Humboldt State University ® Department of Chemistry

Richard A. Paselk

Chem 110

General Chemistry

Summer 2006

Lecture Notes::Lec 4_2 June

© R. Paselk 2006


Aqueous Ion Solubility, cont.

pH Effects on Ionic Solubility: pH can affect solubility dramatically if the anion is a weak acid salt. This is because reaction of the anion with hydrogen ion can convert it to another species and thus remove it from the equilibria. (Remember salts are strong!)

Example: What is the solubility of nickel sulfide in a solution of 0.1 M hydrogen sulfide at pH 6.0? Ksp = 1 x 10-19

Ksp = [Ni2+] [S2-]

and Ka = {[H+]2 [S2-]}/[H2S] = 1 x 10-22

So the concentration of sulfide ion, and therefore nickel ion, is dependent on pH.

From the pH given, [H+] = 1 x 10-6M.

Substituting and rearranging, [S2-] = ([H2S]/[H+]2)(1 x 10-22)

= 0.1(1 x 10-22)/(1 x 10-6)2 = 1 x 10-11

And solubility = [Ni2+] = (Ksp)/(1 x 10-11)

[Ni2+] = (1 x 10-19)(1 x 10-11)-1 = 1 x 10-8M

The effect of pH can be further seen by trying another calculation at pH 2.0:

[S2-] = 0.1(1 x 10-2)-2(1 x 10-22) = 1 x 10-19

[Ni2+] = (1 x 10-19)(1 x 10-19)-1 = 1 M


Complex Ion Equilibrium

Complex ions consist of metal ions surrounded by ligands.

A ligand is an electron pair donor (a Lewis base) that can bind to a metal ion via a weak covalent bond. Of course formation of a covalent bond requires not only an available electron pair on the ligand, but also an empty orbital on the metal ion where they can be shared. As a result complex ions most commonly involve transition metal ions with their available d- and p-orbitals. We will look at the bonding in complex ions in more depth later when we study the transition metals.

Note that for all of these metal ions the hydrated (aquo) ion is also a complex ion with water as the ligand.

The number of ligands around a complex ion is referred to as the coordination number. Some examples are listed below:

Complex Ion  Coordination Number  Kdiss



 5.9 x 10-8
 Cu(NH3)42+  4  1 x 10-12
 Cu(CN)42-  4  1 x 10-25
 Co(NH3)63+  6  6.3 x 10-36
 Fe(CN)63-  6  1 x 10-42

Each of these ions is formed via stepwise equilibrium processes, with the tabulated dissociation constant the products of the individual step constants.

We can look at these processes as formations and get Formation constants (or stability constants), which are just the inverse of the dissociation constants. For example, if one adds ammonia to a solution of silver ions, the silver ammonia complex ion is formed stepwise as shown below:

Ag(H2O)2+ + NH3 Ag(H2O)(NH3)+ K1 = 2.1 x 103
 Ag(H2O)(NH3)+ + NH3 Ag(NH3)2+ K2 = 8.2 x 103

Ag(H2O)2+ + 2NH3 Ag(NH3)2+  K = 1.7 x 107

Thermodynamics, Spontaneity, and Entropy

The First Law of Thermodynamics: This is simply a restatement of the Law of Conservation of Energy - Energy can neither be created nor destroyed. In other words, the energy of the Universe is constant. (Recall also that in a strict sense only the mass-energy of the Universe is conserved, since mass and energy are interconvertable by Einstein's famous equation, E=mc2. For our purposes now we can consider energy to be conserved as energy in chemical processes since the amount of mass-energy conversion is so slight. At the end of the semester when we look at nuclear chemistry we will need to take into account Einstein's equation.)

Even though energy can neither be created nor destroyed, the form of the energy can change.

For our purposes we need to consider three forms of energy:

  1. Kinetic Energy (KE) - the energy of particle motion: KE = 1/2mv2, where m is mass and v is velocity.
  2. Potential Energy (PE) - the energy due to the position of a particle.
  3. Radiation

Another term we must be aware of is Heat. Heat is a general catch-all for energy in transit. Notice that heat always moves spontaneously from hotter to cooler objects.

Recall that temperature is a measure of the average KE of the particles in a system. It can thus serve as an indicator of how readily the energy in a system can be transferred. However, temperature alone does not tell us how much energy is available in a system. We need to know not only the temperature, but also the heat capacity of the system to know how much energy it holds.

Demonstration: Boil water using a gas burner. Note that the energy (heat) released by the burner flame results from the breaking and making bonds between atoms. The energy differences in these bonds is released as KE (the molecules fly off at higher speeds) and radiation. This KE is then transferred to the screen and beaker etc.


Spontaneity and Entropy

The First Law tells us about the energy gained or lost during a process, but not whether that process will occur. So how do we predict which processes are likely to happen?

First we need to define the term spontaneous. In chemistry spontaneous means that a process will occur without outside influence in the direction written. In terms of equilibria we can say that a spontaneous reaction is one where the products are favored over the reactants (there is at least a slight excess).

{Note that saying a reaction is spontaneous says nothing about how fast the reaction or process occurs, just that it will occur! In other words the kinetics is not related to thermodynamics of a reaction. Thermodynamic functions are intrinsic to the overall reaction, and can't be changed without changing the reactants and products (under the same conditions) - they are pathway independent. Kinetics on the other hand can change drastically with path (such as with catalysts).}

So what determines whether a process is spontaneous? Two components contribute under common laboratory and biological conditions:

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© R A Paselk

Last modified 2 June 2006