Transcript Complex Ions and Free Energy

Catalyst
1. Calculate the Ksp for AgC2H3O2 whose mass
solubility is 1.02 g/100 mL
2. A solution containing 0.010 M Ba2+ and 0.10 M
Ag+, which solid will precipitate first when
Na2SO4 is added to the solution? Justify your
reasoning with calculations. NOTE the Ksp
(BaSO4) = 1.1 x 10-10 and
Ksp (AgSO4) = 1.1 x 10-5.
End
Cobalt Complexes
Lecture 8.3 – Complex Ion Formation,
Equilibrium Constants, and ΔG
Today’s Learning Targets
• LT 8.6 – I can discuss how coordination complexes
form between metal ions and ligands. Furthermore,
I can determine the coordination number for a
coordination complex
• LT 8.7 – I can calculate the formation constant for
complex ions and relate that to the Ksp for a slightly
soluble compound.
• LT 8.8 – I can calculate the free energy of a chemical
reaction by utilizing my knowledge the value of K.
Transition Metal Complexes
• When transition metals bond they
are the central atom and they bond
extensively with numerous
substances.
• The complexes that form are known
as metal complexes
• The molecule that bond with metals
are known as ligands
▫ A ligand is always a Lewis Base
• The number of ligands bound to the
metal is the coordination
number
Common Ligands
• The most common ligands are H2O, NH3, Cl-,
and CN- due to their free lone pair of electrons
that can be donated to a metal.
The Metal Ligand Bond
• The bond between a metal and a ligand is a
Lewis acid-base interaction
• Metal = electron pair acceptor (acid)
• Ligand = electron pair donor (base)
Formation of Complex Ions
• The formation of complex ions can drastically
impact the solubility of a metal.
• Ligands act to pull metal ions into solution and
displace solvating H2O, thus, making the
solution more soluble
AgCl (s)
 Ag+ (aq) + Cl- (aq)
Ag+ (aq) + 2NH+3 (aq) 
 Ag(NH3 ) 2 (aq)
Overall: AgCl (s)+ 2NH+3 (aq) 
 Ag(NH3 ) 2  Cl  (aq)
Complex Ions
• The previously mentioned coordination complexes
or metal complexes are referred to here as
complex ions.
• For the equilibrium
Ag (aq)  2NH3 (aq)
 Ag(NH3 )2 (aq)
• The formation constant, Kf, is


3 2
[Ag(NH ) ]
Kf 

[Ag ][NH 3 ]
Class Example
• Calculate the concentration of Ag+ present in
solution at equilibrium when concentrated
ammonia is added to a 0.010 M solution of
AgNO3 to give an equilibrium concentration of
[NH3] = 0.20 M. Neglect the small volume
change that occurs when NH3 is added
Table Talk
• Calculate [Cr3+] in equilibrium with Cr(OH)4when 0.010 mol of Cr(NO3)3 is dissolved in 1 L of
solution buffered at pH = 10.0. The Kf for
Cr(OH)4- is 8 x 1029
Class Example
• The Ksp for AgI is 1 x 10-16 and Kf for Ag(CN)2- is
1 x 1021. Using these values, (a) calculate the molar
solubility of AgI in pure water and (b) calculate the
equilibrium constant for the reaction:


2

AgI(s)  2CN (aq)
 Ag(CN) (aq)  I (aq)
• and, (c) determine the molar solubility of AgI in a
0.100 M NaCN solution.
Table Talk
• The Kf of Ag(NH3)2 is 1.7 x 107 and the Ksp of AgCl (s)
is 1.6 x 10-10. Calculate (a) the molar solubility of AgCl
in pure water, (b) the equilibrium constant for the
reaction:

3

3 2

AgCl(s)  2NH (aq)
 Ag(NH ) (aq)  Cl (aq)
• and, (c) the solubility of AgCl (s) in 0.50 M NH3
Collaborative Poster
• Solve the two problems
that are on the handout
that you and your group
received.
• Explain all steps that you
do so that someone who
was not here could learn
how to do these problems
just from your poster.
Free Energy
• Recall that we calculated standard free energies
(ΔGo), which were when all concentrations were
1 M and temperature was 298 K.
• Most reactions occur at non-standard condition
and we can calculate ΔG at this point with:
G  G  RT lnQ
o
Free Energy and K
• At equilibrium ΔG=0 and Q = K. Therefore:
G  G  RTlnQ
o
0  G  RT lnK
o
G  RTlnK
o
G  RTlnK
o
Class Example
• The equilibrium for the Haber process at 25 oC is
N2 (g)  3H2 (g)
 2NH3 (g)
• What is the Kp for this reaction?

N2
H2
ΔHfo
(kJ/mol
0
0
NH3 -80.29
So (J/ (mol x K)
191.50
130.58
192.5
Table Talk
• Calculate the value of the equilibrium constant for the
following reaction:

H2 (g)  Br2 (l)
 2HBr(g)
H2
Br2
ΔHfo
(kJ/mol
0
0
HBr -36.23
So (J/ (mol x K)
130.58
174.9
198.48
Relay Races
1.If K is calculated to be a value of 0.5 for a reaction at 25 oC, then is the
reaction spontaneous?
2. If ΔG = 52 kJ/mol at 52 oC, then what is the value of K?
3. If K is determined to be 523 at 30 oC, then is the reaction
spontaneous at these conditions?
4. Use the following standard free energy of formation for formic acid
(HCO2). Calculate the Ka for the reaction:
HCO2
H+
ΔGfo (kJ/mol
-372.3
0
HCO2- -351.0
5. The Kf of Ag(NH3)2 is 1.7 x 107 and the Ksp of AgCl (s) is 1.6 x 10-10.
Calculate the equilibrium constant and ΔG for the reaction:

3

3 2

AgCl(s)  2NH (aq)
 Ag(NH ) (aq)  Cl (aq)
Closing Time
• Read 17.6, 19.7, and 20.2
• Homework: 17.60 17.61, 17.63, 17.64, 17.65,
19.79, 19.80, 19.81, 19.83