Property Acid Base Taste sour bitter Feel nothing characteristic soapy

Transcription

1 3 ACIDS AND BASES 3.1 WAT ARE TEY - A ISTORICAL VIEW Acids and bases are among the oldest recognised group of chemicals. Long before atoms and molecules, elements and compounds were known about, the properties and behaviour of certain substances led them to be grouped under these two headings. Table 3.1 lists some of the significant properties of acids and bases. TABLE 3.1 Comparison of properties of acids and bases Property Acid Base Taste sour bitter Feel nothing characteristic soapy Reaction with metals Reaction with each other dissolve many metals, forming salt and hydrogen Formation of a solution lacking in characteristic properties - neutralisation dissolve few metals, forming hydroxide and oxygen???? Formation of a solution lacking in characteristic properties - neutralisation Reaction with litmus turns blue litmus red turns red litmus blue As chemistry developed in the nineteenth century, attempts were made to give a chemical explanation to the group of substances. One of the first theories of the composition of acids and bases was proposed by Arrhenius in 1884, who defined as an acid as a solution containing a high concentration of + ions, and a base as a solution containing a high concentration of O - ions. This theory was limited by its use of solutions as part of the definition (gaseous hydrogen chloride still reacts with solid sodium hydroxide to form sodium chloride and water), and failed to explain why certain substances, such as ammonia and sodium carbonate, were bases, when there was no O in their structure. A better theory was proposed in 1923 by two chemists, Lowry and Brönsted, who were working independently of each other in the same way that Mendeleev and Mayer produced almost identical periodic tables at the same time. owever, in this case, both chemists have gained recognition for their work: the Lowry-Brönsted theory acids and bases. A Lowry-Brönsted acid is any substance that donates a proton to another in a chemical reaction. A Lowry-Brönsted base is any substance that accepts a proton from another in a chemical reaction. Now, there is no mention of solutions or hydroxide ions, but simply proton transfer in a chemical reaction: their definition of an acid-base reaction. ow does this definition help the problem of ammonia behaving as a base? Ammonia can accept a proton, in or out of solution, because of the lone pair of electrons on the nitrogen atom. These are used to form a bond between the nitrogen and the hydrogen ion, as shown in Figure 3.1. In this case, water is acting as the acid, by donating a proton, and in doing so, forms hydroxide ion. ow then does the idea that acids produce protons fit into this definition? In water, an acid donates its proton to a water molecule, forming the hydronium ion, 3 O +, as shown in Equation 3.1, where A is a generalised formula for an acid. Chemistry 2

2 N : O N : + O _ FIGURE 3.1 Ammonia (base) accepting a proton from water (acid) A + 2 O 3 O + + A - Eqn 3.1 If A is acting as an acid, then the water molecule that is accepting the proton must be acting as a base. But hang on, in the ammonia example, water was accepting as an acid. Isn t there something strange about this? Surely, there can t be a substance which acts as both acid and base. The answer is yes: there are substances which can act as proton acceptor in one reaction, and acceptor in another. Such a substance is said to be amphoteric. Water is the most important example. What does this mean for aqueous acid-base reactions? Since the acid transfers its proton to a water molecule, and the base one way or another, liberates hydroxide ion, the real reaction is shown in Equation O + + O O Eqn 3.2 The reaction of an acid or base with water to form hydronium or hydroxide is often called dissociation. Such processes occur only in water to any significant extent, but substances can still behave as acids or bases in other solvents. To conclude, a distinction should be made between the terms base and alkali, which are often used interchangeably. Strictly speaking, a base is any substance which corresponds to the properties and behaviour described in this section, whereas an alkali is a watersoluble base. The term base will be used to describe the general class of compounds, and alkali (or alkaline solution) when referring to aqueous solutions only. The importance of acids and bases in many areas of commerce and industry for centuries has led to names, very commonly used, which were given long before the systematic approach came into being. While you are not expected to memorise these names, you should be familiar with them, because of their widespread use. Table 3.2 lists some of the more important ones. TABLE 3.2 Some common names for acids and bases Formula Systematic Name Common Name Cl hydrochloric acid muriatic acid 2 SO 4 sulfuric acid oil of vitriol NO 3 nitric acid NaO sodium hydroxide caustic soda Na 2 CO 3 sodium carbonate CaO calcium oxide lime KO potassium hydroxide caustic potash K 2 CO 3 potassium carbonate Chemistry 2 3.2

3 PRACTICE QUESTION 1. Explain why Arrhenius would not have understood the distinction between base and alkaline solution. 2. What are the important differences between the theories of Arrhenius and Lowry and Brönsted? 3. Write equations for the dissociation of (a) F 2- (b) CO 3 (assume reaction with 2 water molecules) and (c) ethanoic acid 4. Write an equation which shows hydrogen carbonate ion acting as (a) an acid and (b) a base. What name is given to hydrogen carbonate? 5. Write equations showing the reaction of Cl and NaO where (a) both reactants in are aqueous solution and (b) each reactant is in the pure form (Cl is a gas, NaO a solid). Clearly show the proton transfer in each case. 3.2 ACID-BASE BEAVIOUR IN WATER It was seen in Section 3.1 that acids in water form hydronium ions, while bases form hydroxide by reaction with water molecules. Some acidic or basic compounds undergo complete non-reversible dissociation in water, while others only partially react in an equilibrium process. Acid or base substances that dissociate completely in water to form 3 O + or O - are called strong acids (or bases). Those which only partly dissociate are called weak acids (or bases). Table 5.3 classifies the common acids and bases as strong or weak. TABLE 3.3 Classification of common acids and bases by strength Strong Acids Strong Bases Weak Acids Weak Bases Cl Br 2 SO 4 NO 3 ClO 4 all soluble metal hydroxide salts, e.g. NaO, KO, Ca(O) 2 F all alkanoic acids, e.g. ethanoic acid N 3 metal salts of carbonate and hydrogen carbonate ions At the molecular level, the distinction between strong and weak is clear. When a strong acid, such as Cl, dissolves in water, every single molecule of Cl dissociates into 3 O + and Cl -. On the other hand, a solution of a weak acid, such as ethanoic acid, will contain some 3 O + and C 3 COO -, but also molecular C 3 COO. This is illustrated in Figure 3.2. The extent to which a weak acid dissociates is dependent on the acid itself and also its concentration. This will be discussed further in Section STRONG WEAK molecular form dissociated form FIGURE 3.2 Dissociation of strong and weak acids ow can the distinction between strong and weak be shown in practical terms? By using conductivity measurements to show that one solution has many more ions than another of the same concentration of solute. Conductivity is the ability of a solution to conduct electricity. Chemistry 2 3.3

4 In practical terms of usage, all acid and base solutions should be considered as potentially dangerous, but it is the strong acids and bases in concentrated form, which pose the most risk. owever, F solutions are very hazardous, but not primarily because of their acidity. Do weak acids react completely with strong bases? If a weak acid only partly reacts with water, then maybe you might be thinking that it reacts similarly with NaO. The answer is no: all acids react fully with NaO in water. EXAMPLE 25 ml of a solution of NaO reacts with 25 ml of M Cl. What volume of 0.1 M ethanoic acid would be required to react with another 25 ml of the NaO? Exactly the same volume, since ethanoic and hydrochloric acids each react 1:1 with NaO. The fact that ethanoic acid is a weak acid makes no difference to its reaction with a strong base. PRACTICE QUESTIONS 9. Explain clearly the difference between strong and concentrated, and weak and dilute, in terms of solutions of acids and bases. 10. ow does Le Chatelier s Principle support the statement that a weak acid reacts completely with NaO? 3.3 P - A MEASURE OF ACIDITY The fact that that not all acids (or bases) are equal, the concept of how acidic (or basic) a solution introduces the question: how then do we measure the acidity (or basicity) or a solution? Is it due to concentration or strength of dissociation, or perhaps a combination of each? The answer is the latter: if solution acidity is defined as the concentration of 3 O + ions, then a concentrated solution of a weak acid can be more acidic than a dilute solution of a strong acid. The concentration of hydronium or hydroxide ions could be quoted simply in molarities, but for many solutions, this would involve using very small numbers, e.g M or M, which are inconvenient. So, an acidity scale has been developed which converts the sometimes awkward molarity values to easier numbers. It is known as p: the negative logarithm of the molarity of 3 O + in any aqueous solution. In equation form, this is shown in Equation 3.5. p = -log 10 [ 3 O + ] Eqn 3.5 The basicity of a solution is similarly defined: po: the negative logarithm of the molarity of O - in any aqueous solution (see Equation 3.6). po = -log 10 [O - ] Eqn 3.6 Chemistry 2 3.4

5 EXAMPLES 1. What is the p of a solution of 0.1 M Cl? Since Cl is a strong acid, every Cl molecule produces a 3 O + Therefore, the molarity of 3 O + is 0.1 M. p = -log 10 [0.1] = 1. ion in solution. 2. What is the po of a solution of 0.1 M N 3? Since ammonia is a weak acid, we can t (at least not at this point) say exactly what the concentration of hydroxide ions will be. It will definitely be less than 0.1 M (full dissociation), but the po cannot be calculated with current knowledge. po (and p) values get larger with lower concentrations (because of the negative sign). Therefore, we can say the po of 0.1 M N 3 will be greater than What is the po of M Mg(O) 2? When the strong base magnesium hydroxide dissociates, it liberates two hydroxide ions per molecule. So, [O - ] will be 2 x = M. po = -log 10 [0.0005] = 3.3. The reverse process - converting p or po back to a concentration - simply requires that you undo the logarithm, by using the 10 x button on your calculator, after changing the sign of the p (or po). The px idea is used to simplify small numbers in a number of instances, as you will see later in this chapter. Equations 3.7a and 3.7b summarise the relationship between X (which may be a concentration or an equilibrium constant or some other number) and px. px = -log 10 (X) Eqn 3.7a X = 10 -px Eqn 3.7b EXAMPLE What is the concentration of hydronium ions in a solution with a p of 5.2? Adapting Equation 3.7b, [ 3 O + ] = 10 -p. So, [ 3 O + ] = = 6.3 x 10-6 M. Recalling Equation 3.4, and carrying out the negative log 10 process on both sides of the equation, we obtain a simple relationship between p and po, as in Equation 3.8. pk w = p + po Eqn 3.8 = 14 (at 25 C) EXAMPLES 1. What is the p of a solution with a po of 5? Since p + po = 14, p must be What is the p of a neutral solution? A neutral solution has a 3 O + concentration of 1 x Therefore, its p is 7. Chemistry 2 3.5

6 p has become the predominant scale used, even for neutral and alkaline solutions. Theoretically the limits of the p scale are boundless, but in real terms, the most acidic solution known is concentrated sulfuric acid, with a p of around -1.6, while the most alkaline solution is a saturated solution of cesium hydroxide, which has a p of between 15 and 16 (depending on temperature). Figure 3.3 shows how a range of familiar solutions fit across the p scale. The p scale is logarithmic, which means that a difference of 1 p unit is a difference of ten times in concentration. The difference between the p of stomach acid and lemon juice may not sound much, but it means a difference of eight times in [ 3 O + ]. Acidic solutions are now defined as having p values less than 7, neutral solutions of p 7 exactly, and alkaline solution of p values greater than 7. stomach acid 1.3 wine 3.3 river water 8 ACIDIC NEUTRAL ALKALINE p<7 p=7 p>7-1 battery acid 2.5 lemon juice 6 rain water 7.35 blood 11 washing powder FIGURE 3.3 The p scale Measurement of p Indicators, such as litmus, change from one colour to another, when the solution they are placed in, reaches a certain p level. Different indicators have different p values for their colour changes. There are literally hundreds of indicators, but only a few that are used widely. You will use indicators for the purpose of chemical analysis in a laboratory module. In this subject, you will have a brief introduction to their use in estimating p. Table 3.4 lists some of the more commonly used indicators and their colour change characteristics. TABLE 3.4 Characteristics of some common indicators Indicator Colour At Low p p At Colour Change Colour At igh p Litmus red 5-8 blue Methyl orange red 3-4 yellow Phenolphthalein colourless 8-10 pink Bromothymol blue yellow 6-8 blue Chemistry 2 3.6

7 EXAMPLES Estimate the p of solutions, which give the following indicator colours. 1. Litmus is blue. If litmus is blue, then the p must be greater than Methyl orange and bromothymol blue are both yellow. If methyl orange is yellow, the p is greater than 4, but it is less than 6 because bromothymol blue is yellow. Therefore, the p is between 4 and 6. p can be measured accurately (to at one decimal place), using a special electrode and meter - a p meter. Its use is covered in detail in other modules, including Chemical Laboratory Techniques. PRACTICE QUESTIONS 11. Calculate the p of the following solutions of strong acids or bases. (a) 0.5 M Cl (b) 0.5 M NaO (c) 0.25 M 2 SO 4 (d) M Ca(O) Solution A has a p of 3 and solution B a p of 5. Which is the more acidic? By what factor is it more acidic? 13. Will the p of 0.1 M ethanoic acid be: (a) exactly 1 (b) slightly more than 1 (c) greater than 1 (d) around Explain why the p of 0.1 M N 3 is about the same as M NaO. 15. Calculate the concentration of 3 O + in solutions with p values of (a) 4 and (b) Calculate the concentration of O - in solutions with p values of (a) 2 and (b) pk w at 60 C is What is the p of neutral solutions at this temperature? 18. Draw a po scale, corresponding to Figure Estimate the p of the solutions, given the colour of the indicators below. Phenolphthalein Methyl Orange (a) colourless yellow (b) pink yellow (c) colourless red 3.4 ACID-BASE EQUILIBRIA The partial dissociation of weak acids and bases are equilibrium processes. Therefore, equilibrium constant expressions can be written, as in Equations 3.9a and 3.9b, respectively (A and B are generalised forms of acid and base). In each case, the [ 2 O] part of the expression has been incorporated into the constant value, giving the acid-dissociation constant, K a, and the basedissociation constant, K b. [ 3 O + ][A - ] K a = Eqn 3.9a [A] [B + ][O - ] K b = Eqn 3.9b [B] Chemistry 2 3.7

8 The K a and K b values are a guide to the extent to which a weak acid or base dissociates: the larger the value, the more dissociation occurs. These equations can also be used to calculate the p of weak acid or base solutions. K a and K b values are generally small, and are frequently listed in reference books using the px system: pk a and pk b, respectively. Table 3.5 lists some of the common acids and bases. TABLE 3.5 pk a and pk b values of common weak acids and bases Acid pk a Base pk b ethanoic 4.74 ammonia 4.74 hydrofluoric 3.45 hydrogen carbonate 7.63 phosphoric 2.12, 7.13, carbonate 3.75, 7.63 NOTE: the fact that the dissociation constants for ethanoic acid and ammonia are equal is one of life s coincidences, and nothing more. You will note that phosphoric acid and carbonate ion have more than one dissociation constant. This is because their reaction with water occurs in multiple steps, as shown for phosphoric acid below. Phosphoric acid is known as a triprotic acid, because of the three acidic hydrogen atoms, whereas carbonate is a diprotic base, because it can accept two protons. 3 PO O 2 PO O + pk a1 2 PO O PO O + pk a2 PO O PO O + pk a3 The calculation of p values in solutions of such species is not easy, and will not be expected in this module. EXAMPLE Calculate the p of a 0.1 M C 3 COO solution. The dissociation of ethanoic acid in water can be written as follows. C 3 COO + 2 O 3 O + + C 3 COO - Initial conc The zero concentration for hydronium ions in the water is, of course, not strictly correct, but you will see that the real value (10-7 ) is too small to interfere with our calculations. If y mole/l of the acid dissociates at equilibrium, the new concentrations will be: Equil. conc y y y K a must be calculated from the pk a : K a = 10 -pka = = 1.8 x Substituting these values into the K a expression gives: y.y 1.8 x 10-5 = y Chemistry 2 3.8

9 The equation could be solved for y at this point, but it can be simplified with an assumption: the amount of dissociation (y) is relatively small, meaning that (0.1-y) is almost 0.1. This simplifies the calculations, without introducing too much error. Rearranging the equation now gives: y 2 = 0.1 x 1.8 x 10-5 y = 1.34 x 10-3 M = [ 3 O + ] p = -log 10 (y) = 2.87 NOTE: the assumption that the amount of dissociation was sufficiently small to allow the simplified calculation only holds when (K a concentration of acid) is less than The same applies for calculations relating to weak bases. In this book, you will only be required to perform calculations that allow use of this assumption. Let s now look more carefully at the other product of the dissociation of acids and bases in water: the species produced when a proton is lost or gained, as the case may be. Ethanoic acid will be used as the example. C 3 COO + 2 O C 3 COO O + The reversibility of the equation means that the ethanoate produced in the dissociation could (in the reverse reaction) accept a proton from the hydronium ion. That means that the ethanoate ion is acting as a base. The pair of species, differing by one hydrogen ion - ethanoic acid and ethanoate ion - are called a conjugate acid-base pair. What happens if some of a metal salt of ethanoate (e.g. sodium ethanoate) is dissolved in water? Since water is capable of acting as an acid, it is reasonable to expect that the ethanoate will act as a base in this instance. The sodium ion is a spectator. C 3 COO O C 3 COO + O - DO STRONG ACIDS AVE CONJUGATE BASES? This solution would be expected to be alkaline, and indeed it is. In summary, when a weak acid reacts with a base (strong or weak), the products are another acid and base: the conjugate partners of the reactants. The answer is yes: the conjugate base of Cl is Cl -. The next obvious question is then: why isn t a solution of NaCl alkaline? Cl is a strong acid in water, meaning that its conjugate base in very, very A + B A - + B + weak. So weak in fact, that water is a acid1 base1 base2 acid2 stronger base, and any basic effects of chloride are masked by the water. Thus, the O - contributed by the autoionisation water determines the p of a It is easy to show that there is a simple relationship between K a and K b for a conjugate acid-base pair, as salt solution. shown in Equations 3.10a and 3.10b. If NaCl was dissolved in another solvent, K a K b = K w then it would have measurable basic properties. Eqn 3.10a pk a + pk b = 14 (at 25 C) Eqn 3.10b This simple relationship has led to dissociation constants for bases sometimes being tabulated as dissociation constants for the conjugate acids, meaning the use of one table of data. Chemistry 2 3.9

10 EXAMPLE What is the K b of pyridine, if its conjugate acid has a pk a of 5.25? pk b = 14 - pk a = K b = 10 -pkb = 1.78 x PRACTICE QUESTIONS 20. Which of the following solutions can have their p calculated using the assumption described in the above example. For those that can, calculate the p and the percentage dissociation. (a) 0.01 M C 3 COO (b) 0.1 M N 3 (c) M N 3 (d) 0.05 M F (e) 0.25 M NaCO Calculate the percentage dissociation if 0.01 moles of C 3 COO is dissolved in a litre of 0.01 M Cl. 22. (a) What is the conjugate acid of water? (b) What is the conjugate base of water? 23. What is the conjugate acid of ammonia? Write an equation showing what happen when the chloride salt of this conjugate acid was dissolved in water. 24. Prove Equation 3.10a. 25. Rewrite the bases section in Table 3.5 as pka s of the corresponding conjugate acids. 26. Which base dissociates more: ammonia, ethanoate or carbonate? 27. Explain the following: the weaker the acid, the stronger its conjugate base. 28. Why isn t the fact that the dissociation constant for CO 3 - and the second dissociation constant for CO 3 2- are identical just a coincidence? 3.5 BUFFER SOLUTIONS As you saw in Figure 3.3, the p of our blood is This figure is quite critical: if the p of our blood changes by more than 0.05 units, serious health effects might occur. With all the different types of foods that we eat, how can the bloodstream possibly maintain a consistent p level. Our blood contains a significant concentration of the conjugate pair of CO 2-3 and CO - 3. It has been shown that solutions containing levels of both partners in a conjugate pair (weak acid/bases only) are able to resist changes in p upon addition of acid or base. Such solutions are known as buffer solutions. Why do such solutions resist p changes? The equilibrium between the two species can shift and in doing so, will mop up excess 3 O + or O -. The equations below illustrate this effect. A + 2 O A- + 3 O + If a strong acid is added, the above equilibrium will shift to the left, with some of the conjugate base reacting with the excess hydronium. If a strong base is added, the equilibrium below will shift to the left, with the weak acid reacting with some of the added hydroxide. A O A + O - owever, there is a limit to the amount of acid or base that a volume of buffer solution can absorb without significant p change. This is known as the buffering capacity. A numerical definition of buffer capacity is the moles of strong acid or base that needs to be added to a litre of the buffer solution to cause a change of 1 p unit. Chemistry

11 Buffer capacity increases with higher concentrations of the components of the solutions, since there are more of the molecules to absorb the 3 O + or O -. It has also been found that equal concentrations of the conjugate partners are best for maximising capacity. As has been described, buffer solutions are a critical factor in maintaining our health, but it is not just the ability to resist p change, but also the actual value of the p itself that is important. Do all buffer solutions have the same p? If not, how can that p be determined before preparation of the solution? Not surprisingly, the p of a buffer solution is dependent on the identity of the components and their concentrations. Equation 3.11 below, known unfortunately as the enderson-asselbach equation, allows calculation of the theoretical p of a buffer solution. [base] p = pk a + log 10 Eqn 3.11 [acid] [acid] and [base] are the molarities of the conjugate pair when initially dissolved in the solution. EXAMPLES 1. Calculate the p of a buffer solution prepared by mixing 50 ml of 0.1 M N 3 and 150 ml of 0.2 M N 4 Cl. The initial molarities of each component of the solution must first be calculated, using C 1 V 1 = C 2 V 2. [N 3 ] = 50 x = M [N 4 + ] = 150 x = 0.15 M pk a for ammonium is Substituting these values into Eqn 3.11, p = log 10 ( ) = Using ethanoic acid and ethanoate buffer, what ratio of the two components is required to obtain a solution with a p of 5.0? The pk a of ethanoic acid is Substituting these values into Eqn 3.11, 5.0 = log 10 (base-acid ratio) = base-acid ratio = This means that the concentration of ethanoate must be 1.82 times greater than ethanoic acid. Buffer solutions find application in many areas of laboratory work. One use that you will be familiar with - though you may not have realised the nature of the solutions you were using - is in the calibration of p meters. Typically, three solutions of p 4.2, 7 and 9 are used to check that the meter and electrode are working properly. These solutions - colour coded so that they aren t mixed up - are, in fact, buffer solutions which will hold those p values unless something really drastic is done to them. Some titrations (not acid-base ones) require the use of a buffer solution to maintain a solution p suitable to allow proper reaction of the standard and analyte. An important example is the titration of calcium ions by EDTA - this requires a p 10 buffer, usually prepared from ammonia and ammonia chloride. Chemistry

12 PRACTICE QUESTIONS 29. What is the buffer capacity of a solution, 100 ml of which undergoes a 1 p unit change on the addition of 26.7 ml of 0.05 M Cl? 30. Show that the p of a buffer solution is not affected by the actual concentration of the two components, but by their ratio. 31. Show that the p of a solution where the components are of equal concentration is equal to the pk a. 32. Calculate the p of the following solutions. (a) 250 ml of 0.1 M ethanoic acid and 500 ml of 0.05 M sodium ethanoate (b) 1 L of 0.5 M N3 and 600 ml of 0.1 M N 4 Cl (c) 500 ml of M NaCO 3 and M Na2CO What mass of sodium ethanoate should be added to 250 ml of 0.05 M ethanoic acid to give a p of 4.5? Assume no volume change. WAT YOU NEED TO BE ABLE TO DO explain the theories of Arrhenius and Lowry-Brönsted define and give examples of the term amphoteric distinguish between strong and weak acids and bases calculate p for acid and base solutions classify solutions from their p describe aqueous acid-base equilibria write acid-base dissociation constant expressions perform calculations associated with acid-base equilibria outline the preparation of applications of buffer solutions explain how buffer solutions work perform calculations associated with buffer solutions Chemistry