Acids and Bases. 1. The solvent 2. The salt

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1 Acids and Bases A logical place to go from our study of hydrogen compounds is Acids and Bases. In 1884 Arrhenius devised a theory to explain acid/base behavior. His theory states that Acids contain protons and bases contain hydroxide ions. This neglects two issues that result in flaws in the theory: 1. The solvent 2. The salt This comes to light if we think about HCl in two different solvents. It behaves very differently in Water and Benzene. This implies the solvent has an affect on the chemistry. Salts seem to contradict this rule. Many including carbonates and phosphates are basic while, ammonium and aluminum ions are acid.

2 In come Bronsted and Lowry In come Bronsted and Lowry devise a theory that uses the concepts that acids donate protons and bases accept protons. In this theory the solvent becomes important as it can self-ionize and have acid base characteristics. AUTOIONIZATION H 2 O H 2 O H 3 O + OH -

3 Autoionization of water H 2 O H 2 O H 3 O + OH - In this process water does two things; it accepts a proton and it donates a proton. Hence, it behaves as both an acid and a base. It is said to be amphoteric or amphiprotic. This model relies on the existence of the H 3 O + ion. The first crystal structure of this of this ion was shown in 1924, one year after the development of the theory. It was in the structure of perchloric acid monohydrate.

4 The hydronium ion H 2 O H 2 O H 3 O + OH - + Although the hydronium ion is typically shown as H 3 O +, there are three waters of hydration H- bonded to it. What effect does this have on its stability?

5 The importance of solvent. Acid base behavior relies heavily on the solvent present. The strongest acid in a given solvent is the protonated solvent. This is illustrated well in the case of HF. HF (aq) + H 2 O H 3 O + (aq) + F - Here water is acting as a base and the fluoride ion is the conjugate base of HF. In a reaction with ammonia water is acting as an acid to give the conjugate acid the ammonium ion. H 3 O + (aq) + OH - (aq) 2H 2 O The reaction of a strong acid and strong base may be summarized as the reaction between the hydronium ion and the hydroxide ion. This reaction is viewed as going to completion because water dissociates to such a small degree.

6 A note about aqueous solutions In aqueous solution the strongest acid is the hydrodium ion and the strongest base is the hydroxide ion. What does this mean? What impact does this have on reactions? Generally, stronger acids and bases react upon addition to water to produce either the hydronium ion or the hydroxide ion. O - 2OH H 2 O (aq) HClO 4 + H 2 O H 3 O + (aq) + ClO 4 - (aq) BUT, are we as chemists limited to water? NO!!

7 Solvents other than water.. Any solvent with an ionizable hydrogen atom will work with a B-L acid-base system. How about liquid ammonia? It is rather cold though at -33 o C. NH 3 (l) + NH 3 (l) NH 4+ (NH 3 ) + NH 2- (NH 3 ) Here is an example of an acid base reaction in ammonia. NH 4 Cl(NH 3 ) + NaNH 2 (NH 3 ) NaCl + 2NH 3 Another COOL example is sodium metal in ammonia. Sodium acts as a base and the solution turns bright blue!

8 Equilibrium constants: Acids We understand and quantify acid strength in terms of how easy it is to transfer the H +. We use water as a reference point. The general dissociation process is outlined in the following reaction. HA (aq) + H 2 O (l) H 3 O + + A - (aq) The mathematical expression that describes this process is: k a = [H 3 O+ ][A ] [HA] To simplify things we use the value of pk a to describe the strength of an acid. The more negative the value the stronger the acid.

9 Equilibrium constants: Bases We understand and quantify base strength in terms of how easy it is to transfer the H +. We use water as a reference point. The general dissociation process is outlined in the following reaction. A - (aq) + H 2 O (l) OH - (aq) + HA (aq) The mathematical expression that describes this process is: k b = [HA][OH ] [A ] To simplify things we use the value of pk b to describe the strength of a base. Again the more negative the pk b the stronger the base.

10 Acid/Base Reactions: A Summary Bronstead-Lowry Lewis Acids Proton Donor Electron Acceptor Base Proton Acceptor Electron Donor Neutralization reaction involving HCl and NaOH HCl(aq) + NaOH(aq) NaCl(aq) + H 2 O(l) H + (aq) + Cl - (aq) + Na + (aq) + OH - (aq) Na + (aq) + Cl - (aq) + H 2 O(l) What is the NET ionic equation?

11 Relating K a, K b, and K w The autoionization of water can be related to the K a and K b of acids and their conjugate bases in aqueous solutions. k w = [HA][OH ] [A ] = [H 3 O][OH ] =10 14 Applying the power scale again [H 3 O+ ][A ] [HA] pk w = pk a + pk b =14

12 Trends in acid strength Strong acids K a > 1 (negative pk a values) Examples of these acids are: hydrochloric, nitric, sulfuric, perchloric acids. Weak acids K a < 1 (positive pk a values) Examples of these acids are: nitrous, hydrofluoric. The majority of other inorganic acids are weak. That means there is a significant amount of the molecular species existing in solution. Strong Acids Generally, all strong acids appear equally strong. This means that nearly 100% of the species is ionized in solution. This underscores the importance of the solvent as a protonvehicle (remember the hydronium ion). How would you evaluate the relative strength of acids?

13 Evaluating the strength of strong acids. To qualitatively identify the strength of an acid we dissolve stronger acids in a base weaker than water. What does this tell you about the acidic strength of the new solvent A good example is hydrofluoric acid. If we dissolve perchloric acid in HF, what happens? HClO 4(HF) + HF (l) H 2 F + + ClO 4 - (HF) As with water, the weaker acid (HF) acts as a proton acceptor (base) for perchloric acid. As a result of HF being a weaker base than water the equilibrium doesn t lie 100% to the right. We can do this for a series of acids and evaluate their strength. The strongest of common acids is Perchloric Acid.

14 What is a binary acid? Binary Acids Acids consisting of two simple ions (like hydrohalic acids) HX (aq) + H 2 O (l) H 3 O + + X - (aq) Acid pka Bond Energy (kj mole -1 ) HF HCl HBr HI Generally, the differences in strength can be related to the difference between the strength of the H-X and O-H bonds. The O-H bond energy is 459 kj mole -1. Any reaction tends toward the formation of a stronger bond. How can this be related to acid strength?

15 What is a oxyacid? Oxyacids Ternary acids (contain more than one atomic species) that contain oxygen. You MUST remember that for all common inorganic oxyacids the ionizable hydrogen is bound to an oxygen. Trends in strength. If we study oxyacids of one element we see stronger acids when there are more oxygen atoms.why? Look at nitric and nitrous acids pk a = -1.4 vs. pk a =3.3 How can YOU explain this? Look at the structures.

16 Semiquantitative prediction of Oxyacid strength. We can predict pk a values from the formula (HO) n XO m m pk a Note that ionizable H s are attached to an oxygen!! Can we understand the strength of perchloric acid looking at its structure?

17 What is a polyprotic acid? Polyprotic acids Acids which have more than one ionizable hydrogen. Note that it is increasingly difficult to remove successive protons, hence the strength decreases. Sulfuric acid is good example. H 2 SO 4 what s its structure? H 2 SO 4 (aq) + H 2 O (l) H 3 O + + HSO 4 - pk a = -2 HSO 4 - (aq) + H 2 O (l) H 3 O + + SO 4 2- pk a = 2 This brings an interesting point in regard to the crystallization salts of the +2 and +3 ions with metal cations.

18 Crystallization of metal sulfates. WHY? What crystallizes a +3 or +2 ion is present when crystal growth of the MHSO 4 is attempted? The SULFATE salt of the more positive cation is obtained. NOT THE HYDROGEN SULFATE! You know that the crystal structure is dependant upon lattice energies and that is a function of ionic charges. The lattice energy is greater for a crystal containing 2+ and 2- ions than for a crystal a single 2+ ion and two 1- ions. EXAMPLE The lattice energy for MgF 2 is 2.9MJ mole -1 but for MgO it is 3.9 MJ mole -1. This strongly suggests that the solid metal sulfate is favored by ~1MJ mole -1.

19 Let s look at the reactions. H 2 SO 4 (aq) + H 2 O (l) H 3 O + (aq) + HSO 4 - (aq) HSO 4 - (aq) + H 2 O (l) H 3 O + (aq) + SO 4 2- (aq) As we outlined the crystallization is favored by enthalpy. Hence, the sulfate salt is prepared with these ions. Can you apply Le Chatelier s Principle to this and understand the equilibrium? ONLY +1 METAL IONS FOR STABLE HYDROGEN SULFATE IONS. The salts of these ions result in acidic solutions HSO 4 - (aq) + H 2 O (l) H 3 O + (aq) + SO 4 2- (aq)

20 Sulfuric Acid O 2 O 2 / V 2 O 5 H 2 O S (s) SO 2 SO 3 H 2 SO 4 Contact Process H 2 SO 4 H 2 O H 2 S 2 O 7 Fuming Sulfuric Acid Sulfuric acid is the largest tonnage chemical. It is NOT a strong oxidizing agent but does act as a dehydrating agent very effectively.

21 Nitric Acid Haber Oswald O (+4) 2 H 2 O N 2(g) + 3H 2 2NH 3 NO (g) 3NO 2 2HNO 3 + NO O 2 /Pt (+5) (+2) Some useful notes: Bottles of HNO 3 have brown gas above the solution. Autoionization: hv 2HNO 3 2NO 2 + H 2 O + 1/2 O 2 2HNO 3 H 2 NO NO 3 - H 2 NO 3 + NO 2 + H 2 O This acid is a strong oxidizing agent because of the presence of the nitryl cation NO 2+. This cation is used in organic reactions for nitrating organics. Aqua-Regia is a mixture of 3HCl/1HNO 3 it contains free Cl 2 and ClNO 2. This is a powerful oxidizing environment and will dissolve Au and Pt.

22 Acidic Metal Ions Some metal ions are very acidic in aqueous solutions. Examples are Aluminum and Iron. Why do these metals do this? Both ions are relatively small and highly charged existing as hexahydrates in water. [Al(OH 2 ) 6 ] 3+ and [Fe(OH 2 ) 6 ] 3+ [Fe(OH 2 )] 3+ (aq) + H 2 O (l) H 3 O + (aq) + [Fe(OH 2 ) 5 (OH)] 2+ (aq) pk a = 3.3 Be sure you can do it for Al 3+!! M n+ (aq) + H 2 O (l) H 3 O + (aq) + M (n-1)+ (aq) pk a1 = Na + (14) Mg 2+ (12) Al 3+ (5) Fe 2+ (20) Fe 3+ (3) neutral weakly acidic acidic

23 Bronsted-Lowry Bases The most important B-L base is hydroxide. After this, it is ammonia. It reacts with water to give OH -. NH 3(aq) + H 2 O (l) NH 4 + (aq) + OH - (aq) Ammonia is useful as a glass cleaner as it reacts with fat molecules to make water soluble salts.

24 Other Bronsted-Lowry Bases Other common bases are the conjugate bases of weak acids. PO 3-4 (aq) + H 2 O (l) HPO 2-4 (aq) + OH - (aq) pk b = 1.35 S 2- (aq) + H 2 O (l) HS - (aq) + OH - (aq) pk b = 2.04 F - (aq) + H 2 O (l) HF (aq) + OH - (aq) pk b = Two of these are polyprotic acids what happens there?

25 Conjugate bases of polyprotic acids HS - (aq) + H 2 O (l) H 2 S (aq) + OH - (aq) Can you predict the second step for the sulfide ion? pk b = 6.96 Will its pk b be smaller, larger, or the same and why? HPO 4 2- (aq) + H 2 O (l) H 2 PO 4 - (aq) + OH - (aq) pk b = 6.79 This is an interesting case as the hydrogen phosphate ion behaves as a BASE! NOT AS AN ACID H 2 PO 4 - (aq) + H 2 O (l) H 3 PO 4 (aq) + OH - (aq) pk b = What does the pk b of the third process tell you?

26 Conjugate bases of strong acids. Generally, conjugate bases of strong acids do not interact with water. If in doubt think about a solution of sodium chloride.

27 Lewis Acids and Bases. Acid (LA) : electron pair acceptor Base (LB): electron pair donor Bronsted-Lowry acids and bases are a special case of this theory. HA (aq) + H 2 O (l) H 3 O + + A - (aq) Think about H +, OH - and NH 3. All metal ions are LA, and most ligands are LB. Molecules with incomplete octets are LA. B 2 H 6 + 2N(CH 3 ) 3 2H 3 B N(CH 3 ) 3 Or how about BF 3 and its reaction with NH 3.

28 Lewis Acids and Bases. Molecules or ions that expand their octet are LA. SiF 4 + 2F - SiF 2-6 ; why not CF 4? Si(sp 3 ) Si(sp 3 d 2 ) PF 5 + F - PF 6 - P(sp 3 ) P(sp 3 d 2 ) Molecules or ions with complete octets that can alter their valence electrons to accept electrons are LA. How about the reaction of CO 2 with OH -?