arrhenius theory of electrolytic dissociation
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An acid-base reaction is a chemical reaction that occurs between an acid and a base. Several concepts that provide alternative definitions for the reaction mechanisms involved and their application in solving related problems exist. Despite several differences in definitions, their importance becomes apparent as different methods of analysis when applied to acid-base reactions for gaseous or liquid species, or when acid or base character may be somewhat less apparent. The first of these scientific concepts of acids and bases was provided by the FrenchchemistAntoine Lavoisier, circa 1776.
Historic acid-base theories
Lavoisier's oxygen theory of acids
The first scientific concept of acids and bases was provided by Antoine Lavoisier circa 1776. Since Lavoisier's knowledge of strong acids was mainly restricted to oxoacids, such as link=nitric acid|HNO|3| (nitric acid) and link=sulfuric acid|H|2|SO|4 (sulfuric acid), which tend to contain central atoms in high oxidation states surrounded by oxygen, and since he was not aware of the true composition of the hydrohalic acids (HF, HCl, HBr, and HI), he defined acids in terms of their containing oxygen, which in fact he named from Greek words meaning "acid-former" (from theGreekÎ¿Î¾Ï…Ï‚ (oxys) meaning "acid" or "sharp" and Î³ÎµÎ¹Î½Î¿Î¼Î±Î¹ (geinomai) meaning "engender"). The Lavoisier definition was held as absolute truth for over 30 years, until the 1810 article and subsequent lectures by Sir Humphry Davy in which he proved the lack of oxygen in H2S, H2Te, and the hydrohalic acids. However, Davy failed to develop a new theory, concluding that "acidity does not depend upon any particular elementary substance, but upon peculiar arrangement of various substances". One notable modification of oxygen theory was provided by Berzelius, who stated that acids are oxides of nonmetals while bases are oxides of metals.
Liebig's hydrogen theory of acids
This definition was proposed by Justus von Liebig circa 1838, based on his extensive works on the chemical composition of organic acids. This finished the doctrinal shift from oxygen-based acids to hydrogen-based acids, started by Davy. According to Liebig, an acid is a hydrogen-containing substance in which the hydrogen could be replaced by a metal. Liebig's definition, while completely empirical, remained in use for almost 50 years until the adoption of the Arrhenius definition.
Common acid-base theories
The Arrhenius definition of acid-base reactions is a development of the hydrogen theory of acids, devised by Svante Arrhenius, which was used to provide a modern definition of acids and bases that followed from his work with Friedrich Wilhelm Ostwald in establishing the presence of ions in aqueous solution in 1884, and led to Arrhenius receiving the Nobel Prize in Chemistry in 1903 for "recognition of the extraordinary services... rendered to the advancement of chemistry by his electrolytic theory of dissociation".
As defined by Arrhenius, acid-base reactions are characterized by Arrhenius acids, which dissociate in aqueous solution to form hydrogen ions (H|+), and Arrhenius bases, which form hydroxide (OH|âˆ’) ions. More recent IUPAC recommendations now suggest the newer term "hydronium" be used in favor of the older accepted term "oxonium" to illustrate reaction mechanisms such as those defined in the BrÃ¸nsted-Lowry and solvent system definitions more clearly, with the Arrhenius definition serving as a simple general outline of acid-base character. The Arrhenius definition can be summarised as "Arrhenius acids form hydrogen ions in aqueous solution with Arrhenius bases forming hydroxide ions."
The universal aqueous acid-base definition of the Arrhenius concept is described as the formation of water from hydrogen and hydroxide ions, or hydrogen ions and hydroxide ions from the dissociation of an acid and base in aqueous solution:
- H|+ (aq) + OH|âˆ’ (aq) H|2|O
(In modern times, the use of H|+ is regarded as a shorthand for H|3|O|+, since it is now known that the bare proton H|+ does not exist as a free species in solution.)
This leads to the definition that in Arrhenius acid-base reactions, a salt and water is formed from the reaction between an acid and a base. In other words, this is a neutralization reaction.
- acid+ + baseâˆ’â†’ salt + water
The positive ion from a base forms a salt with the negative ion from an acid. For example, two moles of the base sodium hydroxide (NaOH) can combine with one mole of sulfuric acid (H|2|SO|4) to form two moles of water and one mole of sodium sulfate.
- 2 NaOH + H|2|SO|4 â†’ 2 H|2|O + Na|2|SO|4
The Arrhenius definitions of acidity and alkalinity are restricted to aqueous solutions, and refer to the concentration of the solvent ions. Under this definition, pure H|2|SO|4 or HCl dissolved in toluene are not acidic, and molten KOH and solutions of sodium amide in liquid ammonia are not alkaline.
Solvent system definition
One of the limitations of Arrhenius definition was its reliance on water solutions. E. C. Franklin studied the acid-base reactions in liquid ammonia in 1905 and pointed out the similarities to water-based Arrhenius
An aqueous solution is a solution in which the solvent is water. It is usually shown in chemical equations by appending (aq) to the relevant formula. The word aqueous means pertaining to, related to, similar to, or dissolved in water. As water is an excellent solvent and is also naturally abundant, it is an ubiquitous solvent in chemistry.
Substances which are hydrophobic('water fearing') often do not dissolve well in water whereas those thathydrophilic('water-loving') do. An example of a hydrophilic substance would besodium chloride (ordinary table salt). Acids and bases are aqueous solutions, as part of their Arrhenius definitions.
The ability of a substance to dissolve in water is determined by whether the substance can match or exceed the strong attractive forces that water molecules generate between themselves. If the substance lacks the ability to dissolve in water the molecules form a precipitate.
Aqueous solutions that conduct electric current efficiently contain strong electrolytes, while ones that conduct poorly are considered to have weak electrolytes. Those strong electrolytes are substances that are completely ionized in water, whereas the weak electrolytes exhibit only a small degree of ionization in water.
Nonelectrolytes are substances that dissolve in water, but which maintain their molecular integrity (do not dissociate into ions). Examples include sugar, urea, glycerol, and methylsulfonylmethane (MSM).
When performing calculations regarding the reacting of one or more aqueous solutions, one generally must know the concentration, or molarity, of the aqueous solutions. Solution concentration is given in terms of the form of the solute prior to it dissolving.
acids and bases two related classes of chemicals; the members of each class have a number of common properties when dissolved in a solvent, usually water. Properties Acids in water solutions exhibit the following common properties: they taste sour; turn litmus paper red; and react with certain metals, such as zinc, to yield hydrogen gas. Bases in water solutions exhibit these common properties: they taste bitter; turn litmus paper blue; and feel slippery. When a water solution of acid is mixed with a water solution of base, water and a salt are formed; this process, called neutralization , is complete only if the resulting solution has neither acidic nor basic properties. Classification Acids and bases can be classified as organic or inorganic. Some of the more common organic acids are: citric acid , carbonic acid , hydrogen cyanide , salicylic acid, lactic acid , and tartaric acid . Some examples of organic bases are: pyridine and ethylamine. Some of the common inorganic acids are: hydrogen sulfide , phosphoric acid , hydrogen chloride , and sulfuric acid . Some common inorganic bases are: sodium hydroxide , sodium carbonate , sodium bicarbonate , calcium hydroxide , and calcium carbonate . Acids, such as hydrochloric acid, and bases, such as potassium hydroxide, that have a great tendency to dissociate in water are completely ionized in solution; they are called strong acids or strong bases. Acids, such as acetic acid, and bases, such as ammonia, that are reluctant to dissociate in water are only partially ionized in solution; they are called weak acids or weak bases. Strong acids in solution produce a high concentration of hydrogen ions, and strong bases in solution produce a high concentration of hydroxide ions and a correspondingly low concentration of hydrogen ions. The hydrogen ion concentration is often expressed in terms of its negative logarithm, or p H (see separate article). Strong acids and strong bases make very good electrolytes (see electrolysis ), i.e., their solutions readily conduct electricity. Weak acids and weak bases make poor electrolytes. See buffer ; catalyst ; indicators, acid-base ; titration . Acid-Base Theories There are three theories that identify a singular characteristic which defines an acid and a base: the Arrhenius theory, for which the Swedish chemist Svante Arrhenius was awarded the 1903 Nobel Prize in chemistry; the BrÃ¶nsted-Lowry, or proton donor, theory, advanced in 1923; and the Lewis, or electron-pair, theory, which was also presented in 1923. Each of the three theories has its own advantages and disadvantages; each is useful under certain conditions. The Arrhenius Theory When an acid or base dissolves in water, a certain percentage of the acid or base particles will break up, or dissociate (see dissociation ), into oppositely charged ions. The Arrhenius theory defines an acid as a compound that can dissociate in water to yield hydrogen ions, H + , and a base as a compound that can dissociate in water to yield hydroxide ions, OH - Â . For example, hydrochloric acid, HCl, dissociates in water to yield the required hydrogen ions, H + , and also chloride ions, Cl - Â . The base sodium hydroxide, NaOH, dissociates in water to yield the required hydroxide ions, OH - , and also sodium ions, Na + . The BrÃ¶nsted-Lowry Theory Some substances act as acids or bases when they are dissolved in solvents other than water, such as liquid ammonia. The BrÃ¶nsted-Lowry theory, named for the Danish chemist Johannes BrÃ¶nsted and the British chemist Thomas Lowry, provides a more general definition of acids and bases that can be used to deal both with solutions that contain no water and solutions that contain water. It defines an acid as a proton donor and a base as a proton acceptor. In the BrÃ¶nsted-Lowry theory, water, H 2 O, can be considered an acid or a base since it can lose a proton to form a hydroxide ion, OH - , or accept a proton to form a hydronium ion, H 3 O + (see amphoterism ). When an acid loses a proton, the remaining species can be a proton acceptor and is called the conjugate base of the acid. Similarly when a base accepts a proton, the resulting species can be a proton donor and is called the conjugate acid of that base. For example, when a water molecule loses a proton to form a hydroxide ion, the hydroxide ion can be considered the conjugate base of the acid, water. When a water molecule accepts a proton to form a hydronium ion, the hydronium ion can be considered the conjugate acid of the base, water. The Lewis Theory Another theory that provides a very broad definition of acids and bases has been put forth by the American chemist Gilbert Lewis. The Lewis theory defines an acid as a compound that can accept a pair of electrons and a base as a compound that can donate a pair of electrons. Boron trifluoride, BF 3 , can be considered a Lewis acid and ethyl alcohol can be considered a Lewis base.
Acids and bases have been known by their properties since the early days of experimental chemistry. The word "acid" comes from the Latin acidus, meaning "sour" or "tart," since water solutions of acids have a sour or tart taste. Lemons, grapefruit, and limes taste sour because they contain citric acid and ascorbic acid (vitamin C). Another common acid is vinegar, which is the sour liquid produced when apple cider, grape juice, or other plant juices ferment beyond the formation of alcohol. Vinegar is a 5 percent water solution of acetic acid. Besides having a sour taste, acids react with active metals to give hydrogen, they change the colors of indicators (for example, litmus turns from blue to red), and they neutralize bases. Bases change the colors of indicators (litmus turns from red to blue) and they neutralize acids. Hence, bases are considered the chemical opposite of acids. Most common acid-base reactions take place in water solutions (commonly referred to as aqueous solutions ). One of the earliest definitions of acids, advanced by the Swedish physicist and chemist Svante Arrhenius in 1887, stated that acid ionizes in aqueous solution to produce hydrogen ions (which are protons), H+, and anions ; and a base ionizes in aqueous solution to produce hydroxide ions (OHâˆ’) and cations. Later studies of aqueous solutions provided evidence of a small, positively charged hydrogen ion combining with a water molecule to form a hydrated proton, H+(H2O) or H3O+, which is called the hydronium ion. Often, the hydronium ion or hydrated proton is represented as H+ (aq ). Hydrogen chloride (HCl), a gas, is an acid because it dissolves in water to yield hydrogen ions and chloride ions. This water solution of HCl is referred to as hydrochloric acid. Â Â Â Â Â (1) A typical base, according to the Arrhenius definition, is sodium hydroxide (NaOH). It dissolves in water to give sodium ions and hydroxide ions. Â Â Â Â Â (2) In the reaction of an acid with a base in aqueous solution, the hydrogen ions of the acid react with the hydroxide ions of the base to give water. The second product is a salt, which is composed of the positive metal ion from the base and the negative ion from the acid. For example, HCl (aq ) + KOH (aq ) â†’ H2O (l ) + KCl (aq ) Â Â Â Â Â (3) Since HCl (aq ) and KOH (aq ) are fully ionized in solution, the preceding equation can be written as H+ (aq ) + Clâˆ’ (aq ) K+ (aq ) + OHâˆ’ (aq ) â†’ H2O (l ) K+ (aq ) + Clâˆ’ (aq ) Â Â Â Â Â (4) Ions common to both sides can be canceled to yield H+ (aq ) + OHâˆ’ (aq ) â†’ H2O (l ) Â Â Â Â Â (5) This is referred to as the net ionic equation for the neutralization reaction. If H3O+ is substituted for H+ (aq ), the neutralization equation becomes H3O+ (aq ) + OHâˆ’ (aq ) â†’ 2 H2O (l ) Â Â Â Â Â (6) The strength of an acid or base is determined by the extent of its ionization in aqueous solution. Strong acids, such as hydrochloric acid, are 100 percent ionized in aqueous solution, whereas weak acids, such as acetic acid, are less than 5 percent ionized. Experimentally, the extent of ionization is determined by measuring the electrical conductance of solutions. Strong acids and bases are strong electrolytes, and weak acids and bases are weak electrolytes. Table 1 lists some common acids and bases and indicates whether they are strong or weak. For weak acids and bases, partial ionization is a dynamic equilibrium between unionized molecules and its ion, as indicated by the double arrow in equation (7). For example, acetic acid is only partially ionized in aqueous solution CH3COOH (aq ) â‡† H+ (aq ) + CH3COOâˆ’ (aq ) Â Â Â Â Â (7) In acetic acid, hydrogen ions and acetate ions recombine to form acetic acid molecules. The double arrow signifies that at any given instant, less than 5 percent of acetic acid molecules dissociate into hydrogen ions and acetate ions, while the hydrogen ions and acetate ions recombine to form acetic acid molecules. Ammonia (NH3) is a weak base, and although it does not have OHâˆ’ ions in its formula, it produces the ion on reaction with water. NH3 (aq ) + H2O (l ) â‡† NH4+ (aq ) + OHâˆ’ (aq ) Â Â Â Â Â (8) A major problem with Arrhenius's acid-base theory is that some substances, like ammonia, produce basic solutions and react with acids, but do not contain hydroxide ions. In 1923 Johannes BrÃ¸nsted, a Danish chemist, and Thomas Lowry, an English chemist, independently proposed a new way to define acids and bases. An acid donates hydrogen ions (also called a proton donor); a base accepts hydrogen ions (also called a proton acceptor). These definitions not only explain all the acids and bases covered by Arrhenius's theory, they also explain the basicity of ammonia and ions such as carbonate, CO32âˆ’, and phosphate, PO43. The BrÃ¸nsted-Lowry theory includes water as a reactant and considers its acidity or basicity in the reaction. In the partial ionization of acetic acid, water is a base because it accepts the hydrogen ion to form hydronium ion. Â Â Â Â Â (9) A meticulous experimenter, Thomas Lowry is best known for his conceptualization of acidâ€“base chemistry. Studies of nitrogenous compounds led Lowry to question fundamental aspects of the role of hydrogen during acidâ€“base reactions. Three months before BrÃ¸nsted published his theory, Lowry released his own similar thoughts on proton acceptors and donors in print. â€”Valerie Borek In the reaction, a new acid and a new base are formed, which are called the conjugate acid and conjugate base, respectively. The hydronium ion, H3O+, is the conjugate acid of the base, H2O, and the acetate ion, CH3COOâˆ’, is the conjugate base of acetic acid, CH3COOH. A pair of molecules or ions related to one another by the gain or loss of a single hydrogen ion is called a conjugate acid-base pair. In the reaction of ammonia, water is an acid because it donates a hydrogen ion to ammonia. Â Â Â Â Â (10) This ability of water to donate or accept hydrogen ions, depending on whether it reacts with a base or an acid, is referred to as "amphiprotic." The conjugate acid-base pairs in this reaction are NH3/NH4+ and H2O/OHâˆ’. The BrÃ¸nsted-Lowry definitions also explain why carbonate salts such as sodium carbonate (washing soda) dissolve in water to give basic solutions. Carbonate ion removes a hydrogen ion from a water molecule, which leaves behind a hydroxide ion: Â Â Â Â Â (11) In the preceding reaction, water and hydroxide ion are a conjugate acid-base pair, whereas carbonate ion and bicarbonate ion are a conjugate base-acid pair. Every BrÃ¸nsted-Lowry acid has a conjugate base, and every BrÃ¸nsted-Lowry base has a conjugate acid. Familiarity with conjugate acid-base pairs is important to understanding the relative strengths of acids and bases. Table 2 lists some conjugate acid-base pairs and their relative strengths. Strong acids have weak conjugate bases, and weak acids have strong conjugate bases. Several common acids have more than one ionizable hydrogen ion (Table1). Each successive hydrogen ion in these polyprotic acids ionizes less readily. For example, sulfuric acid is a strong acid because of the complete ionization of the first hydrogen ion. H2SO4 (aq ) + H2O (l ) â†’ H3O+ (aq ) + HSO4âˆ’ (aq ) Â Â Â Â Â (12) The HSO4âˆ’ also acts as an acid, but it is not 100 percent ionized, so HSO4âˆ’ is an acid of moderate strength. For example, sodium hydrogen sulfate is used to increase the acidity of swimming pools, whereas sodium carbonate is used to increase the basicity of swimming pools. HSO4âˆ’ (aq ) + H2O (l ) â‡† H3O+ (aq ) + SO42âˆ’ (aq ) Â Â Â Â Â (13) Phosphoric acid has three ionizable hydrogen ions. Each stepwise ionization of phosphoric acid occurs to a lesser extent than the one before it. Phosphoric acid is stronger than acetic acid because the first step ionizes to a greater extent than acetic acid. H3PO4 (aq ) + H2O (l ) â‡† H3O+ (aq ) + H2PO4âˆ’ (aq ) Â Â Â Â Â (14) However, H2PO4âˆ’ is a weaker acid than acetic acid because the second ionization is much smaller (by a factor of 105) than the first step. H2PO4âˆ’ (aq ) + H2O (l ) â‡† H3O+ (aq ) + HPO42âˆ’ (aq ) Â Â Â Â
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Answers:arrhenius base: substance that dissociate OH in water arrhenius acid: substance that dissociate H+ in water example: NaOH, HCl
Answers:The Lewis theory is applicable to the largest number of compounds. This is because of the following. * Arrhenius acids are equivalent to Br nsted-Lowry acids in aqueous solution. However, a Br nsted-Lowry acid can also be a compound which donates protons in non-aqueous solvents. * Arrhenius bases are those compounds which dissociate to give OH- in aqueous solution. Br nsted-Lowry bases, on the other hand, include Arrhenius bases because they accept a proton from water, converting it into OH-, but can also accept protons from other solvents, or produce OH- without dissociation (as is the case with bases such as NH3.) Therefore, * The Br nsted-Lowry definition encompasses every Arrhenius acid/base, and more. Then, consider the Lewis definition. * A Br nsted-Lowry acid is a compound which donates protons (H+). A Lewis acid is a compound which can accept an electron pair, forming a bond. Note that every Br nsted-Lowry acid is also a Lewis acid, since H+ has the capacity to accept an electron pair. However, compounds such as BF3 are Lewis acids but not Br nsted-Lowry acids - it accepts an electron pair, but not by virtue of hydrogen ion. * A Lewis base, similarly, is a compound which donates an electron pair. A Br nsted-Lowry base combines with H+, forming a bond to the hydrogen atom by donating its electron pair. Therefore, every Br nsted-Lowry base is also a Lewis base. Therefore, the number of compounds covered by each theory increases as Arrhenius < Br nsted-Lowry < Lewis.
Answers:Theory: Arrhenius According to Arrhenius, if a molecular substance ionizes (forms ions) in water solution, and the only positive ion turns out to be a hydrogen ion (H+), then the substance is an acid. Example: HCl(aq) --> H+(aq) + Cl-(aq) In the above example, a molecule of hydrogen chloride, when dissolved in water, breaks up to produce a positive hydrogen ion and a negative chloride ion. Of course, I am leaving the role of water from this equation. Hydrogen chloride molecules and water molecules are highly polar. That means, that they both have oppositely charged ends that attract each other. When the positive end of HCl (the hydrogen end) attracts the negative end of H2O (the oxygen end), the oxygen rips the hydrogen atom (actually just the proton) off the hydrogen chloride molecule leaving the hydrogen's electron behind it. This extra electron gives the chlorine atom its negative charge transforming it into a chloride ion, while the water "adopts" the hydrogen (proton) that begins to share a pair of oxygen's electrons forming a covalent bond with it. Since a proton carries a positive charge, the entire resulting molecule becomes a positive polyatomic ion called the "hydronium ion". The complete equation for this reaction that includes water is HCl(aq) + H2O(l) --> H3O+(aq) + Cl-(aq) H+ and H3O+ can be used interchangeably for the hydrogen ion, because hydronium ion (H3O+) is just a proton attached to a water molecule. A base, on the other hand, is a substance that either dissociates or ionizes in water to release hydroxide ions (OH-) as the only negative ions in water solution. Most bases are already ionic, since they usually contain a metallic (positive) ion. An ionic substance does not need to ionize (form ions) in water solution. It only dissociates. Polar water molecules are attracted to both the positve and the negative ions in the crystal, and surround them, separating them from one another, making the crystal dissolve and become part of the solution. Example: NaOH(s) --> Na+(aq) + OH-(aq) Note that the only negative ion that is released in water is OH- (hydroxide ion. This makes sodium hydroxide a base. Some bases (the weak kind) start out as molecules. When those molecules interact with water, some of them capture a proton (H+) from a water molecule and become positive polyatomic ions, while the rest of the water molecules without a proton become a hydroxide ions (OH-). Example: NH3(aq) + H2O(l) --> NH4+(aq) + OH-(aq) In the above equation, an ammonia molecule takes a proton from a water molecule and becomes a positive ammonium ion, while the remaining part of the water molecule becomes a negative hydroxide ion. This makes ammonia a base. Theory: Bronsted-Lowry According to Bronsted-Lowry, an acid is a particle that donates a proton to another particle. The "particle" may be a molecule, an atom, or an ion. Example: HCl(aq) + H2O(l) --> H3O+(aq) + Cl-(aq) In the above equation, hydrogen chloride molecule "donates" a proton (H+) to the water molecule. Therefore, HCl is defined as an "acid". However, donating something means donating it to something else that accepts the donation. That something is called a base. In the above example, water accepts the proton, and in so doing, it acts as a "base". Another example: NH3(aq) + H2O(l) --> NH4+(aq) + OH-(aq) In the equation above, an ammonia molecule (NH3) takes (accepts) a proton from a water molecule. This makes the ammonia molecule a base, while the water molecule acts like an acid since it supplies (donates) the proton to the ammonia molecule. Note, that in the above examples, water acts sometimes as an acid, and sometimes as a base, depending what it is that water is interacting with. Theory: Lewis According to Lewis, an acid is any particle that accepts an electron pair. Example: H2O(l) + H+(aq) --> H3O+(aq) In the above example, a water molecule, accepts a proton. A proton does not have any electrons of its own. A water molecule is made up of a central oxygen atom that is surrounded by four pairs of electrons. Two of these pairs are shared with two hydrogen atoms respectively, leaving two unshared pairs. A proton moves in and begins to share one of those unshared pairs of electrons and forms a so-called coordinate-covalent bond with oxygen. Hydrogen ion (proton) is seen as a proton acceptor, and defined as an acid, while the water molecule that provides the electron pair is seen as and "electron pair donor", and therefore, a base.