acid (HClO), chlorous acid (HClO2), chloric acid (HClO3), perchloric acid (HClO4), and corresponding analogs for bromine and iodine
    • Hypofluorous acid (HFO), the only known oxoacid for fluorine.
  • Sulfuric acid (H2SO4)
  • Fluorosulfuric acid (HSO3F)
  • Nitric acid (HNO3)
  • Phosphoric acid (H3PO4)
  • Fluoroantimonic acid (HSbF6)
  • Fluoroboric acid (HBF4)
  • Hexafluorophosphoric acid (HPF6)
  • Chromic acid (H2CrO4)
  • Boric acid (H3BO3)

Sulfonic acids[edit]

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

  • Methanesulfonic acid (or mesylic acid, CH3SO3H)
  • Ethanesulfonic acid (or esylic acid, CH3CH2SO3H)
  • Benzenesulfonic acid (or besylic acid, C6H5SO3H)
  • p-Toluenesulfonic acid (or tosylic acid, CH3C6H4SO3H)
  • Trifluoromethanesulfonic acid (or triflic acid, CF3SO3H)
  • Polystyrene sulfonic acid (sulfonated polystyrene, [CH2CH(C6H4)SO3H]n)

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.

  • Acetic acid (CH3COOH)
  • Citric acid (C6H8O7)
  • Formic acid (HCOOH)
  • Gluconic acid HOCH2-(CHOH)4-COOH
  • Lactic acid (CH3-CHOH-COOH)
  • Oxalic acid (HOOC-COOH)
  • Tartaric acid (HOOC-CHOH-CHOH-COOH)

Halogenated carboxylic acids[edit]

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

  • Fluoroacetic acid
  • Trifluoroacetic acid
  • Chloroacetic acid
  • Dichloroacetic acid
  • Trichloroacetic 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.

  • Ascorbic acid

Nucleic acids[edit]

  • Deoxyribonucleic acid (DNA)
  • Ribonucleic acid (RNA)

References[edit]

  1. ^ a b c IUPAC Gold Book - acid
  2. ^ Petrucci, R. H.; Harwood, R. S.; Herring, F. G. (2002). General Chemistry: Principles and Modern Applications (8th ed.). Prentice Hall. p. 146. ISBN 0-13-014329-4.
  3. ^ Merriam-Webster's Online Dictionary: acid
  4. ^ a b c d Otoxby, D. W.; Gillis, H. P.; Butler, L. J. (2015). Principles of Modern Chemistry (8th ed.). Brooks Cole. p. 617. ISBN 978-1305079113.
  5. ^ a b c d Ebbing, Darrell; Gammon, Steven D. (1 January 2016). General Chemistry (11th ed.). Cengage Learning. ISBN 9781305887299.
  6. ^ Stahl PH, Nakamo M (2008). "Pharmaceutical Aspects of the Salt Form". In Stahl PH, Warmth CG (eds.). Handbook of Pharmaceutical Salts: Properties, Selection, and Use. Weinheim: Wiley-VCH. pp. 92–94. ISBN 978-3-906390-58-1.
  7. ^ Vogel G. C.; Drago, R. S. (1996). "The ECW Model". Journal of Chemical Education. 73: 701–707. Bibcode:1996JChEd..73..701V. doi:10.1021/ed073p701.
  8. ^ Laurence, C. and Gal, J-F. Lewis Basicity and Affinity Scales, Data and Measurement, (Wiley 2010) pp 50-51 IBSN 978-0-470-74957-9
  9. ^ Cramer, R. E.; Bopp, T. T. (1977). "Graphical display of the enthalpies of adduct formation for Lewis acids and bases". Journal of Chemical Education. 54: 612–613. doi:10.1021/ed054p612. The plots shown in this paper used older parameters. Improved E&C parameters are listed in ECW model.
  10. ^ de Levie, Robert (1999). Aqueous Acid–Base Equilibria and Titrations. New York: Oxford University Press.
  11. ^ Jameson, Reginald F. (1978). "Assignment of the proton-association constants for 3-(3,4-dihydroxyphenyl)alanine (L-dopa)". Journal of the Chemical Society, Dalton Transactions (1): 43–45. doi:10.1039/DT9780000043.
  12. ^ Helfferich, Friedrich G. (1 January 1962). Ion Exchange. Courier Corporation. ISBN 9780486687841.
  13. ^ "Titration of Diprotic Acid". dwb.unl.edu. Archived from the original on 7 February 2016. Retrieved 24 January 2016.
  14. ^ a b Kotz, John C.; Treichel, Paul M.; Townsend, John; Treichel, David (24 January 2014). Chemistry & Chemical Reactivity. Cengage Learning. ISBN 9781305176461.
  15. ^ Lehninger, Albert L.; Nelson, David L.; Cox, Michael M. (1 January 2005). Lehninger Principles of Biochemistry. Macmillan. ISBN 9780716743392.
  16. ^ "The Top 10 Industrial Chemicals - For Dummies". dummies.com. Retrieved 5 February 2016.
  17. ^ "Sulfuric acid". essentialchemicalindustry.org. Retrieved 6 February 2016.
  18. ^ McMillin, John R.; Tracy, Gene A.; Harvill, William A.; Credle, William S., Jr. (8 December 1981), Method of and apparatus for making and dispensing a carbonated beverage utilizing propellant carbon dioxide gas for carbonating, retrieved 6 February 2016
  19. ^ Barrett, G. C.; Elmore, D. T. (June 2012). 8 - Biological roles of amino acids and peptides - University Publishing Online. doi:10.1017/CBO9781139163828. ISBN 9780521462921.
  20. ^ Graham, Timur (2006). "Acid Buffering". Acid Base Online Tutorial. University of Connecticut. Archived from the original on 13 February 2016. Retrieved 6 February 2016.
  21. ^ Voet, Judith G.; Voet, Donald (2004). Biochemistry. New York: J. Wiley & Sons. pp. 496–500. ISBN 978-0-471-19350-0.
  • Listing of strengths of common acids and bases
  • Zumdahl, Steven S. (1997). Chemistry (4th ed.). Boston: Houghton Mifflin. ISBN 9780669417944.
  • Pavia, D. L.; Lampman, G. M.; Kriz, G. S. (2004). Organic Chemistry Volume I. Mason, OH: Cengage Learning. ISBN 0759347271.

External links[edit]

  • Curtipot: Acid–Base equilibria diagrams, pH calculation and titration curves simulation and analysis – freeware