Chapter 17 notes!
Chapter 17 notes! CHEM 130 - 003
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This 7 page Class Notes was uploaded by Izabella Nill Gomez on Monday October 26, 2015. The Class Notes belongs to CHEM 130 - 003 at University of Tennessee - Knoxville taught by Bin Zhao in Summer 2015. Since its upload, it has received 30 views. For similar materials see General Chemistry II in Chemistry at University of Tennessee - Knoxville.
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Date Created: 10/26/15
Chemistry Chapter 17 Notes C H3 COOH C H3 HOONa Consider both solutions contain 2 substances that share a common ion CH3C00quot aq zaqCH3C00 disociates completely CH3COONaaq gtNa quot o The weak electrolyte ionizes partially and the addition of CH3C00quot shifts the equilibrium reducing the number of H ions the presence of an added acetate ion causes the acetic acid to ionize less than it normally would Whenever a weak electrolyte and a strong electrolyte containing a common ion are together in a solution the weak electrolyte ionizes less than it would if it were alone in the solutioncommon ion effect onization of a weak base is also decreased by the addition of a common ion laq aqOH NH3aQH20ZNH Z Addition of NH shifts the equilibrium reducing the amount of OH ions Soutions with weak conjugate acidbase pairs resist drastic changes in pH when small amounts of strong acidbase are addedbuffered solutions buffers Ex human blood resists changes in pH because of acid to neutralize added OH ions and a base to neutralize H ions Acid and base that make up the buffer must not consume each other through a neutralization reaction The requirements ful lled by Z O a weak acidbase conjugate pair such as CH3C00H CH3CO 0quot or 8 H6 3 4 X6 3 2 Z NH3 NHf1 The buffered solution example fz X2 3 Z ZaqKaZ aqXz HXaqlt gtHquot pH is determined by the value of K0 and the ratio of concentrations of conjugate 4 H6 3 acidbase pairs Z If OH ions are added to the buffer they react X6 3 aq ZaqHXaq gtH20aqXquot OH X Z Causes HX to decrease and 26 to increase 3 aq gtHXaq If H ions are added WHY H6 4 Causes Xquot to decrease and HX to increase If the ratio of HXX is smathe pH is small 4 X6 3 X6 ogHpHlog K0 pKa pH z 3 Z lHXl pKa log H K 1 base HendersonHasselbach Equation p p a Ogacid used to calculate the pH of buffers Buffer capacity the amount of acid or base the buffer can neutralize before the pH begins to change to an appreciable degree Depends on the amount of acid and f CH3 COOH base used to prepare the buffer Ex pH of 1 M o and 1 M CHBCOONa is the same as a solution with the 1 M of each The second has greater buffering capacity because there39s more of the compounds pH range pH range over which a buffer acts effectively Buffers effectively resist a change in pH in either direction when the concentrations of weak acid and 19K conjugate base are the same When pH the buffer pH is optimal Buffers 19K usually have a useabe range within or 1 pH unit of pKa pH or 1 Reactions between strong acids and weak bases proceed essentially to completion as do those between strong bases and weak acids As long as buffer capacity is not exceeded the strong acidbase is completely consumed To calculate how the pH of the buffer responds to the addition of strong acidbase 1 Consider acidbase neutralization reaction determine the effect of HX and X stoichiometry calculations 2 Use values HX and X with the K0 to calculate H Equilibrium calculation use HendersonHasselbach Strong acid conjugate base of buffer LHX l lt ZHquot X6 gt i 3 Buffer with HX and Xquot Calculate new HX and 3 Xquot values use Kw to nd pH ZH20 gt T Z gtX T lt HX0Hquot Strong base conjugate acid of buffer Acidbase indicators can nd the equivalence point of titration point at which stoichiometrically equivalent quantities of acid and base have been brought together pH titration curve can monitor the progress of reaction The curve can determine equilibrium point and can also determine the Ka of a weak acid or Kb of a weak base 1 Strong acidstrong base titration 1 Initial pH pH of a solution before the addition of any base is through determination of concentration of the strong acid pH is low 2 Between initial pH and equilibrium point As NaOH base is added the pH increases slowly then rapidly pH levels are not yet neutralized 3 Equivalence point at the equivalence point there is an equal number of moles of base as acid reacted leaving only salt pH is 7 because cations and anions do not affect pH 4 After equivalence point pH is determined by the concentration of excess NaOH in the solution 2 Weak acidstrong base K 1 Initial pH Ex 50 mL of 1 M acetic acid 1M NaOH Use a pH289 ZH20 gt 2 Between CH3COOHCH3COOL use the neutralization reaction to nd CH 2 553C00H and CH3 OO Calculate the pH of the buffer 3 t ZH20 gt Solution with weak acid Zr X L calculate the new HX and Xuse Ka HX and gtHX0H X add strong base neutralization reaction HX decreases X increases to nd H 3Equivaence point is reached when 50 mL of the base is added to 50 mL of the acid 4 After equivalence point excess base pH is determined by the excess OH and pOH Weak acids have higher initial pH than strong acids pH change for rapid rise portion of the titration graph is smaller for weak acids than strong acids pH for equivalence point is above 7 for weak acid titration When weak acids have more than 1 ionizable H atom the reaction with OH occurs in steps aqlH20ll Zaq gtH2PO H3P03aqOHZ EX 2 ZaqH20l Zaq gtHPO ZaqOHZ H2P0 3 When neutralization steps of a polyprotic acid or polybasic base are suf ciently separated the titration has multiple equivalence points Phenophtaein is a good indicator for acidbase titrations Saturated solution one in which the solution is in contact with the undissolved solute 24 2 aq Bf 2ZaqSOZ 2Z BaSO4sltgtBaz 50 K Z SP Solubility product constant indicates how soluble the solid is in water and is denoted Ksp In general the solubility product Ksp of a compound equals the product of the concentration of the ions involved in the equilibrium each raised to the power of its coef cient in the equilibrium equation Coef cient for each ion in the equilibrium equation also equals its subscript in the compound s chemical formula Molar solubility is the number of moles of solute that dissolve in forming 1 L of saturated solution of the solute molL Soubiity of a slightly soluble salt is decreased by the presence of a second solute that furnishes a common ion pH of a solution affects the solubility of any substance whose anion is basic Soubiity of a compound containing a basic anion anion of a weak acid increases as the solution becomes more acidic Soubiity of a slightly soluble salt containing basic anions increases as H increases or as pH is lowered The more basic the anion the more solubility is in uenced by pH Solubility of salts with negligible basicity such as ClBr l and N0 is unaffected by pH changes Characteristic property of metal ions is their ability to act as Lewis acids toward H20 molecules which act as Lewis bases Bases other than water can also interact with metal ions such as transition metal ions Assembly of Lewis bases bonded to a metal ion is a complex ion Stability can be judged by size at K equilibrium Constant for formation from a hydrated metal ion f6 formation constant Soubiity of metal salts increases in the presence of Lewis bases such as quot LCUJ A inHV Amphoteric oxides and hydroxides soluble in strong acids and bases because they 2Z 2z5n 3 3Zan Z Crz Ag can behave as an acid or base Ex Z 2 Lv0Hquot react with acids 0 These dissolve because their anions The use of reaction quotient Q can be used to determine the direction in which a reaction must proceed to reach equilibrium lf Q gt Ksp precipitation occurs reducing the ion concentrations until 0 KW lf Qlt Kspa the solid dissolves increasing ion concentrations until 0 KW lf 0 Kspa equilibrium exists saturated solution lons can be separated from each other based on the solubilities of their salts selective precipitation
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