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Acid - Wikipedia. An acid is a molecule or ion capable of donating a hydron (proton or hydrogen ion H+), or, alternatively, capable of forming a covalent bond with an electron pair (a Lewis acid). In the special case of aqueous solutions, proton donors form the hydronium ion H3.
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O+ and are known as Arrhenius acids. The word acid is derived from the Latinacidus/ac. A lower p. H means a higher acidity, and thus a higher concentration of positive hydrogen ions in the solution. Chemicals or substances having the property of an acid are said to be acidic.
Common aqueous acids include hydrochloric acid (a solution of hydrogen chloride which 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.
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 which 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 Br. However, hydrogen chloride, acetic acid, and most other Br.
In modern terminology, an acid is implicitly a Br. Interestingly, although alcohols and amines can be Br. An Arrhenius acid is a substance that, when added to water, increases the concentration of H+ ions in the water.
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 HCl and acetic acid. An Arrhenius base, on the other hand, is a substance which increases the concentration of hydroxide (OH. This decreases the concentration of hydronium because the ions react to form H2. O molecules: H3. O+(aq) + OH.
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 1. Since p. H is defined as the negative logarithm of the concentration of hydronium ions, acidic solutions thus have a p. H of less than 7.
Br. 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 1. 92. 3 chemists Johannes Nicolaus Br. Consider the following reactions of acetic acid (CH3.
COOH), the organic acid that gives vinegar its characteristic taste: CH3. COOH + H2. O . In the second example CH3. COOH 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, CH3. COOH is both an Arrhenius and a Br. Hydrogen chloride (HCl) and ammonia combine under several different conditions to form ammonium chloride, NH4. Cl. 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: H3. O+(aq) + Cl. 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. A third, only marginally related concept was proposed in 1. 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. Many Lewis acids are not Br. Contrast how the following reactions are described in terms of acid- base chemistry: In the first reaction a fluoride ion, F. BF3 is a Lewis acid because it accepts the electron pair from fluoride. This reaction cannot be described in terms of Br.
The second reaction can be described using either theory. A proton is transferred from an unspecified Br. The species that gains the electron pair is the Lewis acid; for example, the oxygen atom in H3.
O+ gains a pair of electrons when one of the H—O bonds is broken and the electrons shared in the bond become localized on oxygen. Depending on the context, a Lewis acid may also be described as an oxidizer or an electrophile. Few, if any, of the acids discussed in the following are Lewis acids. Dissociation and equilibrium. Reactions of acids are often generalized in the form HA .
This reaction is referred to as protolysis. The protonated form (HA) of an acid is also sometimes referred to as the free acid. 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+ . 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 . 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+. Ka=. Because the range of possible values for Ka spans many orders of magnitude, a more manageable constant, p.
Ka is more frequently used, where p. Ka = . Stronger acids have a smaller p. Ka than weaker acids. Experimentally determined p. Ka at 2. 5 . 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, . Thus, for hydrogen chloride, the IUPAC name would be aqueous hydrogen chloride. 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. 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 (HCl.
O4), nitric acid (HNO3) and sulfuric acid (H2. SO4). In water each of these essentially ionizes 1. 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 Ka and a more negative p. Ka 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 1. Examples of superacids are fluoroantimonic acid, magic acid and perchloric acid.
Superacids can permanently protonate water to give ionic, crystalline hydronium . 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 p. H, which is an indication of the concentration of hydronium in the solution. The p. H of a simple solution of an acid compound in water is determined by the dilution of the compound and the compound's Ka.
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) + H2. O(l) . 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 (CH3.
COOH) and benzoic acid (C6. H5. COOH). 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 acid (two potential protons to donate) and triprotic acid (three potential protons to donate). Online streaming The Property in english with subtitles in 4K 16:9.
A diprotic acid (here symbolized by H2. A) can undergo one or two dissociations depending on the p. H. Each dissociation has its own dissociation constant, Ka. Ka. 2. H2. A(aq) + H2. O(l) . For example, sulfuric acid (H2. SO4) can donate one proton to form the bisulfate anion (HSO.
The large Ka. 1 for the first dissociation makes sulfuric a strong acid. In a similar manner, the weak unstable carbonic acid (H2. CO3) can lose one proton to form bicarbonate anion (HCO. Both Ka values are small, but Ka. Ka. 2 . Entreolivos movie online in english FULLHD 16:9 on this page. A triprotic acid (H3.
A) can undergo one, two, or three dissociations and has three dissociation constants, where Ka. Ka. 2 > Ka. 3. H3. A(aq) + H2. O(l) . All three protons can be successively lost to yield H2. PO. 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.