Acid-Base Equilibria
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Transcript of Acid-Base Equilibria
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Acid-Base Equilibria
Chapter 16
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The __________________is the shift in equilibrium caused by the addition of a compound having an __________in common with the______________ substance.
The presence of a common ion _____________the ionization of a ________acid or a __________base.
Consider mixture of CH3COONa (_______electrolyte) and CH3COOH (_______ acid).
CH3COONa (s) Na+ (aq) +)
CH3COOH (aq) H+ (aq) +)
common ion
16.2
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A ________________is a solution of:
1. A weak ______or a weak _______and
2. 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 __________.
16.3
CH3COOH (aq) H+ (aq) + CH3COO- (aq)
Consider an equal molar mixture of CH3COOH and CH3COONa
Adding more _______creates a shift left IF enough _________ions are present
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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
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What is the pH of a solution containing 0.30 M HCOOH and 0.52 M HCOOK?
HCOOH (aq) H+ (aq) + HCOO- (aq)
Initial (M)
Change (M)
Equilibrium (M)
16.2
Mixture of weak acid and conjugate base!
KKaa for HCOOH = 1.8 x 10 for HCOOH = 1.8 x 10 -4-4
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What is the pH of a solution containing 0.30 M HCOOH and 0.52 M HCOOK?
HCOOH (aq) H+ (aq) + HCOO- (aq)
Initial (M)
Change (M)
Equilibrium (M)
Common ion effect
16.2
Mixture of weak acid and conjugate base!
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HCl H+ + Cl-
HCl + CH3COO- CH3COOH + Cl-
16.3
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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
End
16.3
Change
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pH = 9.20
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)
final volume = 80.0 mL + 20.0 mL = 100 mL
16.3
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=
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
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Chemistry In Action: Maintaining the pH of Blood
16.3
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TitrationsIn 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 baseto unknown acid
UNTIL
The indicatorchanges color
(________) 4.7
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Strong Acid-Strong Base Titrations
NaOH (aq) + HCl (aq) H2O (l) + NaCl (aq)
OH- (aq) + H+ (aq) H2O (l)
16.4
100% ionization!
____________
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_________Acid-________ Base Titrations
CH3COOH (aq) + NaOH (aq) CH3COONa (aq) + H2O (l)
CH3COOH (aq) + OH- (aq) CH3COO- (aq) + H2O (l)
CH3COO- (aq) + H2O (l) OH- (aq) + CH3COOH (aq)
At equivalence point (pH ______ 7):
16.4
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_________Acid-__________ Base Titrations
HCl (aq) + NH3 (aq) NH4Cl (aq)
NH4+ (aq) + H2O (l) NH3 (aq) + H+ (aq)
At equivalence point (pH _____7):
16.4
H+ (aq) + NH3 (aq) NH4Cl (aq)
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Acid-Base Indicators
16.5
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The titration curve of a strong acid with a strong base.
16.5
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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
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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 ___
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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)
start (moles)
end (moles)
NO2- (aq) + H2O (l) OH- (aq) + HNO2 (aq)
Initial (M)
Change (M)
Equilibrium (M)
Final volume = 200 mL
Kb = = =
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_____________Ion Equilibria and Solubility
A __________is an ion containing a central metal cation bonded to one or more molecules or ions.
Co2+ (aq) + 4Cl- (aq) CoCl4 (aq)2-
Co(H2O)62+ CoCl4
2-
16.10
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16.10
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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)4
2+ (ammonia is used as a test for Cu2+ ions), and Ag(NH3)2
+.• Memorize the common ligands.
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Common LigandsLigands Names used in the ion
H2O aqua
NH3 ammine
OH- hydroxy
Cl- chloro
Br- bromo
CN- cyano
SCN- thiocyanato (bonded through sulphur) isothiocyanato (bonded through nitrogen)
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Names• Names: ligand first, then cation
Examples:– tetraamminecopper(II) ion: Cu(NH3)4
2+
– 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.
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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)4
2+, 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.
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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.
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Some Coordination Complexes
molecular formula
Lewis base/ligand
Lewis acid donor atom
coordination number
Ag(NH3)2+ NH3 Ag+ N 2
[Zn(CN)4]2- CN- Zn2+ C 4
[Ni(CN)4]2- CN- Ni2+ C 4
[PtCl6] 2- Cl- Pt4+ Cl 6
[Ni(NH3)6]2+ NH3 Ni2+ N 6