Acid-Base Equilibria

Chapter 16

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

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

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

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

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!

HCl H+ + Cl-

HCl + CH3COO- CH3COOH + Cl-

16.3

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

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

=

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

Chemistry In Action: Maintaining the pH of Blood

16.3

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

Strong Acid-Strong Base Titrations

NaOH (aq) + HCl (aq) H2O (l) + NaCl (aq)

OH- (aq) + H+ (aq) H2O (l)

16.4

100% ionization!

____________

_________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

_________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)

Acid-Base Indicators

16.5

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

16.5

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

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 ___

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 = = =

_____________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

16.10

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.

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)

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.

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.

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.

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