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An acid (...from the Latin acidus meaning sour) is traditionally considered any chemical compound that, when dissolved in water, gives a solution with a hydrogen ion activity greater than in pure water, i.e. a pH less than 7.0 in its standard state. That approximates the modern definition of Johannes Nicolaus Brønsted and Martin Lowry, who independently defined an acid as a compound which donates a hydrogen ion (H+) to another compound (called a base). Common examples include acetic acid (in vinegar) and sulfuric acid (used in car batteries). Acid/base systems are different from redox reactions in that there is no change in oxidation state. Acids can occur in solid, liquid or gaseous form, depending on the temperature. They can exist as pure substances or in solution. Chemicals or substances having the property of an acid are said to be acidic.

Contents

Definitions and concepts

Arrhenius acids

The Swedish chemist Svante Arrhenius attributed the properties of acidity to hydrogen in 1884. An Arrhenius acid is a substance that increases the concentration of the hydronium ion, H3O+, when dissolved in water. This definition stems from the equilibrium dissociation of water into hydronium and hydroxide (OH-) ions:

H2O(l) + H2O(l) is in equilibrium with H3O+(aq) + OH-(aq)

In pure water the majority of molecules exist as H2O, but a small number of molecules are constantly dissociating and re-associating. Pure water is neutral with respect to acidity or basicity because the concentration of hydroxide ions is always equal to the concentration of hydronium ions. An Arrhenius base is a molecule which increases the concentration of the hydroxide ion when dissolved in water. Note that chemists often write H+(aq) and refer to the hydrogen ion when describing acid-base reactions but the free hydrogen nucleus, a proton, does not exist alone in water, it exists as the hydronium ion, H3O+.

Brønsted acids

While the Arrhenius concept is useful for describing many reactions, it is also quite limited in its scope. In 1923 chemists Johannes Nicolaus Brønsted and Thomas Martin Lowry independently recognized that acid-base reactions involve the transfer of a proton. A Brønsted-Lowry acid (or simply Brønsted acid) is a species that donates a proton to a Brønsted-Lowry base. Brønsted-Lowry acid-base theory has several advantages over Arrhenius theory. Consider the following reactions of acetic acid (CH3COOH), the organic acid that gives vinegar its characteristic taste:

Acid-base.png

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, donating a proton to ammonia (NH3), but cannot be described using the Arrhenius definition of an acid because the reaction does not produce hydronium. Brønsted-Lowry theory can also be used to describe molecular compounds, whereas Arrhenius acids must be ionic compounds. 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' definition:

1.) H3O+(aq) + Cl-(aq) + NH3 → Cl-(aq) + NH4+(aq)
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. The next two reactions do not involve the formation of ions but can still be viewed as 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

A third concept was proposed by Gilbert N. Lewis which includes reactions with acid-base characteristics that do not involve a proton transfer. A Lewis acid is a species that accepts a pair of electrons from another species; in other words, it is an electron pair acceptor. Brønsted acid-base reactions are proton transfer reactions while Lewis acid-base reactions are electron pair transfers. All Brønsted acids are also Lewis acids, but not all Lewis acids are Brønsted acids. Contrast the following reactions which could be described in terms of acid-base chemistry.

LewisAcid.png

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.

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.

Dissociation and equilibrium

Reactions of acids are often generalized in the form HA is in equilibrium with H+ + A-, where HA represents the acid and A- is the conjugate base. Acid-base conjugate pairs differ by one proton, and can be interconverted by the addition or removal of a proton (protonation and deprotonation, respectively). Note that 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+ is in equilibrium with 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 concentration 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+.

K_a = \frac{[\mbox{H}^+] [\mbox{A}^-]}{[\mbox{HA}]}

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

In the classical naming system, acids are named according to their anions. That ionic suffix is dropped and replaced with a new suffix (and sometimes prefix), according to the table below. For example, HCl has chloride as its anion, so the -ide suffix makes it take the form hydrochloric acid. In the IUPAC naming system, "aqueous" is simply added to the name of the ionic compound. Thus, for hydrogen chloride, the IUPAC name would be aqueous hydrogen chloride. The prefix "hydro-" is added only if the acid is made up of just hydrogen and one other element.

Classical naming system:

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

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 higher 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. Superacids can permanently protonate water to give ionic, crystalline hydronium "salts". They can also quantitatively stabilize carbocations.

Polarity and the inductive effect

Polarity refers to the distribution of electrons in a bond, the region of space between two atomic nuclei where a pair of electrons is shared. When two atoms have roughly the same electronegativity (ability to attract electrons) the electrons are shared evenly and spend equal time on either end of the bond. When there is a significant difference in electronegativities of two bonded atoms, the electrons spend more time near the nucleus of the more electronegative element and an electrical dipole, or separation of charges, occurs, such that there is a partial negative charge localized on the electronegative element and a partial positive charge on the electropositive element. Hydrogen is an electropositive element and accumulates a slightly positive charge when it is bonded to an electronegative element such as oxygen or bromine. As the electron density on hydrogen decreases it is more easily abstracted, in other words, it is more acidic. Moving from left to right across a row on the periodic table elements become more electronegative (excluding the noble gases), and the strength of the binary acid formed by the element increases accordingly:

Formula Name pKa[1]
HF hydrofluoric acid 3.17
H2O water 15.7
NH3 ammonia 38
CH4 methane 48

The electronegative element need not be directly bonded to the acidic hydrogen to increase its acidity. An electronegative atom can pull electron density out of an acidic bond through the inductive effect. The electron-withdrawing ability diminishes quickly as the electronegative atom moves away from the acidic bond. The effect is illustrated by the following series of halogenated butanoic acids. Chlorine is more electronegative than bromine and therefore has a stronger effect. The hydrogen atom bonded to the oxygen is the acidic hydrogen. Butanoic acid is a carboxylic acid.

Structure Name pKa[2]
Butanoic.png butanoic acid ≈4.8
4chlorobutanoic.png 4-chlorobutanoic acid 4.5
3chlorobutanoic.png 3-chlorobutanoic acid ≈4.0
2bromobutanoic.png 2-bromobutanoic acid 2.93
2chlorobutanoic.png 2-chlorobutanoic acid 2.86

As the chlorine atom moves further away from the acidic O—H bond, its effect diminishes. When the chlorine atom is just one carbon removed from the carboxylic acid group the acidity of the compound increases significantly, compared to butanoic acid (a.k.a. butyric acid). However, when the chlorine atom is separated by several bonds the effect is much smaller. Bromine is much more electronegative than either carbon or hydrogen, but not as electronegative as chlorine, so the pKa of 2-bromobutanoic acid is slightly greater than the pKa of 2-chlorobutanoic acid.

Perchloric acid (HClO4) is an oxoacid and a strong acid.

The number of electronegative atoms adjacent an acidic bond also affects acid strength. Oxoacids have the general formula HOX where X can be any atom and may or may not share bonds to other atoms. Increasing the number of electronegative atoms or groups on atom X decreases the electron density in the acidic bond, making the loss of the proton easier. Perchloric acid is a very strong acid (pKa ≈ -8) and completely dissociates in water. Its chemical formula is HClO4 and it comprises a central chlorine atom with three chlorine-oxygen double bonds (Cl=O) and one chlorine-oxygen single bond (Cl—O). The singly bonded oxygen bears an extremely acidic hydrogen atom which is easily abstracted. In contrast, chloric acid (HClO3) is a weaker acid, though still quite strong (pKa = -1.0), while chlorous acid (HClO2, pKa = +2.0) and hypochlorous acid (HClO, pKa = +7.53) acids are weak acids.[3]

Carboxylic acids are organic acids that contain an acidic hydroxyl group and a carbonyl (C=O bond). Carboxylic acids can be reduced to the corresponding alcohol; the replacement of an electronegative oxygen atom with two electropositive hydrogens yields a product which is essentially non-acidic. The reduction of acetic acid to ethanol using LiAlH4 (lithium aluminium hydride or LAH) and ether is an example of such a reaction.

Reduction.png

The pKa for ethanol is 16, compared to 4.76 for acetic acid.[2][4]

Atomic radius and bond strength

Another factor that contributes to the ability of an acid to lose a proton is the strength of the bond between the acidic hydrogen and the atom that bears it. This, in turn, is dependent on the size of the atoms sharing the bond. For an acid HA, as the size of atom A increases, the strength of the bond decreases, meaning that it is more easily broken, and the strength of the acid increases. Bond strength is a measure of how much energy it takes to break a bond. In other words, it takes less energy to break the bond as atom A grows larger, and the proton is more easily removed by a base. This partially explains why hydrofluoric acid is considered a weak acid while the other hydrohalic acids (HCl, HBr, HI) are strong acids. Although fluorine is more electronegative than the other halogens, its atomic radius is also much smaller, so it shares a stronger bond with hydrogen. Moving down a column on the periodic table atoms become less electronegative but also significantly larger, and the size of the atom tends to dominate its acidity when sharing a bond to hydrogen. Hydrogen sulfide, H2S, is a stronger acid than water, even though oxygen is more electronegative than sulfur. Just as with the halogens, this is because sulfur is larger than oxygen and the H—S bond is more easily broken than the H—O bond.

Chemical characteristics

Monoprotic acids

Monoprotic 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) is in equilibrium with 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 carboxyl 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

Polyprotic 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 acid (two potential protons to donate) and triprotic 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) is in equilibrium with H3O+(aq) + HA(aq)       Ka1
HA(aq) + H2O(l) is in equilibrium with 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 (HSO4-), for which Ka1 is very large; then it can donate a second proton to form the sulfate anion (SO42-), 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 (HCO3-) and lose a second to form carbonate anion (CO32-). 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) is in equilibrium with H3O+(aq) + H2A(aq)        Ka1
H2A(aq) + H2O(l) is in equilibrium with H3O+(aq) + HA2−(aq)       Ka2
HA2−(aq) + H2O(l) is in equilibrium with 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 H2PO4-, then HPO42-, and finally PO43-, the orthophosphate ion, usually just called phosphate. An organic example of a triprotic acid is citric acid, which can successively lose three protons to finally form the citrate ion. Even though the positions of the protons on the original molecule may be equivalent, the successive Ka values will differ since it is energetically less favorable to lose a proton if the conjugate base is more negatively charged.

Neutralization

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 equilibria

In order to lose a proton, it is necessary that the pH of the system rise above the pKa of the protonated 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, or to protonate its conjugate base (the deprotonated form).

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

Applications of acids

There are numerous uses for acids. 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.

Strong acids, sulfuric acid in particular, are widely used in mineral processing. 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 used as catalysts; for example, sulfuric acid is used in very large quantities in the alkylation process to produce gasoline. Strong acids, such as sulfuric, phosphoric and hydrochloric acids also effect dehydration and condensation reactions.

Acids are used as additives to drinks and foods, as they alter their taste and serve as preservatives. Phosphoric acid, for example, is a component of cola drinks.

Biological occurrence

Basic structure of an amino acid.

Many biologically important molecules are acids. Nucleic acids, including DNA and RNA contain the genetic code that determines much 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 the 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 physiologic pH, typically around 7, free fatty 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 (-NH3+). The entire molecule has a net neutral charge and is a zwitterion.

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 is in equilibrium with H2CO3 is in equilibrium with H+ + HCO3-

It is the decrease in pH that signals the brain to breath 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

Mineral acids

Sulfonic acids

Carboxylic acids

Vinylogous carboxylic acids

Nucleic acids

See also

Chemistry
Environment

References

  1. ^ http://www2.lsdiv.harvard.edu/labs/evans/pdf/evans_pKa_table.pdf
  2. ^ a b http://www.scribd.com/doc/6792576/638478
  3. ^ pKa values for HClOn from C.E.Housecroft and A.G.Sharpe "Inorganic Chemistry" (Pearson, 2d edn 2005), p.171
  4. ^ http://research.chem.psu.edu/brpgroup/pKa_compilation.pdf
  • Listing of strengths of common acids and bases
  • Zumdahl, Chemistry, 4th Edition.
  • Ebbing, D.D., & Gammon, S. D. (2005). General chemistry (8th ed.). Boston, MA: Houghton Mifflin. ISBN 0618511776
  • Pavia, D.L., Lampman, G.M., & Kriz, G.S. (2004). Organic chemistry volume 1: Organic chemistry 351. Mason, OH: Cenage Learning. ISNB: 9780759342724

External links


Study guide

Up to date as of January 14, 2010
(Redirected to Acids and bases article)

From Wikiversity

Topics in acids and bases
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In chemistry, a compound could be an acid or a base. Acids and bases are very important in chemistry. They describe a large class of substances. Understanding their properties allows us to rationalize patterns that occur in chemistry.

Contents

Contrasting properties of acids and bases

Here is a list which contrasts their properties:

  • Acids turn blue litmus paper to red. Bases turn red litmus paper to blue.
  • Acids give off H+ ions in water; bases give off OH- ions in water
  • Acids generally taste sour due to the sour H+ ion; bases taste bitter due to the OH- ion; (NEVER TASTE CHEMICALS IN A LAB) but they may have other tastes depending on the other part of the molecule
  • Acids corrode active metals ; Bases denature protein
  • Acids have a pH less than 7 ; Bases have a pH greater than 7

Similar properties of acids and bases

Here is a list of properties that are common to both acids and bases

  • When you dissolve an acid or base in water, it makes ions. This makes the water conduct electricity better.
  • The stronger the acid (or base), the more ions are produced, so the conductivity of the solution increases.

Things unique to acids

Here are some properties of acids:

Acids react with some metals to generate hydrogen gas. Acids generally taste sour.

Examples of common things which are acidic

  • Milk
  • Saliva
  • Lemon juice
  • Apple juice
  • Vinegar
  • Water and soft drinks (Carbon dioxide reacts in water to form Carbonic acid)
  • Bleach
  • oranges
  • pineapples

Things unique to bases

Bases feel slippery and soapy because they hydrolyse the fatty acid ester of the skin (saponification reaction).

Examples of common things which are bases

  • Hand soap
  • Ammonia
  • Lye
  • Soda
  • Salt Water
  • lahat ng malapot

1911 encyclopedia

Up to date as of January 14, 2010

From LoveToKnow 1911

ACID (from the Lat. root ac-, sharp; acere, to be sour), the name loosely applied to any sour substance; in chemistry it has a more precise meaning,denoting a substance containing hydrogen which may be replaced by metals with the formation of salts. An acid may therefore be regarded as a salt of hydrogen. Of the general characters of acids we may here notice that they dissolve alkaline substances, certain metals, &c., neutralize alkalies and redden many blue and violet vegetable colouring matters.

The ancients probably possessed little knowledge indeed of acids. Vinegar (or impure acetic acid), which is produced when wine is allowed to stand, was known to both the Greeks and Romans, who considered it to be typical of acid substances; this is philologically illustrated by the words OEbs, acidus, sour, and duos, acetus, vinegar. Other acids became known during the alchemistic period; and the first attempt at a generalized conception of these substances was made by Paracelsus, who supposed them to contain a principle which conferred the properties of sourness and solubility. Somewhat similar views were promoted by Becher, who named the principle acidum primogenium, and held that it was composed of the Paracelsian elements " earth " and "water." At about the same time Boyle investigated several acids; he established their general reddening of litmus, their solvent power of metals and basic substances, and the production of neutral bodies, or salts, with alkalies. Theoretical conceptions were revived by Stahl, who held that acids were the fundamentals of all salts, and the erroneous idea that sulphuric acid was the principle of all acids.

The phlogistic theory of the processes of calcination and combustion necessitated the view that many acids, such as those produced by combustion, e.g. sulphurous, phosphoric, carbonic, &c., should be regarded as elementary substances. This principle more or less prevailed until it was overthrown by Lavoisier's doctrine that oxygen was the acid-producing element; Lavoisier being led to this conclusion by the almost general observation that acids were produced when non-metallic elements were burnt. The existence of acids not containing oxygen was, in itself, sufficient to overthrow this idea, but, although Berthollet had shown, in 1789, that sulphuretted hydrogen (or hydrosulphuric acid) contained no oxygen, Lavoisier's theory held its own until the researches of Davy, Gay-Lussac and Thenard on hydrochloric acid and chlorine, and of Gay-Lussac on hydrocyanic acid, established beyond all cavil that oxygen was not essential to acidic properties.

In the Lavoisierian nomenclature acids were regarded as binary oxygenated compounds, the associated water being relegated to the position of a mere solvent. Somewhat similar views were held by Berzelius, when developing his dualistic conception of the composition of substances. In later years Berzelius renounced the " oxygen acid " theory, but not before Davy, and, almost simultaneously, Dulong, had submitted that hydrogen and not oxygen was the acidifying principle. Opposition to the " hydrogen-acid " theory centred mainly about the hypothetical radicals which it postulated; moreover, the electrochemical theory of Berzelius exerted a stultifying influence on the correct views of Davy and Dulong. In Berzelius' system + potassium sulphate is to be regarded as K 2 0.S0 3; electrolysis should simply effect the disruption of the positive and negative components, potash passing with the current, and sulphuric acid against the current. Experiment showed, however, that instead of only potash appearing at the negative electrode, hydrogen is also liberated; this is inexplicable by Berzelius's theory, but readily explained by the " hydrogen-acid " theory. By this theory potassium is liberated at the negative electrode and combines immediately with water to form potash and hydrogen.

Further and stronger support was given when J. Liebig promoted his doctrine of polybasic acids. Dalton's idea that elements preferentially combined in equiatomic proportions had as an immediate inference that metallic oxides contained one atom of the metal to one atom of oxygen, and a simple expansion of this conception was that one atom of oxide combined with one atom of acid to form one atom of a neutral salt. This view, which was specially supported by Gay-Lussac and Leopold Gmelin and accepted by Berzelius, necessitated that all acids were monobasic. The untenability of this theory was proved by Thomas Graham's investigation of the phosphoric acids; for he then showed that the ortho- (ordinary), pyroand metaphosphoric acids contained respectively 3, 2 and I molecules of " basic water " (which were replaceable by metallic oxides) and one molecule of phosphoric oxide, P2 05. Graham's work was developed by Liebig, who called into service many organic acids - citric, tartaric, cyanuric, comenic and meconic - and showed that these resembled phosphoric acid; and he established as the criterion of polybasicity the existence of compound salts with different metallic oxides. In formulating these facts Liebig at first retained the dualistic conception of the structure of acids; but he shortly afterwards perceived that this view lacked generality since the halogen acids, which contained no oxygen but yet formed salts exactly similar in properties to those containing oxygen, could not be so regarded. This and other reasons led to his rejection of the dualistic hypothesis and the adoption, on the ground of probability, and much more from convenience, of the tenet that " acids are particular compounds of hydrogen, in which the latter can be replaced by metals "; while, on the constitution of salts, he held that " neutral salts are those compounds of the same class in which the hydrogen is replaced by its equivalent in metal. The substances which we at present term anhydrous acids (acid oxides) only become, for the most part, capable of forming salts with metallic oxides after the addition of water, or they are compounds which decompose these oxides at somewhat high temperatures." The hydrogen theory and the doctrine of polybasicity as enunciated by Liebig is the fundamental characteristic of the modern theory. A polybasic acid contains more than one atom of hydrogen which is replaceable by metals; moreover, in such an acid the replacement may be entire with the formation of normal salts, partial with the formation of acid salts, or by two or more different metals with the formation of compound salts (see Salts). These facts may be illustrated with the aid of orthophosphoric acid, which is tribasic: - Acid. Normal salt. Acid salts.

H3P04. Ag3P04. Na2HP04; NaH2P04.

Phosphoric Silver phosphate. Acid sodium acid. phosphates. Compound salts.

Mg(NH4)P04; Na(NH4)HP04.

Magnesium ammonium Microcosmic phosphate; salt.

Reference should be made to the articles Chemical Action, Thermochemistry and Solutions, for the theory of the strength or avidity of acids.

Organic Acids. - Organic acids are characterized by the presence of the monovalent group - CO. OH, termed the carboxyl group, in which the hydrogen atom is replaceable by metals with the formation of salts, and by alkyl radicals with the formation of esters. The basicity of an organic acid, as above defined, is determined by the number of carboxyl groups present. Oxy-acids are carboxylic acids which also contain a hydroxyl group; similarly we may have aldehyde-acids, ketone-acids, &c. Since the more important acids are treated under their own headings, or under substances closely allied to them, we shall here confine ourselves to general relations.

Classification

It is convenient to distinguish between aliphatic and aromatic acids; the first named being derived from open-chain hydrocarbons, the second from ringed hydrocarbon nuclei. Aliphatic monobasic acids are further divided according to the nature of the parent hydrocarbon. Methane and its homologues give origin to the " paraffin " or " fatty series " of the general formula C,H 2 ,,+ 1 000H, ethylene gives origin to the acrylic acid series, C n H 27 ,- 1 000H, and so on. Dibasic acids of the paraffin series of hydrocarbons have the general formula C n H 2 (000H) 2 "; malonic and succinic acids are important members. The isomerism which occurs as soon as the molecule contains a few carbon atoms renders any classification based on empirical molecular formulae somewhat ineffective; on the other hand, a scheme based on molecular structure would involve more detail than it is here possible to give. For further information, the reader is referred to any standard work on organic chemistry. A list of the acids present in fats and oils is given in the article Oils.

Syntheses of Organic Acids

The simplest syntheses are undoubtedly those in which a carboxyl group is obtained directly from the oxides of carbon, carbon dioxide and carbon monoxide. The simplest of all include: (z) the synthesis of sodium oxalate by passing carbon dioxide over metallic sodium heated to 3500 - 360'; (2) the synthesis of potassium formate from moist carbon dioxide and potassium, potassium carbonate being obtained simultaneously; (3) the synthesis of potassium acetate and propionate from carbon dioxide and sodium methide and sodium ethide; (4) the synthesis of aromatic acids by the interaction of carbon dioxide, sodium and a bromine substitution derivative; and (5) the synthesis of aromatic oxy-acids by the interaction of carbon dioxide and sodium phenolates (see Salicylic Acid). Carbon monoxide takes part in the syntheses of sodium formate from sodium hydrate, or soda lime (at 200 0 -2 20 0), and of sodium acetate and propionate from sodium methylate and sodium ethylate at 160 0 -200°. Other reactions which introduce carboxyl groups into aromatic groups are: the action of carbonyl chloride on aromatic hydrocarbons in the presence of aluminium chloride, acid-chlorides being formed which are readily decomposed by water to give the acid; the action of urea chloride Cl�CO�NH 2, cyanuric acid (CONH) 3, nascent cyanic acid, or carbanile on hydrocarbons in the presence of aluminium chloride, acid-amides being obtained which are readily decomposed to give the acid. An important nucleus-synthetic reaction is the saponification of nitriles, which may be obtained by the interaction of potassium cyanide with a halogen substitution derivative or a sulphonic acid.

Acids frequently result as oxidation products, being almost invariably formed in all cases of energetic oxidation. There are certain reactions, however, in which oxidation can be successfully applied to the synthesis of acids. Thus primary alcohols and aldehydes, both of the aliphatic and aromatic series, readily yield on oxidation acids containing the same number of carbon atoms. These reactions may be shown thus: - � R�CH 2 OH -> R. CHO - R�CO.OH.

In the case of aromatic aldehydes, acids are also obtained by means of " Cannizzaro's reaction " (see Benzaldehyde). An important oxidation synthesis of aromatic acids is from hydrocarbons with aliphatic side chains; thus toluene, or methylbenzene, yields benzoic acid, the xylenes, or dimethyl-benzene, yield methyl-benzoic acids and phthalic acids. Ketones, secondary alcohols and tertiary alcohols yield a mixture of acids on oxidation. We may also notice the disruption of unsaturated acids at the double linkage into a mixture of two acids, when fused with potash.

In the preceding instances the carboxyl group has been synthesized or introduced into a molecule; we have now to consider syntheses from substances already containing carboxyl groups. Of foremost importance are the reactions termed the malonic acid and the aceto-acetic ester syntheses; these are discussed under their own headings. The electrosyntheses call for mention here. It is apparent that metallic salts of organic acids would, in aqueous solution, be ionized, the positive ion being the metal, and the negative ion the acid residue. Esters, however, are not ionized. It is therefore apparent that a mixed salt and ester, for example KO 2 C�CH 2 �CH 2 �CO 2 C 2 H 5, would give only two ions, viz. potassium and the rest of the molecule. If a solution of potassium acetate be electrolysed the products are ethane, carbon dioxide, potash and hydrogen; in a similar manner, normal potassium succinate gives ethylene, carbon dioxide, potash and hydrogen; these reactions may be represented: CH 3 �CO 2;K CH 3 CO 2 K' CH 2 �CO 2 1K CH 2 CO 2 K' --> I + + I I -i iI + CH 3 �CO 21 K CH 3 CO 2 K' CH 2 �CO 2 iK CH 2 CO 2 K' By electrolysing a solution of potassium ethyl succinate, KO 2 C�(CH 2) 2 CO 2 C 2 H 5, the KO 2 C� groups are split off and the two residues �(CH 2) 2 CO 2 C 2 H 5 combine to form the ester (CH2)4(C02C2H5)2. In the same way, by electrolysing a mixture of a metallic salt and an ester, other nuclei may be condensed; thus potassium acetate and potassium ethyl succinate yield CH 3 * CH2 � CH2 � C02 C2H5.

Reactions

Organic acids yield metallic salts with bases, and ethereal salts or esters, R. CO. OR', with alcohols. Phosphorus chlorides give acid chlorides, R�CO�C1, the hydroxyl group being replaced by chlorine, and acid anhydrides, (R�CO) 2 0, a molecule of water being split off between two carboxyl groups. The ammonium salts when heated lose one molecule of water and are converted into acid-amides, R. CO�NH2, which by further dehydration yield nitriles, R�CN. The calcium salts distilled with calcium formate yield aldehydes r distilled with soda-lime, ketones result.


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Wiktionary

Up to date as of January 15, 2010

Definition from Wiktionary, a free dictionary

See also acid
See also ACID

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Proper noun

Acid

  1. a musical genre influenced by the typical electronic sounds of the Roland TR-808 and TR-909 synthesizers.

Noun

Acid n. (genitive Acids, no plural)

  1. (slang) LSD

Simple English

The article about certain chemical substances is at Acid

ACID is a concept used with Databases. It is an abbreviation. It stands for

  • Atomicity - something is either done, or not. The state of doing it is not visible outside the database.
  • Consistency - The database is in a legal state at all times. When a transaction occurs, it can not break the rules. These rules are about integrity, what is allowed and what is not allowed in certain locations of the database.
  • Isolation - There can be more than one transaction occurring at the same time. A certain transaction will not see changes made by other transactions.
  • Durability - When a transaction is done, it will be committed. After it is committed, it can no longer be undone.








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