Understanding the Ionization of Salts

In summary: What would the AgI do?In summary, the conversation is about solubility calculations, specifically regarding the solubility of Pb(OH)2 at pH 10. The participants discuss the Ksp value, the concentration of OH- and Pb2+, and the use of ICE tables to solve the problem. They also discuss the effect of the common ion on solubility and how to calculate the remaining concentrations of Ag+ and I- after precipitation.
  • #1
Kqwert
160
3
Hello,

I am a bit confused re. solubility calculations.

Calculate the solubility of Pb(OH)2 at pH 10. Setting up the expression for Ksp:

Ksp = [Pb2+][OH-]2 = 8* 10^-17 (Ksp value from SI chemical data)

pOH = 14 - 10 = 4, i.e. [OH-] = 10-4.

8*10-17 = [Pb2+][OH-]2. Solving for [Pb2+] we get that [Pb2+] = 8*10-9.

My question is: why couldn't we use the relationship between the concentrations of Pb2+ and OH- like we do with ICE tables, to get the result? I mean, couldn't we just divide the concentration of OH- by 2 and get the concentration of Pb2+..?
 
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  • #2
Doesn't the bulk of the [OH] come from water ?
 
  • #3
Yeah, hadn't thought about that. Guess that might be the answer?
 
  • #4
Compare concentration of OH- already present pH 10 with the concentration of OH- that would be produced just by Pb(OH)2 dissolution. Does the addition of OH- from the dissolution matter?
 
  • #5
So, at pH 10 the [OH-] concentration will be 10-4. Whereas the concentration from *just* the Pb(OH)2 dissolution would be 5.42*10-6, right? (At neutral pH) But due the significantly higher concentration of [OH-] from the pH 10, we have to include that in the calculations?
 
  • #6
You can use the ICE table to show the process. It may help to better understand this problem by considering the difference between solubility of the Pb(OH)2 in, say DI water at pH = 7 and OH- treated water at pH = 10. Set up the ICE table and in the 'Initial' row enter '0' Molar for the Pb+2 concentration and 10-4 M for the OH- concentration as the common ion in the solution. (pOH = 14 - pH => pOH = 14 - 10 = 4 => [OH-] = 10-4) Fill in the rest of the table with variables per standard practice. Since the OH- is already present at 10-4 & additional OH- from further ionization of Pb(OH)2 is negligible, equilibrium row of the ICE table would be 'x' for [Pb+2 and 10-4 for [OH-. Apply to the Ksp expression for Pb(OH)2 and solve for x = [Pb+2] at pH = 10. Compare this with solubility of Pb(OH)2 in DI water ( for 1:2 ionization in DI water S = (Ksp/4)1/3. You should see a decrease in solubility on the order of 102.
 
Last edited:
  • #7
Kqwert said:
Hello,

I am a bit confused re. solubility calculations.

Calculate the solubility of Pb(OH)2 at pH 10. Setting up the expression for Ksp:

Ksp = [Pb2+][OH-]2 = 8* 10^-17 (Ksp value from SI chemical data)

pOH = 14 - 10 = 4, i.e. [OH-] = 10-4.

8*10-17 = [Pb2+][OH-]2. Solving for [Pb2+] we get that [Pb2+] = 8*10-9.

My question is: why couldn't we use the relationship between the concentrations of Pb2+ and OH- like we do with ICE tables, to get the result? I mean, couldn't we just divide the concentration of OH- by 2 and get the concentration of Pb2+..?

Check your Ksp value... I found Ksp to be different at => https://users.stlcc.edu/gkrishnan/ksptable.html
 
  • #8
Kqwert said:
So, at pH 10 the [OH-] concentration will be 10-4. Whereas the concentration from *just* the Pb(OH)2 dissolution would be 5.42*10-6, right? (At neutral pH) But due the significantly higher concentration of [OH-] from the pH 10, we have to include that in the calculations?

It is not like we to include it, we actually have to base calculations on it. Concentration of OH- already present is two orders of magnitude higher than what would be added from the Pb(OH)2 dissolution, so the effect of the dissolution can be rather safely ignored, while the effect of the ions present can not.
 
  • #9
You are exactly right ... I am just offering a 'perspective' approach to solving/visualizing the problem by 'assuming' the OH- ions are already present as a common ion and Pb(OH)2 is added into that solution. The high OH- concentration would of course inhibit ionization of the Pb(OH)2 as opposed to ionization in pH = 7 water. Calculation shows ~100x drop in solubility. at [OH-] ~ 10-4M levels. I often ask students to consider 'How much of the compound/salt would deliver ions into solution with vs without the common ion'. Of course, your right in that the ions are of primary consideration and the lower the concentration of ions delivered into solution the lower the solubility; i.e., as you say 'safely ignored'. Have a great day! :-)
 
  • #10
Thanks, both of you!

I have another question re. solubility:

You have a solution of 0.1M AgNO3 and 0.09M KI, and AgI is precipitated. What is [Ag+] and [I-] in the solution after precipitation?

Given the higher concentration of AgNO3 I understand there must be 0.1 - 0.09 = 0.01 M [Ag+] left in the solution. But what about the [I-? Isn't all this consumed in the precipitation of AgI?
 
  • #11
Please start a new thread for new questions.

Kqwert said:
Isn't all this consumed in the precipitation of AgI?

Imagine putting some solid AgI in the 0.01 M solution of Ag+. What would happen?
 
  • #12
Some [I-] would react and precipitate AgI whereas the rest would just stay as Ag+ and I- ions..?
 
  • #13
There were no I- ions in the case I have suggested - please reread what I wrote.
 
  • #14
Ok, but won't some of the added AgI dissolve in your case? That was what I meant by I-.
 
  • #15
Kqwert said:
Ok, but won't some of the added AgI dissolve in your case? That was what I meant by I-.

Yes, some would dissolve, but it would not precipitate any additional Ag+ - quite the opposite, when AgI dissolves concentration of Ag+ will (very slightly) increase. But you are right that some I- will show in the solution. Can you calculate (using Ksp) how much?
 
  • #16
Borek said:
Yes, some would dissolve, but it would not precipitate any additional Ag+ - quite the opposite, when AgI dissolves concentration of Ag+ will (very slightly) increase. But you are right that some I- will show in the solution. Can you calculate (using Ksp) how much?
Isn't (bold part) a contradiction? I know how I can calculate the amount of I-, knowing the Ksp of AgI and the concentration of [Ag+] = 0.01M, and the result is a veery small amount of [I-]. What I don't understand is where this [I-] comes from, because I thought all of it was consumed in the precipitation of AgI.
 
  • #17
Kqwert said:
all of it was consumed in the precipitation of AgI
If that were the case, what would [Ag]##\times## ##[##I] be ? And what is the litterature value ?
 
  • #18
You need to review the concept of salt ionization. All salts ionize to some degree. It might help to compare the solubility of the salt of interest (in this case AgI) in pure water vs in the presence of a common ion (0.01M Ag+ for example). You mentioned you knew how to determine ion concentrations. There is a significant drop in solubility of a given salt in the presence of a common ion but ionization to some degree still occurs. Remember not all salts have the same solubility but whatever does go into solution will ionize 100%.
 

1. What is ionization and how does it relate to salts?

Ionization is the process by which a compound or molecule gains or loses electrons, creating ions. Salts, which are ionic compounds, undergo ionization when they dissolve in water, breaking into their constituent ions.

2. Why do salts readily dissolve in water?

Salts have an ionic bond, which is a strong attraction between positively and negatively charged ions. Water molecules are polar, meaning they have a positive and negative end, making them able to attract and surround the ions, causing the salt to dissolve.

3. How does the concentration of salt affect its ionization?

The concentration of salt in a solution affects the degree of ionization. Higher concentrations of salt result in more ions being present, increasing the likelihood of ionization. Conversely, lower concentrations of salt result in fewer ions and therefore less ionization.

4. Can salts ionize in substances other than water?

Yes, salts can ionize in substances other than water. However, the degree of ionization may vary depending on the properties of the solvent. For example, salts may not fully ionize in nonpolar solvents, such as oil, due to the lack of polar interactions.

5. How does the ionization of salts affect their properties?

The ionization of salts affects their properties in several ways. It can alter their solubility, conductivity, and reactivity. For example, highly ionized salts will have a higher solubility and conductivity compared to salts with low ionization. Additionally, the presence of ions can also affect the chemical reactions in which the salt is involved.

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