2/3/2015 Acid-base Introduction Chem1 General Chemistry Virtual Textbook → Acidbase concepts → Introduction Introduction What is an acid? What is a base? ⇐ index intro | pH | Brønsted | competition | electrons | structures | gallery< This page is best viewed with a standardscompliant browser such as Firefox, Opera, Camino, or Safari. The concepts of an acid, a base, and a salt are very old On this page: ones that have undergone several major refinements as Acids Acids and the hydrogen ion Bases chemical science has evolved. Our treatment of the subject at this stage will be mainly conceptual and qualitative, emphasizing the definitions and fundamental ideas associated with acids and bases. We will not cover Neutralization What you should be able to do calculations involving acid-base equilibria in these lessons. Concept map 1 Acids The term acid was first used in the seventeenth century; it comes from the Latin root ac, meaning “sharp”, as in acetum, vinegar. Some early writers suggested that acidic molecules might have sharp corners or spinelike projections that irritate the tongue or skin. Acids have long been recognized as a distinctive class of compounds whose aqueous solutions exhibit the following properties: A characteristic sour taste (think of lemon juice!); ability to change the color of litmus* from blue to red; react with certain metals to produce gaseous H2; *Litmus is a natural dye found in certain lichens. The name is of Scandinavian origin, e.g. lit (color) + mosi (moss) in Icelandic. "Litmus test" has acquired a meaning that transcends both Chemistry and science to denote any kind of test giving a yes/no answer. react with bases to form a salt and water. How oxygen got misnamed The first chemistrybased definition of an acid turned out to be wrong: in 1787, Antoine Lavoisier, as part of his masterful classification of substances, identified the known acids as a separate group of the “complex substances” (compounds). Their special nature, he postulated, derived from the presence of some common element that embodies the “acidity” file://localhost/Users/Bill/Desktop/chem1vt/webtext/acid1/abcon-1.html 1/6 2/3/2015 Acid-base Introduction principle, which he named oxygen, derived from the Greek for “acid former”. Lavoisier had recently assigned this name to the new gaseous element that Joseph Priestly had discovered a few years earlier as the essential substance that supports combustion. Many combustion products (oxides) do give acidic solutions, and oxygen is in fact present in most acids, so Lavoisier’s mistake is understandable. In 1811 Humphrey Davy showed that muriatic (hydrochloric) acid (which Lavoisier had regarded as an element) does not contain oxygen, but this merely convinced some that chlorine was not an element but an oxygencontaining compound. Although a dozen oxygenfree acids had been discovered by 1830, it was not until about 1840 that the hydrogen theory of acids became generally accepted. By this time, the misnomer oxygen was too well established a name to be changed. The root oxy comes from the Greek word οξνς, which means "sour". 2 Acids and the hydrogen ion The key to understanding acids (as well as bases and salts) had to await Michael Faraday’s midnineteenth century discovery that solutions of salts (known as electrolytes) conduct electricity. This implies the existence of charged particles that can migrate under the influence of an electric field. Faraday named these particles ions (“wanderers”). Later studies on electrolytic solutions suggested that the properties we associate with acids are due to the presence of an excess of hydrogen ions in the solution. By 1890 the Swedish chemist Svante Arrhenius (18591927) was able to formulate the first useful theory of acids: anion." [image link] "an acidic substance is one whose molecular unit contains at least one hydrogen atom that can dissociate, or ionize, when dissolved in water, producing a hydrated hydrogen ion and an hydrochloric acid HCl → H+(aq) + Cl–(aq) sulfuric acid H2SO4→ H+(aq) + HSO4–(aq) hydrogen sulfite ion HSO3–(aq) → H+(aq) + SO32–(aq) acetic acid H3CCOOH → H+(aq) + H3CCOO–(aq) Strictly speaking, an “Arrhenius acid” must contain hydrogen. However, there are substances that do not themselves contain hydrogen, but still yield hydrogen ions when dissolved in water; the hydrogen ions come from the water itself, by reaction with the substance. A more useful operational definition of an acid is therefore the following: An acid is a substance that yields an excess of hydrogen ions when dissolved in water. file://localhost/Users/Bill/Desktop/chem1vt/webtext/acid1/abcon-1.html 2/6 2/3/2015 Acid-base Introduction There are three important points to understand about hydrogen in acids: Although all Arrhenius acids contain hydrogen, not all hydrogen atoms in a substance are capable of dissociating; thus the –CH3 hydrogens of acetic acid are “nonacidic”. An important part of knowing chemistry is being able to predict which hydrogen atoms in a substance will be able to dissociate into hydrogen ions; this topic is covered in a later lesson of this set. Those hydrogens that do dissociate can do so to different degrees. The strong acids such as HCl and HNO3 are effectively 100% dissociated in solution. Most organic acids, such as acetic acid, are weak; only a small fraction of the acid is dissociated in most solutions. HF and HCN are examples of weak inorganic acids. Acids that possess more than one dissociable hydrogen atom are known as polyprotic acids; H2SO4 and H3PO4 are wellknown examples. Intermediate forms such as HPO42–, being capable of both accepting and losing protons, are called ampholytes. H2SO4 sulfuric acid HSO4– hydrogen → sulfate ("bisulfate") ion H2S HS– hydrosulfide ion → SO42– sulfate ion S2– → sulfide ion hydrosulfuric acid → H3PO4 H2PO4– HPO42– → dihydrogen phosphate → hydrogen phosphate ion ion phosphoric acid HOOCCOOH oxalic acid → HOOCCOO– hydrogen oxalate ion → –OOCCOO– oxalate ion → PO43– phosphate ion You will find out in a later section of this lesson that hydrogen ions cannot exist as such in water, but don't panic! It turns out that chemists still find it convenient to pretend as if they are present, and to write reactions that include them. 3 Bases The name base has long been associated with a class of compounds whose aqueous solutions are characterized by: a bitter taste; a “soapy” feeling when applied to the skin; ability to restore the original blue color of litmus that has been turned red by acids; ability to react with acids to form salts. react with certain metals to produce gaseous H2; Just as an acid is a substance that liberates hydrogen ions into solution, a base yields hydroxide ions when dissolved in water: file://localhost/Users/Bill/Desktop/chem1vt/webtext/acid1/abcon-1.html The word “alkali” is often applied to strong inorganic bases. It is of Arabic origin, from alkali ("the ashes") which refers to the calcined wood ashes that were boiled with water to obtain potash 3/6 2/3/2015 Acid-base Introduction NaOH(s) → Na+(aq) + OH–(aq) which contains the strong base KOH, used in soap making. The element name potassium and its symbol K (from the Latin kalium) derive from these sources. Sodium hydroxide is an Arrhenius base because it contains hydroxide ions. However, other substances which do not contain hydroxide ions can nevertheless produce them by reaction with water, and are therefore also classified as bases. Two classes of such substances are the metal oxides and the hydrogen compounds of certain nonmetals: Na2O(s) + H2O → [2 NaOH] → 2 Na+(aq) + 2 OH–(aq) NH3 + H2O → NH4+(aq) + OH–(aq) We can therefore define a base as follows: A base is a substance that yields an excess of hydroxide ions when dissolved in water. 4 Neutralization Acids and bases react with one another to yield two products: water, and an ionic compound known as a salt. This kind of reaction is called a neutralization reaction. This "molecular" equation is convenient to write, but we need to recast it as a net ionic equation to reveal what is really going on here when the reaction takes place in water, as is almost always the case. H+ + Cl– + Na+ + OH–→ Na+ + Cl– + H2O If we cancel out the ions that appear on both sides (and therefore don't really participate in the reaction), we are left with the net equation H+(aq) + OH–(aq) → H2O (1) which is the fundamental process that occurs in all neutralization reactions. Confirmation that this equation describes all neutralization reactions that take place in water is provided by experiments indicating that no matter what acid and base are combined, all liberate the same amount of heat (57.7 kJ) per mole of H+ neutralized. In the case of a weak acid, or a base that is not very soluble in water, more than one step might be required. For example, a similar reaction can occur between acetic acid and calcium hydroxide to produce calcium acetate: 2 CH3COOH + Ca(OH)2 → CH3COOCa + 2 H2O file://localhost/Users/Bill/Desktop/chem1vt/webtext/acid1/abcon-1.html 4/6 2/3/2015 Acid-base Introduction If this takes place in aqueous solution, the reaction is really between the very small quantities of H+ and OH– resulting from the dissociation of the acid and the dissolution of the base, so the reaction is identical with (1): H+(aq) + OH–(aq) → H2O If, on the other hand, we add solid calcium hydroxide to pure liquid acetic acid, the net reaction would include both reactants in their "molecular" forms: 2 CH3COOH(l) + Ca(OH)2 (s) → 2 CH3COO– + Ca2+ + 2 H2O The “salt” that is produced in a neutralization reaction consists simply of the anion and cation that were already present. The salt can be recovered as a solid by evaporating the water. What you should be able to do Make sure you thoroughly understand the following essential ideas which have been presented above. Suggest simple tests you could carry out to determine if an unknown substance is an acid or a base. State the chemical definitions of an acid and a base in terms of their behavior in water. Write the formula of the salt formed when a given acid and base are combined. Concept Map file://localhost/Users/Bill/Desktop/chem1vt/webtext/acid1/abcon-1.html 5/6 2/3/2015 Acid-base Introduction Page last modified: 09.09.2010 ⇐ index intro | pH | Brønsted | competition | electrons | structures | gallery © 2007 by Stephen Lower Simon Fraser University Burnaby/Vancouver Canada For information about this Web site or to contact the author, please see the Chem1 Virtual Textbook home page. This work is licensed under a Creative Commons AttributionNonCommercial 2.5 License. file://localhost/Users/Bill/Desktop/chem1vt/webtext/acid1/abcon-1.html 6/6 2/3/2015 pH and titration Chem1 General Chemistry Virtual Textbook →Acidbase concepts →pH and titration pH and titration Aqueous solutions of acids and bases ⇐ index | intro pH Brønsted | competition | electrons | structures | gallery< This page is best viewed with a standardscompliant browser such as Firefox, Opera, Camino, or Safari. On this page: Ion product of water pH definition The pH scale Acidbase titration What you should be able to do Concept map As you will see in the lesson that follows this one, water plays an essential role in acid-base chemistry as we ordinarily know it. To even those who know very little about chemistry, the term pH is recognized as a measure of "acidity", so the major portion of this unit is devoted to the definition of pH and of the pH scale. But since these topics are intimately dependant on the properties of water and its ability do dissociate into hydrogen and hydroxyl ions, we begin our discussion with this topic. We end this lesson with a brief discussion of acid-base titration— probably the most frequently carried-out chemistry laboratory operation in the world. 1 Dissociation of water The ability of acids to react with bases depends on the tendency of hydrogen ions to combine with hydroxide ions to form water: H+(aq) + OH–(aq) → H2O (1) All chemical reactions that take place in a single phase (such as in a solution) are theoretically "incomplete" and are said to be reversible. This tendency happens to be very great, so the reaction is practically complete — but not "completely" complete; a few stray H+ and OH– ions will always be present. What's more, this is true even if you start with the purest water attainable. This means that in pure water, the reverse reaction, the "dissociation" of water H2O → H+(aq) + OH–(aq) (2) will proceed to a very slight extent. Both reactions take place simultaneously, but (1) is so much faster than (2) that only a minute fraction of H2O molecules are dissociated. Liquids that contain ions are able to conduct an electric current. Pure water is practically an insulator, but careful experiments show that even the most highly purified water exhibits a very slight conductivity that corresponds to a concentration of both the H+ ion and OH– ions of almost exactly 1.00 × 10– 7mol L–1 at 25°C. file://localhost/Users/Bill/Desktop/chem1vt/webtext/acid1/abcon-2.html 1/8 2/3/2015 pH and titration Problem Example 1 What fraction of water molecules in a litre of water are dissociated Solution: 1 L of water has a mass of 1000 g. The number of moles in 1000 g of H2O is (1000 g)/(18 g mol–1) = 55.5 mol. This corresponds to (55.5 mol)(6.02E23 mol1) = 3.34E25 H2O molecules. An average of 107 mole, or (107)(6.02E23) = 6.0E16 H2O molecules will be dissociated at any time. The fraction of dissociated water molecules is therefore (6.0E16)/(3.3E25) = 1.8E–9. Thus we can say that only about two out of every billion (109) water molecules will be dissociated. Ion product of water The degree of dissociation of water is so small that you might wonder why it is even mentioned here. The reason stems from an important relationship that governs the concentrations of H+ and OH– ions in aqueous solutions: [H+][OH–] = 1.00 × 10–14 (3) must know this! in which the square brackets [ ] refer to the concentrations (in moles per litre) of the substances they enclose. The quantity 1.00 x 10–14 is commonly denoted by Kw. Its value varies slightly with temperature, pressure, and the presence of other ions in the solution. This expression is known as the ion product of water, and it applies to all aqueous solutions, not just to pure water. The consequences of this are far reaching, because it implies that if the concentration of H+ is large, that of OH– will be small, and vice versa. This means that H+ ions are present in all aqueous solutions, not just acidic ones. This leads to the following important definitions, which you must know: acidic solution [H+] > [OH–] alkaline ("basic") solution [H+] < [OH–] neutral solution [H+] = [OH–] = 1.00×10–7 mol L–1 Take special note of the followng definition: A neutral solution is one in which the concentrations of H+ and OH– ions are identical. The values of these concentrations are constrained by Eq. 3. Thus, in a neutral solution, both the hydrogen and hydroxide ion concentrations are 1.00 × 10– 7 mol L–1: file://localhost/Users/Bill/Desktop/chem1vt/webtext/acid1/abcon-2.html 2/8 2/3/2015 pH and titration [H+][OH–] = [1.00 × 10–7][1.00 × 10–7] =1.00 × 10–14 Hydrochloric acid is a typical strong acid that is totally dissociated in solution: HCl → H+(aq) + Cl–(aq) A 1.0M solution of HCl in water therefore does not really contain any significant concentration of HCl molecules at all; it is a solution in of H+ and Cl– in which the concentrations of both ions are 1.0 mol L–1. The concentration of hydroxide ion in such a solution, according to Eq 2, is [OH–] = (Kw)/[H+] = (1.00 x 10–14) / (1 mol L–1) = 1.00 x 10–14 mol L–1. Similarly, the concentration of hydrogen ion in a solution made by dissolving 1.0 mol of sodium hydroxide in water will be 1.00 x 10–14 mol L–1. 2 pH When dealing with a range of values (such as the variety of hydrogen ion concentrations encountered in chemistry) that spans many powers of ten, it is convenient to represent them on a more compressed logarithmic scale. By convention, we use the pH scale to denote hydrogen ion concentrations: pH = – log10 [H+] (4) must know this! or conversely, [H+] = 10–pH . This notation was devised by the Danish chemist Søren Sørensen (18681939) in 1909. There are several accounts of why he chose "pH"; a likely one is that the letters stand for the French term pouvoir hydrogène, meaning "power of hydrogen"— "power" in the sense of an exponent. It has since become common to represent other small quantities in "p notation". Two that you need to know in this course are the following: pOH = – log10 [OH–] pKw = – log Kw (= 14 when Kw = 1.00 × 10–14) Note that pH and pOH are expressed as numbers without any units, since logarithms must be dimensionless. Recall from Eq 3 that [H+][OH–] = 1.00 × 10–14; if we write this in "p notation" it becomes pH + pOH = 14 (5) must know this! In a neutral solution at 25°C, pH = pOH = 7.0. As pH increases, pOH diminishes; a higher pH corresponds to an alkaline solution, a lower pH to an acidic solution. In a solution with [H+] = 1 M , the pH would be 0; in a 0.00010 M solution of H+, it would be 4.0. Similarly, a 0.00010 M solution of file://localhost/Users/Bill/Desktop/chem1vt/webtext/acid1/abcon-2.html 3/8 2/3/2015 pH and titration NaOH would have a pOH of 4.0, and thus a pH of 10.0. It is very important that you thoroughly understand the pH scale, and be able to convert between [H+] or [OH–] and pH in both directions. Problem Example 2 The pH of blood must be held very close to 7.40. Find the hydroxide ion concentration that corresponds to this pH. Solution: The pOH will be (14.0 – 7.40) = 6.60. [OH–] = 10–pOH = 10–6.6 = 2.51 x 10–7 M The pH scale The range of possible pH values runs from about 0 to 14. The word "about" in the above statement reflects the fact that at very high concentrations (10 M hydrochloric acid or sodium hydroxide, for example,) a significant fraction of the ions will be associated into neutral pairs such as H+·Cl–, thus reducing the concentration of “available” ions to a smaller value which we will call the effective concentration. It is the effective concentration of H+ and OH– that determines the pH and pOH. For solutions in which ion concentrations don't exceed 0.1 M, the formulas pH = –log [H+] and pOH = –log[OH–] are generally reliable, but don't expect a 10.0 M solution of a strong acid to have a pH of exactly –1.00! The table shown here will help give you a general feeling for where common substances fall on the pH scale. Notice especially that most foods are slightly acidic; the principal "bodily fluids" are slightly alkaline, as is seawater— not surprising, since early animal life began in the oceans. the pH of freshlydistilled water will drift downward as it takes up carbon dioxide from the air; CO2 reacts with water to produce carbonic acid, H2CO3. the pH of water that occurs in nature varies over a wide range. Groundwaters often pick up additional CO2 respired by organisms in the soil, but can also become alkaline if they are in contact with carbonatecontaining sediments. "Acid" rain is by definition more acidic than pure water in equilibrium with atmospheric CO2, owing mainly to sulfuric and nitric acids that originate from fossilfuel emissions of nitrogen oxides and SO2. pH indicators file://localhost/Users/Bill/Desktop/chem1vt/webtext/acid1/abcon-2.html 4/8 2/3/2015 pH and titration The colors of many dyelike compounds depend on the pH, and can serve as useful indicators to determine whether the pH of a solution is above or below a certain value. Here is a list of ordinary foods and household substances that can serve as indicators Natural indicator dyes The best known of these is of course litmus, which has served as a means of distinguishing beween acidic and alkaline substances since the early 18th century. Many flower pigments are also dependent on the pH. You may have noticed that the flowers of some hydrangea shrub species are blue when grown in acidic soils, and white or pink in alkaline soils. Red cabbage is a popular makeityourself indicator. Here is a typical recipe. Universal indicators Most indicator dyes show only one color change, and thus are only able to determine whether the pH of a solution is greater or less than the value that is characteristic of a particular indicator. By combining a variety of dyes whose color changes occur at different pHs, a "universal" indicator can be made. Commerciallyprepared pH test papers of this kind are available for both wide and narrow pH ranges. 3 Titration Since acids and bases readily react with each other, it is experimentally quite easy to find the amount of acid in a solution by determining how many moles of base are required to neutralize it. This operation is called titration, and you should already be familiar with it from your work in the Laboratory. We can titrate an acid with a base, or a base with an acid. The substance whose concentration we are determining (the analyte) is the substance being titrated; the substance we are adding in measured amounts is the titrant. The idea is to add titrant until the titrant has reacted with all of the analyte; at this point, the number of moles of titrant added tells us the concentration of base (or acid) in the file://localhost/Users/Bill/Desktop/chem1vt/webtext/acid1/abcon-2.html 5/8 2/3/2015 pH and titration solution being titrated. 36.00 ml of a solution of HCl was titrated with 0.44 M KOH. The volume of KOH solution required to neutralize the acid solution was 27.00 ml. What was the concentration of the HCl? Solution: The number of moles of titrant added was (.027 L)(.44 mol L–1) = .0119 mol. Because one mole of KOH reacts with one mole of HCl, this is also the number of moles of HCl; its concentration is therefore (.0119 mol) ÷ (.036 L) = 0.33 M . Titration curves The course of a titration can be followed by plotting the pH of the solution as a function of the quantity of titrant added. The figure shows two such curves, one for a strong acid (HCl) and the other for a weak acid, acetic acid, denoted by HAc. Looking first at the HCl curve, notice how the pH changes very slightly until the acid is almost neutralized. At that point, which corresponds to the vertical part of the plot, just one additional drop of NaOH solution will cause the pH to jump to a very high value— almost as high as that of the pure NaOH solution. Compare the curve for HCl with that of HAc. For a weak acid, the pH jump near the neutralization point is less steep. Notice also that the pH of the solution at the neutralization point is greater than 7. These two characteristics of the titration curve for a weak acid are very important for you to know. If the acid or base is polyprotic, there will be a jump in pH for each proton that is titrated. In the example shown here, a solution of carbonic acid H2CO3 is titrated with sodium hydroxide. The first equivalence point (at which the H2CO3 has been converted entirely into bicarbonate ion HCO3–) occurs at pH 8.3. The solution is now identical to one prepared by dissolving an identical amount of sodium bicarbonate in water. Addition of another mole equivalent of hydroxide ion converts the bicarbonate into carbonate ion and is complete at pH 10.3; an identical solution could be prepared by dissolving the appropriate amount of sodium carbonate in water. Finding the equivalence point: indicators When enough base has been added to react completely with the hydrogens of a monoprotic acid, the equivalence point has been reached. If a strong acid and strong base are titrated, the pH of the solution will be 7.0 at the equivalence point. However, if the acid is a weak one, the pH will be greater than 7; the file://localhost/Users/Bill/Desktop/chem1vt/webtext/acid1/abcon-2.html 6/8 2/3/2015 pH and titration “neutralized” solution will not be “neutral” in terms of pH. For a polyprotic acid, there will be an equivalence point for each titratable hydrogen in the acid. These typically occur at pH values that are 45 units apart, but they are occasionally closer, in which case they may not be readily apparent in the titration curve. The key to a successful titration is knowing when the equivalance point has been reached. The easiest way of finding the equivalence point is to use an indicator dye; this is a substance whose color is sensitive to the pH. One such indicator that is commonly encountered in the laboratory is phenolphthalein; it is colorless in acidic solution, but turns intensely red when the solution becomes alkaline. If an acid is to be titrated, you add a few drops of phenolphthalein to the solution before beginning the titration. As the titrant is added, a local red color appears, but quickly dissipates as the solution is shaken or stirred. Gradually, as the equivalence point is approached, the color dissipates more slowly; the trick is to stop the addition of base after a single drop results in a permanently pink solution. Different indicators change color at different pH values; see here for an illustrated list. Since the pH of the equivalance point varies with the strength of the acid being titrated, one tries to fit the indicator to the particular acid. One can titrate polyprotic acids by using a suitable combination of several indicators, as is illustrated above for carbonic acid. What you should be able to do Make sure you thoroughly understand the following essential ideas which have been presented above. It is especially imortant that you know the precise meanings of all the highlighted terms in the context of this topic. Define the ion product of water, and know its roomtemperature value. State the criteria, in terms of H+ and OH– concentrations, of an acidic, alkaline, and a neutral solution. Given the effective hydrogen ion concentration in a solution, calculate the pH. Conversely, find the hydrogen or hydroxide ion concentration in a solution having a given pH. Find the pH or pOH of a solution when either one is known. Describe the process of acidbase titration, and explain the significance of the equivalence point. Sketch a typical titration curve for a monoprotic or polyprotic acid. Concept Map file://localhost/Users/Bill/Desktop/chem1vt/webtext/acid1/abcon-2.html 7/8 2/3/2015 pH and titration Page last modified: 17.02.2007 ⇐ index | intro pH Brønsted | competition | electrons | structures | gallery< © 2007 by Stephen Lower Simon Fraser University Burnaby/Vancouver Canada For information about this Web site or to contact the author, please see the Chem1 Virtual Textbook home page. This work is licensed under a Creative Commons AttributionNonCommercial 2.5 License. file://localhost/Users/Bill/Desktop/chem1vt/webtext/acid1/abcon-2.html 8/8
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