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4. Acid Base Chemistry 4.1

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4. Acid Base Chemistry 4.1 4. Acid Base Chemistry 4.1. Terminology: 4.1.1. Bronsted / Lowry Acid: "An acid is a substance which can donate a hydrogen ion (H+) or a proton, while a base is a substance that accepts a proton. B- + HA <---> HB + A- base1 + acid1 <---> acid2 + base 2 Where HB = the conjugate acid A- = the conjugate base Examples: HCl + water; carbonate + water; H2S in Water Note that water can act as either an acid or a base. Such substances are called ampholytes, and are amphoteric. 4.1.2. Lewis Acid: This is a broader definition that is stated as a substance that accepts an electron, while a Lewis base donates a proton. All Bronsted/Lowry acids are Lewis acids, not all Lewis acids anr Bronsted acids. 4.1.3. Acids in Aqueous Solution: --Acids are usually refered to as donating protons, or H+, while bases donate OH-, or hydroxyls --The proton is strongly bound to water forming the basic unit of H3O+, the + Hydronium ion. This species in turn binds to other waters forming H9O4 --A similar structure is formed with OH, H7O4- + --the useage of H3O is usually reserved to reactions involving water. • The reaction for water: + - + - (H )(OH ) H2O <---> H + OH K = (H2O) + - Ka = (H )(OH ) = Kw (Dilute solution) ∆ ∆ -14 --We derived the value for Keq from G°f; G°r = +19.09 kcal/mol, Keq 25°C 10 4.1.4. The Dissociation (Equilibrium) Constant, Ka: • The Ka refers to the Keq for the reaction where an acid donates a proton to water; • The Kb is the reaction in which a base accepts a proton from a water molecule. • The value of Ka or Kb is the strength of the acid or base. Large numbers are strong acids, small numbers are weak acids --A strong acid has a stronger tendency to donate a proton --A strong base has a greater tendency to accept a proton. • Ka is usually expressed in log. pKa is the negative log of the equilibrium constant 0.8 • A strong acid has a pK a that is >0.8, a weak acid is <0.8. The Ka of HI is 10 • A strong base has a pK b of greater than 1.4 • A very strong acid dissociates completely at almost all concentrations • A very weak acid does not dissociate completely, leaving HA in solution • Borderline acids have HA present at high total A concentrations. • Rule of Thumb: Strong acids produce weak conjugate bases; weak acids produce - strong conjugate bases: e.g. HCl v. HCO3 4.1.5. Relationship between Kb and Ka: - - (HCN)(OH-) CN + H2O <---> HCN + OH Kb = (CN-)(H2O) + - + - (H3O )(OH ) HCN + H2O <----> H3O + CN Ka = (HCN)(H2O) By adding these two expressions, we get 2 H2O <---> H3O+ + OH-; (H3O+)(CN-)(HCN)(OH-) (H3O+)(OH-) K = KaKb = = (HCN)(H2O)(CN-)(H2O) (H2O) **This is a legal and useful approach. Remember, the K is an algebraic, numerical constant. Because the activity of water is 1: + - -14 KaKb = (H3O ) + (OH ) = Kw = 10 Where Kw is the equilibrium constant water. This relationship allows the calculation of either Ka or Kb if the other is known by: pH + pOH = pKw = 14 --therefore, in pure water when the (H+) = (OH-), the pH = 7, the pOH = 7 --Acid solution have a pH < 7, basic solutions have a pH > 7, or a pOH < 7 4.1.6. Neutrality, and the affect of temperature: -14 --Neutral solutions are pH 7 only at 25°C, where the Kw is 10 --At other temperatures, the pH changes because the equilibrium changes. As temperature increases, since there is a positive ∆H, dissociation increases. The ∆H° of the dissociation of water is +13.36 kcal/mole ∆H° 1 1 ln K @ 50°C = ln K + - = 5.73 x 10-14 w w, 25°C R 298 323 So at 50°C, neutrality is at pH = 6.62. 4.1.7 Muliprotic acids: • A monoprotic acid has one proton that can be donated under normal circumstances (i.e. environmental conditions. A diprotic acid has 2, a triprotic 3, etc.) • For groundwater systems, triprotic acids are the most complicated to deal with, and there are several of these, e.g. phosphoric acid, boric acid, and citric acid. • Each proton has a unique Ka, or acid dissociation constant, with names of Ka,1, Ka,2, etc. • A general rule for Bronsted acids is that the first proton to be donated from a multiprotic acid is much more acidic than the second, the second more acidic than the third, etc. In organic acids, the second and third may be similar. • The conjugate base of the first deprotonation of a diprotic acid is amphoteric, and is both a proton donor and a proton acceptor. 2 4.2. Equilibrium Calculations This is a systematic approach to solving acid-base equilibrium that uses 4 equations. The goal is to use enough equations, and the right equations, to solve the specific problem being examined. An important consideration, and one of the most difficult to master, is the use of specific a assumptions, used to simplify the equations. This a trial and error approach. 4.2.4. Mass Balance: In acid-base reactions the reacting species are conserved, assuming a closed system. When HA is added to water, the acid ionizes (partly or completely): HA + H2O ----> A- + H2O A mass balance equation would be: CT,A = [HA] + [A-] CT,A is the analytical concentration of all species containing A, with units being moles per liter. Note that concentration is used. When C moles of the salt MA are added to a liter of solution: MA <------> M+ + A- A- + H2O <---> HA + OH- CT,M = [M+] + [MA], or if MA dissociates completely, CT,M = [M+] 4.2.4. Equilibrium Relationships This utilizes the known equilibrium constant of the acids involved to help solve for concentration of an unknown species. The first equation is Kw. -14 o Kw = (H+)(OH-) = 10 at 25 C Note that here activity must be used, and the difference between activity and concentration used in the mass balance equation must be reckoned. The dissociation of HA is described by : ()()HA+− ()()HA OH− K = K = a ()HA b ()A− Since Kw = KaKb , both of these equations can not be used since they are not independent. 4.2.4. Charge Balance (electroneutrality) All solutions are electrically neutral. Unlike the analytical charge balance equation, the CB equation used here is an absolute equality – The total molar concentration of cations is always equal to the total anions. [M+] + 2[M2+] + [H+] = [A-] + + 2[A2-] + [OH-] NOTE: concentration is used, not activity. Also, note that [HA] is not used, as it is a neutral species. 3 4.2.4. Proton Condition: • The proton condition is a mass balance on protons, and is essential part of solving for protons involved in equilibria. • The proton mass balance is established with reference to a "zero level" or reference level for protons, called the PRL. The PRL is established as the species with which the solution was prepared, and species having protons in excess of the PRL are equated with species having less than the PRL. • This means that protons are conserved, and if protons are transfered from the initial reactants, that the result is some species gain protons (>PRL), while others lose protons (<PRL). 4.3. Strong Acid-Base Calculations. • Examples of strong acids in ground water are nitric and sulfuric acid from acid rain or industrial discharge, hydrochloric acid used to open up a well (dissolve calcite) • Strong acids have a large Ka, indicating that there are more products than reactants, i.e. that the reaction proceeds very far toward to products. • The dissociation of strong acids in water is exothermic and has a large positive entropy. The release of heat is from the hydration of proton by water to make hydronium ion. (H + )(Cl − ) HCl -Æ Cl- + H+ Ka = K = ~103, pK~-3 a (HCl) So the product of proton and chloride concentration is 1000 times the concentration of free HCl. This leads to the basic assumption used to simplify acid-base calculations: STRONG ACIDS DISSOCIATE COMPLETELY in dilute solutions. This results simplification of the mass balance calculation as: [Cl-]>>[HCl] CT,Cl = [HCL]+[Cl] Strong and intermediate acids important in natural systems: Acid Name Formula pK Perchloric HClO4 -7 Hydrogen Chloride HCl -3 Sulfuric H2SO4 -3 (pK1) Nitric HNO3 -1 + Hydronium H3O 0 - Bisulfate HSO4 1.9 (pK2) Phosphoric H3PO4 2.1 (pK1) 3+ Aquo ferric ion** Fe(H2O)6 2.2 (pK1) Arsenic Acid H3AsO4 2.25 (pK1) **The ferric iron ion in solution is coordinated to 6 waters of hydration, and you can think of these waters as having acidic protons. Another way of thinking of it is the exchange of a water for a hydroxyl ion – that decreases the hydroxyl activity, thus decreasing pH – i.e. an acid. So adding FeCl3 to water decreases pH. Aluminum acts in the same way. 4 4.4. Weak Acid-Base Calculations • Weak acids and bases do not dissociate completely, so while the approach to solving the equations is similar to strong-acid systems, the complication of the Ka is added. The use of simplifying assumptions is even more important for this system. Acid Name Formula pK Hydrofluoric HF 3.45 Acetic CH3COOH 4.7 - Arsenic Acid H2AsO4 6.77 Hydrogen Sulfide H2S 7.1 - Dihydrogen Phosphate H2PO4 7.2 (pK2) Hypochlorous Acid HOCl 7.6 Hydrogen Cyanide HCN 9.2 Boric Acid H3BO3 9.3 + Ammonium Ion NH4 9.3 - Bicarbonate Ion HCO3 10.3 - Arsenic Acid H2AsO4 11.6 (pK3) - Bisulfide HS 14 (pK2) Water H2O 14 • The small pK values mean that the concentration of the reactants is greater than the products, and the reaction does not proceed very far to the right.
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