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acids and bases

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 Brönsted-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 -  . For example, hydrochloric acid, HCl, dissociates in water to yield the required hydrogen ions, H + , and also chloride ions, Cl -  . The base sodium hydroxide, NaOH, dissociates in water to yield the required hydroxide ions, OH - , and also sodium ions, Na + . The Brönsted-Lowry Theory Some substances act as acids or bases when they are dissolved in solvents other than water, such as liquid ammonia. The Brönsted-Lowry theory, named for the Danish chemist Johannes Brönsted 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 Brönsted-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 ACIDS AND BASES

The name "acid" calls to mind vivid sensory images—of tartness, for instance, if the acid in question is meant for human consumption, as with the citric acid in lemons. On the other hand, the thought of laboratory-and industrial-strength substances with scary-sounding names, such as sulfuric acid or hydrofluoric acid, carries with it other ideas—of acids that are capable of destroying materials, including human flesh. The name "base," by contrast, is not widely known in its chemical sense, and even when the older term of "alkali" is used, the sense-impressions produced by the word tend not to be as vivid as those generated by the thought of "acid." In their industrial applications, bases too can be highly powerful. As with acids, they have many household uses, in substances such as baking soda or oven cleaners. From a taste standpoint, (as anyone who has ever brushed his or her teeth with baking soda knows), bases are bitter rather than sour. How do we know when something is an acid or a base? Acid-base indicators, such as litmus paper and other materials for testing pH, offer a means of judging these qualities in various substances. However, there are larger structural definitions of the two concepts, which evolved in three stages during the late nineteenth and early twentieth centuries, that provide a more solid theoretical underpinning to the understanding of acids and bases. Prior to the development of atomic and molecular theory in the nineteenth century, followed by the discovery of subatomic structures in the late nineteenth and early twentieth centuries, chemists could not do much more than make measurements and observations. Their definitions of substances were purely phenomenological—that is, the result of experimentation and the collection of data. From these observations, they could form general rules, but they lacked any means of "seeing" into the atomic and molecular structures of the chemical world. The phenomenological distinctions between acids and bases, gathered by scientists from ancient times onward, worked well enough for many centuries. The word "acid" comes from the Latin term acidus, or "sour," and from an early period, scientists understood that substances such as vinegar and lemon juice shared a common acidic quality. Eventually, the phenomenological definition of acids became relatively sophisticated, encompassing such details as the fact that acids produce characteristic colors in certain vegetable dyes, such as those used in making litmus paper. In addition, chemists realized that acids dissolve some metals, releasing hydrogen in the process. The word "alkali" comes from the Arabic al-qili, which refers to the ashes of the seawort plant. The latter, which typically grows in marshy areas, was often burned to produce soda ash, used in making soap. In contrast to acids, bases—caffeine, for example—have a bitter taste, and many of them feel slippery to the touch. They also produce characteristic colors in the vegetable dyes of litmus paper, and can be used to promote certain chemical reactions. Note that today chemists use the word "base" instead of "alkali," the reason being that the latter term has a narrower meaning: all alkalies are bases, but not all bases are alkalies. Originally, "alkali" referred only to the ashes of burned plants, such as seawort, that contained either sodium or potassium, and from which the oxides of sodium and potassium could be obtained. Eventually, alkali came to mean the soluble hydroxides of the alkali and alkaline earth metals. This includes sodium hydroxide, the active ingredient in drain and oven cleaners; magnesium hydroxide, used for instance in milk of magnesia; potassium hydroxide, found in soaps and other substances; and other compounds. Broad as this range of substances is, it fails to encompass the wide array of materials known today as bases—compounds which react with acids to form salts and water. The reaction to form salts and water is, in fact, one of the ways that acids and bases can be defined. In an aqueous solution, hydrochloric acid and sodium hydroxide react to form sodium chloride—which, though it is suspended in an aqueous solution, is still common table salt—along with water. The equation for this reaction is HCl(aq ) + NaOH(aq ) →H2O + NaCl(aq ). In other words, the sodium (Na) ion in sodium hydroxide switches places with the hydrogen ion in hydrochloric acid, resulting in the creation of NaCl (salt) along with water. But why does this happen? Useful as this definition regarding the formation of salts and water is, it is still not structural—in other words, it does not delve into the molecular structure and behavior of acids and bases. Credit for the first truly structural definition of the difference goes to the Swedish chemist Svante Arrhenius (1859-1927). It was Arrhenius who, in his doctoral dissertation in 1884, introduced the concept of an ion, an atom possessing an electric charge. His understanding was particularly impressive in light of the fact that it was 13 more years before the discovery of the electron, the subatomic particle responsible for the creation of ions. Atoms have a neutral charge, but when an electron or electrons depart, the atom becomes a positive ion or cation. Similarly, when an electron or electrons join a previously uncharged atom, the result is a negative ion or anion. Not only did the concept of ions greatly influence the future of chemistry, but it also provided Arrhenius with the key necessary to formulate his distinction between acids and bases. Arrhenius observed that molecules of certain compounds break into charged particles when placed in liquid. This led him to the Arrhenius acid-base theory, which defines an acid as any compound that produces hydrogen ions (H+) when dissolved in water, and a base as any compound that produces hydroxide ions (OH−) when dissolved in water. This was a good start, but two aspects of Arrhenius's theory suggested the need for a definition that encompassed more substances. First of all, his theory was limited to reactions in aqueous solutions. Though many acid-base reactions do occur when water is the solvent, this is not always the case. Second, the Arrhenius definition effectively limited acids and bases only to those ionic compounds, such as hydrochloric acid or sodium hydroxide, which produced either hydrogen or hydroxide ions. However, ammonia, or NH3, acts like a base in aqueous solutions, even though it does not produce the hydroxide ion. The same is true of other substances, which behave like acids or bases without conforming to the Arrhenius definition. These shortcomings pointed to the need for a more comprehensive theory, which arrived with the formulation of the Brønsted-Lowry definition by English chemist Thomas Lowry (1874-1936) and Danish chemist J. N. Brønsted (1879-1947). Nonetheless, Arrhenius's theory represented an important first step, and in 1903, he was awarded the Nobel Prize in Chemistry for his work on the dissociation of molecules into ions. The Brønsted-Lowry acid-base theory defines an acid as a proton (H+) donor, and a base as a proton acceptor, in a chemical reaction. Protons are represented by the symbol H+, and in representing acids and bases, the symbols HA and A−, respectively, are used. These symbols indicate that an acid has a proton it is ready to give away, while a base, with its negative charge, is ready to receive the positively charged proton. Though it is used here to represent a proton, it should be pointed out that H+ is also the hydrogen ion—a hydrogen atom that has lost its sole electron and thus acquired a positive charge. It is thus really nothing more than a lone proton, but this is the one and only case in which an atom and a proton are exactly the same thing. In an acid-base reaction, a molecule of acid is "donating" a proton, in the form of a hydrogen ion. This should not be confused with a far more complex process, nuclear fusion, in which an atom gives up a proton to another atom. The most fundamental type of acid-base reaction in Brønsted-Lowry theory can be symbolized thus HA(aq ) + H2O(l ) →H3O+(aq )


From Yahoo Answers

Question:The initial solution in a titration is 20ml of 0.2M HNO3. The nitric acid is titrated with 0.15M KOH in 5.0ml volumes. 1) What is the initial pH of the nitric acid? 2) What is the pH after the addition of 10.0ml KOH 3) How many times of titration are required to reach equivalence point

Answers:1) HNO3-->H+ NO3- strong electrolyte completely dissociated [H+] =0.2 pH = log1/[H+]= log 1/0.2=log 5 =0.7 2) HNO3+KOH---> KNO3+H2O moles in 20ml of 0.2M = 20*(10^-3)*0.2 = 4*10^-3mole 10mL of 0.15M KOH = 0.15*0.01=0.0015=1.5 *10^-3mole So, you obtain1.510^-3 mole of KNO3 and have( 4-1.5)10^-3 mole of HNO3 left (this in a volume of10+20 =30ml Concentraton of HNO3 unreacted= 2.5*10^-3/0.03=0.083M pH = log 1/0.083=log12=1.08 3)to neutralize, you need 20*0.2/0.15 =26.6mL you need 6 times at least to reach equivalence point

Question:Although HCl(aq) exhibits Arrehenius acidic properties, pure HCl gas and HCl dissolved in a nonpolar solvent exhibit no acidic properties in the Arrhenius sense. Explain why.

Answers:In order for something to be an Arrhenius acid, it has to be able to donate protons to the solvent. That's what happens with HCl(aq), but HCl gas cannot give up its H+ to air and HCl dissolved in nonpolar solvent cannot give H+ to the solvent b/c neither the air nor the nonpolar solvent have any affinity (attaction) for the H+.

Question:Can someone direct me to some animation/tutorial site that can teach me about acids and bases without the bore? I don't get indicators (in titration curve) and the whole titration curve stuff. Thanks.

Answers:Try this site: http://www2.wwnorton.com/college/chemistry/gilbert/tutorials/ch16.htm Hope this is what you are looking for!


From Youtube

Acids and Bases

Created by brettfishbin Using Memoov online animation studio. Go to memoov.com to create Your own Animations.

Acids and Bases

properties

Acids and Bases

Acids and Bases Verse 1: The first chemist to inform us, His name was Savant Arrhenius. He said acids give the ion H Plus, Bases give the ion OH Minus. HCl shows us a great example Hydronium forms in stock sample. NaOH is the other story Hydroxide makes the base inventory. Chorus: Acids and bases whats left to say? 3 different theories are together today Arrhenius, Bronsted-Lowry, Lewis All their ideas added to this Acids have a pH less than 7 pH 0 is as strong as heaven the pH of 7 is neutral Above which is base, must be well controlled. Verse 2: There was a problem with the first theory Solved by two chemists Bronsted and Lowry They proposed acids were proton donors And that bases were but proton receptors The result of this was acid-base reactions Always wanting equilibrium action Then the rise of conjugate pairs More acids and bases put out there Bridge: These two things can both be quite caustic If it hits you, you should test diagnostics Dont ever mess with the six strong acids They dissociate completely and will leave you roasted Verse 3: A guy called Lewis in 1963 He made the concept much harder for me Lewis bases donate electron pairs Lewis acids accept them there An adduct forms between the compounds This time the protons dont have to move around This is unlike Bronsted-Lowrys theory It makes the subject far from easy

Acids & Bases

Please to enjoy: Will, Maya, and Lois's AP Chemistry Acids & Bases commercial. Oh how lame this video is. Alright, eat it alll up.


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