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Acid-Base Equilibria

Acid-Base Equilibria. Chapter 16. The presence of a common ion _____________the ionization of a ________acid or a __________base. CH 3 COONa ( s ) Na + ( aq ) + ). common ion. CH 3 COOH ( aq ) H + ( aq ) + ).

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Acid-Base Equilibria

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  1. Acid-Base Equilibria Chapter 16

  2. The presence of a common ion _____________the ionization of a ________acid or a __________base. CH3COONa (s) Na+(aq) +) common ion CH3COOH (aq) H+(aq) +) The __________________is the shift in equilibrium caused by the addition of a compound having an __________in common with the______________ substance. Consider mixture of CH3COONa (_______electrolyte) and CH3COOH (_______ acid). 16.2

  3. Consider an equal molar mixture of CH3COOH and CH3COONa • A ________________is a solution of: • A weak ______or a weak _______and • The _________of the weak acid or weak base • Both must be present! A ____________has the ability to _________changes in ___upon the addition of small amounts of either _______or __________. CH3COOH (aq) H+(aq) + CH3COO-(aq) Adding more _______creates a shift left IF enough _________ions are present 16.3

  4. Which of the following are buffer systems? (a) KF/HF (b) KCl/HCl, (c) Na2CO3/NaHCO3 (a) KF is a weak acid and F- is its conjugate base _______________ (b) HCl is a strong acid _______________ (c) CO32- is a weak base and HCO3- is it conjugate acid _______________ 16.3

  5. HCOOH (aq) H+(aq) + HCOO-(aq) Initial (M) Change (M) Equilibrium (M) What is the pH of a solution containing 0.30 M HCOOH and 0.52 M HCOOK? Mixture of weak acid and conjugate base! Ka for HCOOH = 1.8 x 10 -4 16.2

  6. HCOOH (aq) H+(aq) + HCOO-(aq) Initial (M) Change (M) Equilibrium (M) What is the pH of a solution containing 0.30 M HCOOH and 0.52 M HCOOK? Mixture of weak acid and conjugate base! Common ion effect 16.2

  7. HCl H+ + Cl- HCl + CH3COO- CH3COOH + Cl- 16.3

  8. Calculate the pH of the 0.30 M NH3/0.36 M NH4Cl buffer system. What is the pH after the addition of 20.0 mL of 0.050 M NaOH to 80.0 mL of the buffer solution? Initial Change End  16.3

  9. Calculate the pH of the 0.30 M NH3/0.36 M NH4Cl buffer system. What is the pH after the addition of 20.0 mL of 0.050 M NaOH to 80.0 mL of the buffer solution? final volume = 80.0 mL + 20.0 mL = 100 mL start (M) end (M) pH = 9.20 16.3

  10. [ [ [ Calculate the pH of the 0.30 M NH3/0.36 M NH4Cl buffer system. What is the pH after the addition of 20.0 mL of 0.050 M NaOH to 80.0 mL of the buffer solution? start (M) end (M) = 16.3

  11. Chemistry In Action: Maintaining the pH of Blood 16.3

  12. Titrations In a ____________a solution of accurately known _____________is added gradually added to another solution of __________concentration until the chemical reaction between the two solutions is complete. ___________________– the point at which the reaction is complete ___________– substance that changes color at the ________(hopefully close to the equivalence point) Slowly add base to unknown acid UNTIL The indicator changes color (________) 4.7

  13. NaOH (aq) + HCl (aq) H2O (l) + NaCl (aq) OH-(aq) + H+(aq) H2O (l) Strong Acid-Strong Base Titrations 100% ionization! ____________ 16.4

  14. CH3COOH (aq) + NaOH (aq) CH3COONa (aq) + H2O (l) CH3COOH (aq) + OH-(aq) CH3COO-(aq) + H2O (l) CH3COO-(aq) + H2O (l) OH-(aq) + CH3COOH (aq) _________Acid-________ Base Titrations At equivalence point (pH ______ 7): 16.4

  15. HCl (aq) + NH3(aq) NH4Cl (aq) H+(aq) + NH3(aq) NH4Cl (aq) NH4+(aq) + H2O (l) NH3(aq) + H+(aq) _________Acid-__________ Base Titrations At equivalence point (pH _____7): 16.4

  16. Acid-Base Indicators 16.5

  17. The titration curve of a strong acid with a strong base. 16.5

  18. Which indicator(s) would you use for a titration of HNO2 with KOH ? _______acid titrated with _________base. At equivalence point, will have __________________acid. At equivalence point, pH _____ 7 Use __________or__________________ 16.5

  19. Finding the Equivalence Point(calculation method) • ________Acid vs. ________Base • ________% ionized! pH ____No__________! • ________Acid vs. _______ Base • ______is____________; Need ____for conjugate ________equilibrium • _________Acid vs. ________Base • Base is________________; Need ____for conjugate ________equilibrium • _________ Acid vs. ______ Base • Depends on the ______of both; could be conjugate______, conjugate ______, or pH ___

  20. Exactly 100 mL of 0.10 M HNO2 are titrated with 100 mL of a 0.10 M NaOH solution. What is the pH at the equivalence point ? HNO2(aq) + OH-(aq) NO2-(aq) + H2O (l) Initial (M) Change (M) NO2-(aq) + H2O (l) OH-(aq) + HNO2(aq) Equilibrium (M) = Kb = start (moles) end (moles) Final volume = 200 mL =

  21. 2- Co2+(aq) + 4Cl-(aq) CoCl4(aq) 2+ Co(H2O)6 2- CoCl4 _____________Ion Equilibria and Solubility A __________is an ion containing a central metal cation bonded to one or more molecules or ions. 16.10

  22. 16.10

  23. Complex Ion Formation • These are usually formed from a transition metal surrounded by ligands (polar molecules or negative ions). • As a "rule of thumb" you place twice the number of ligands around an ion as the charge on the ion... example: the dark blue Cu(NH3)42+ (ammonia is used as a test for Cu2+ ions), and Ag(NH3)2+. • Memorize the common ligands.

  24. Common Ligands

  25. Names • Names: ligand first, then cation Examples: • tetraamminecopper(II) ion: Cu(NH3)42+ • diamminesilver(I) ion: Ag(NH3)2+. • tetrahydroxyzinc(II) ion: Zn(OH)4 2- • The charge is the sum of the parts (2+) + 4(-1)= -2.

  26. When Complexes Form • Aluminum also forms complex ions as do some post transitions metals. Ex: Al(H2O)63+ • Transitional metals, such as Iron, Zinc and Chromium, can form complex ions. • The odd complex ion, FeSCN2+, shows up once in a while • Acid-base reactions may change NH3 into NH4+ (or vice versa) which will alter its ability to act as a ligand. • Visually, a precipitate may go back into solution as a complex ion is formed. For example, Cu2+ + a little NH4OH will form the light blue precipitate, Cu(OH)2. With excess ammonia, the complex, Cu(NH3)42+, forms. • Keywords such as "excess" and "concentrated" of any solution may indicate complex ions. AgNO3 + HCl forms the white precipitate, AgCl. With excess, concentrated HCl, the complex ion, AgCl2-, forms and the solution clears.

  27. Coordination Number • Total number of bonds from the ligands to the metal atom. • Coordination numbers generally range between 2 and 12, with 4 (tetracoordinate) and 6 (hexacoordinate) being the most common.

  28. Some Coordination Complexes

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