pH, Brønsted-Lowry acidity, Lewis acids/bases. 2. Behavior of ions in aqueous solution This formalism alows the acid-base concept to be extended to.
25 pages
42 KB – 25 Pages
PAGE – 1 ============
GG325 L7 & 8, F2013 Lectures 7 and 8 Aqueous Inorganic Geochemistry of Natural Waters Pease read chapter White Ch6 (just 217 through page 2 4 Œthe second half of the chapter is for next week) 1.pH, Brønsted-Lowry acidity, Lewis acids/bases 2.Behavior of ions in aqueous solution Lecture 8 3. More about complexes and metal solubility 4. Quantifying aqueous solubility and total dissolved solids (TDS) Lecture 7 Acids and Bases GG325 L7 & 8, F2013
PAGE – 2 ============
Two Types of Acids and Bases: a. Brönstead acids and bases contain H +and OH -b. Lewis acids and bases Contain electron deficient ions (acids) and as electron ex cessive ions (bases) GG325 L7 & 8, F2013 GG325 L7 & 8, F2013 Two Types of Acids and Bases: a. Brönstead acids and bases , which contain H +and OH -b. Lewis acids and bases a different perspective on acidity/basicity involves the electron density around chemicals in aqueous solution. Brønstead acids can be thought of as electron deficient ions Brønstead bases can be thought of as electron excessive ions This formalism alows the acid-base concept to be exte nded to non protic, non hydroxyl, and in fact non-aqueous compounds.
PAGE – 3 ============
GG325 L7 & 8, F2013 a. Strong Brønsted acid and base examples: HC l H++ + Cl-(acid) NaOH Na ++ OH -(base). water is WEAK acid and base simultaneously, H2O H++ OH -. an “acid” neutral solution has equal amounts of H +and OH -, such H 2O without anything dissolved in it. Brønstead acidity: Kwis the equilibrium constant for the dissociation of water. H2O H++ OH -Kw= [H +][OH -] = 10 -14 at 25 °CKwhas a slight temperature dependence: Temp ( °C)K w010-14.94 less dissociated 2510 -14 60 10-13.02 more dissociated GG325 L7 & 8, F2013
PAGE – 4 ============
Brønstead acidity: Kw= [H +][OH -]An acid neutral solution always has [H +] = [OH -]. Setting [H +] = x, yields x 2= 10 -14 x = 10 -7 = moles of H +in this solution. pH = -log a H+ ~ -log [H +]pH = -log [10 -7 ] = 7 in an acid neutral at 25 °C A less common but sometimes useful variable is: pOH = -log a OH- ~ -log [OH -]In an acid neutral solution pOH = pH = 7 In any water at 25EC, pH + pOH = 14 GG325 L7 & 8, F2013 Non-acid neutral aqueous solutions “low” pH is <7 = acid solution = large H +"high" pH is >7 = base solution = small H +How much do natural waters deviate from acid neutrality? GG325 L7 & 8, F2013
PAGE – 5 ============
Lewis Acids and Bases Lewis acid example :BF 3has an empty orbital on B and only 6 electrons involved in the 3 B-F bonds. It is electron deficient so it is known as a Lewis Acid (B needs 2e -to achieve a Ne electronic configuration). BF 3will react with H 2O to get its needed electrons, creates new ions, and acidify the solution acid in a Brönstead sense. BF 3+ H 2O [OH-BF 3]-+ H+Lewis base example :In a similar fashion, NH 3(Ammonia) can be shown to be a lewis base, such that NH 3+ H 2O [H-NH 3]++ OH -note: NH 4+ is an ammonium ion BF 3NH 3GG325 L7 & 8, F2013 GG325 L7 & 8, F2013 The acidity of metal oxides Lewis Acid/Base and Bond character concepts help us understand why some elemental oxides form acids and some form bases in the hydrosphere. O2- ions are typically produced upon reaction with water; O 2- ions will react to take an H +from H 2O leaving two OH -in its wake (and therefore a basic solution). Acid Oxides These are covalent oxides. Ex ample: (SiO 2)When O is in a more covalent bond the “metal” shares more of thee -. To be stable in dissociated form in water, both elements in the bond try to get an electron by dissociating water and taking an OH -ion out of circulation. This leave an excess of H+(an acid solution). Basic Oxides These are ionic oxides. Example: Na 2O. In a compound where O has its e -held tightly to itself in a highly electronegative bond with a metallic element (little sharing),
PAGE – 6 ============
GG325 L7 & 8, F2013 Acidity from the aqueous CO 2system Equilibria involving carbon dioxide and its conjugate bases in water set the baseline pH of many waters at the Earth’s surface other chemical acids and bases usually serve to alter this carbonate equilibrium pH. Environmental chemists often speak of CO 2(aq ) = aqueous carbon dioxide in all of its aqueous forms. Recall from lecture 3 that these are: a. dissolved gaseous carbon dioxideCO 2(aq )b. carbonic acid H2CO 3(aq )c. bicarbonate anionHCO 3-(aq )d. carbonate anion CO 32- (aq )–and that Calcium Carbonate solubility also plays a role GG325 L7 & 8, F2013 Note that one of these 3 chemical species essentially dominates the mixture in each of three domains, which are separated by vertical lines at values of pH = pK a1 and pH = pKa2. These vertical lines are the locus of maxima in buffering capability of the solution. First let’s examine the pH relationships in a mixtur e of just water and CO 2components, meaning we ignore for now the effects of CaCO 3precipitation/dissolution. The Bjerrum plot shows dissolved CO 2, HCO 3-and CO 32- as a function of pH The curves are drawn using equations from last lecture.
PAGE – 8 ============
1GG325 L7 & 8, F2013 You can reason through the maximum and minimum buffering cases for CO 2in water yourself using: H2CO3: H++ HCO3-Ka1= [H+][HCO3-]/[H2CO3]pK a1= pH -log[HCO 3-][H2CO3]HCO3-: H++ CO32-Ka2= [H+][CO32-]/[HCO3-]pK a2= pH -log[CO 32-] [HCO3-]In general when we add “x” mol/L acid (H +) to a solution of H-Ac in water, the pH changes and [Ac-]/[HAc] changes to [Ac –x]/[HAc+ x]. When [Ac -]/[HAc] is close to 1 (i.e. at pH = pKa), the solution is less sensitive to the added x so the solution is buffered at pH = pKa .At [Ac-]/[HAc] <1 or > 1, the solution is not buffered. A plot of the pH sensitivity of an acetic acid solution as an external strong acid is added shows that near pHm= pKa, pH/molH+ added is at a minimum
PAGE – 9 ============
GG325 L7 & 8, F2013 pH Summary pH is a function of the ratio of conjugate base/acid. buffered: when acid and conjugate base are close to equal, log(conjugate base/acid) goes to 0 and pH=pK a.deviating from this point, each mole of H +added or subtracted goes into changing (conjugate base/acid). continue to change pH and approach an endpoint, this ratio changes more with each successive unit of pH change because almost all of the acid or conjugate base is consumed. oxides of elements can be thought of as Lewis acids and bases based on their bond character with O, and will make water acidic or basic accordingly GG325 L7 & 8, F2013 Alkalinity the acid-neutralizing capacity of an aqueous solution It is the sum of all the titratable bases in solution that can be neutralized with strong acid. Alkalinity is a significant environmental variable for natural wa ters and waste waters. The alkalinity of many natural waters is primarily a function of CO 2(aq ) and OH -. [Alk] [HCO 3-] + 2[CO 32-] + [OH -] -[H +]But the measured value can also include contributions from borates, phosphates, organic acid anions, silicates or other bases if present.
PAGE – 10 ============
Behavior of Ions in Water GG325 L7 & 8, F2013 GG325 L7 & 8, F2013 Behavior of ions in water: Aqueous stability of ions is the primary determinant of the “distribution” of m any elements between solids (minerals and organic matter) and water in surficial environments. The form that the ion takes in aqueous solution is the fundamental control on el ement solubility. Form is mostly a function of how the ion interacts with water molecules (as well as OH -, H 3O+and dissolved oxygen, aka fi DOxfl). These interactions are essentially dictated by Ion-O bonding characteristics, particularly in very fresh waters. During hydration (lewis acid/base interaction with H 2O molecules), Electronegativity and ion size determine the bonding preference of a cation for DOx or water (and its conjugate bases: OH -, O 2- )
PAGE – 11 ============
GG325 L7 & 8, F2013 Cation electronegativity determines how “ionic” or “covalent” the resulting O-cation “bonds” is: electropositive elements make ionic “bonds” whereas electronegative elements make relatively covalent “bonds”. Un-hydrated ion size affects O-cation ligand “bond character” and the stability of the hydration complex, since the geometric “fit” of electrons in “bonds” worsens as cation size increases. Cations that can form a stable covalent bond with O will do so in water. Those that don™t will make ionic bonds. Behavior of ions in water GG325 L7 & 8, F2013 The relative solubility of the cation-oxygen compound depends on the relative stability in water of the resulting oxy-or hydroxy ion versus a solid composed of the original cation and oxygen. For instance, the product would be favored in the reaction below for more covalent X-O bonds. Note that as H +is released the water becomes acidic. Behavior of ions in water Arrows depict the flow of electrons in breaking the H-O and forming the X-O bonds
42 KB – 25 Pages