# Solving Solubility Problem: AgNO3 + Na2SO4

• caragheen
In summary, the conversation revolved around predicting and testing chemical equations in class. The individual was stuck on one equation involving AgNO3 and Na2SO4 and was unsure why their predicted product, Ag2SO4, did not form a precipitate in the lab. It was determined that the solubility rules for sulfate salts are only an approximation and the actual solubility can vary. The reason for not getting a precipitate in the lab was due to the solution being unsaturated. It was also clarified that a supersaturated solution is when there are more ions dissolved than the solution can hold.
caragheen
In class, we've been working on predicting equations given particular chemicals, and later, actually testing some of them.
I've been doing just fine, but I'm stuck on one of them--

1. Homework Statement + attempted solutions

I was given: AgNO3 + Na2SO4, and was told to predict what the products would be.
Here is what my equation looks like--I know it's a double replacement reaction, and I matched up the ions, balanced them, and determined their solubility(using my rules):

2AgNO3(aq) + Na2SO4(aq) yields 2NaNO3 (aq) + Ag2SO4 (precipitate)

However, when I did the lab, the 2 solutions, when mixed together, did not form a precipitate/did not form a solid. Thus, this means that the products of this reaction should all be aqueous, so that Ag2SO4 (precipate) should actually be aqueous.
This does not seem to make sense to me though, as in my rules, it states that "sulfate salts are soluble/aqueous, but there are exceptions with Ag, Hg, Pb, Ca, Ba, and Sr"

I repeated the lab 3 more times, getting the same results.
Is there some sort of extra exception in this solubility rule? Why didn't I get a precipitate when my equation states that I should?

Hmm...
The concentration is something I'm not really sure of--we were told to mix 2 drops of each solution with each other.

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Part of what you wrote is right, part is either wrong or my English fails me.

In general classification of substances as soluble/insoluble is only an approximation. All substances are soluble in all solvents. However, they differ in solubility. Saturated solution of AgI is about 10-9M, saturated solution of calcium sulfate is about 10-2M - both are considered insoluble, but solubility of one is 107 times larger. There is no such thing as "low saturated solution" - solution either is saturated, or is not. If solution is not saturated, there is no precipitate (or, in other words, there is no solid in equilibrium with the solution).

--

ideasrule said:
Silver sulfate has a solubility of 1.2 g/100 mL: http://en.wikipedia.org/wiki/Silver_sulfate

It is actually 0.84 g/100mL water at 25oC. Wiki is wrong a bit. Your point is still valid.

Caraqheen, are you sure you used sodium sulfate and not ammonium sulfate?

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@ chemisttree--
Yes, I am positive that I used sodium sulfate.
Also, just to be on the safe side, I checked with one of my classmates today and we both had the same results--a predicted product side where one was a precipitate while in the actual lab, there was no precipitate evident.

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caragheen said:
Thank you for clearing up some of the parts that were confusing to me--I apologize for any inconveniences I gave you with my conclusion.

No need to apologize, as long as you are learning you have every right to be occasionally wrong

Thus, the silver sulfate turns out to be aqueous because the solution was not saturated.

Not sure if that's a correct way of stating things, or at least I have never seen it stated this way. I would not say "sulafte is aqueous", rather "sulfate was dissolved".

If the AgSO4 solution had been saturated, or more concentrated, it's number of ions would have enabled a precipitate to be formed.

I don't like the wording, but the idea behind is not incorrect.

And the reason why it was unsaturated in the first place was because the two original solutions (the silver nitrate and the sodium sulfate) weren't saturated?

No, it doesn't matter. You can mix two unsaturated solutions and observe precipitate, that's quite common. Solution is saturated for each salt (pair of ions) separately.

--

Okay--thanks for answering my side question.

Just as a note, if a solution has more ions dissolved than it can hold, wouldn't that be called a supersaturated solution?

If not 100% correct, at least I know I'm heading in the right path. :)

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Looks much better now

And yes, that would be a supersaturated solution.

--

Thank you so much for helping me, Borek!
My question has been answered--thanks to all who helped me here as well.

## 1. How do you determine the solubility of AgNO3 + Na2SO4?

The solubility of AgNO3 + Na2SO4 can be determined by conducting a solubility test. This involves mixing small amounts of the two substances together in a controlled environment and observing how much of the substances dissolve in the solution. The amount of substance that dissolves is the solubility limit.

## 2. What factors affect the solubility of AgNO3 + Na2SO4?

The solubility of AgNO3 + Na2SO4 can be affected by a variety of factors, including temperature, pressure, and the presence of other substances. Higher temperatures generally increase solubility, while higher pressures can either increase or decrease solubility depending on the specific substances involved. The presence of other substances can also affect solubility through processes such as complexation or ion pairing.

## 3. How does the solubility of AgNO3 + Na2SO4 change with pH?

The solubility of AgNO3 + Na2SO4 can be affected by pH through the formation of different ionic species. For example, at low pH levels, AgNO3 and Na2SO4 will dissociate into Ag+ and Na+ ions, which are highly soluble. However, at higher pH levels, these ions may combine with hydroxide ions to form less soluble species, resulting in a decrease in solubility.

## 4. How can you increase the solubility of AgNO3 + Na2SO4?

The solubility of AgNO3 + Na2SO4 can be increased by changing the temperature or pH of the solution, or by adding a substance that can complex with one of the ions present. For example, adding ammonia can increase the solubility of AgNO3 by forming a soluble complex with Ag+. Additionally, increasing the surface area of the substances through grinding or mixing can also increase solubility.

## 5. Why is it important to determine the solubility of AgNO3 + Na2SO4?

Determining the solubility of AgNO3 + Na2SO4 is important for several reasons. It can provide valuable information for industrial processes, such as in the production of silver or sodium sulfate. It can also help predict the behavior of the substances in various environments, such as in natural water systems. Additionally, understanding the solubility can aid in the development of new products and technologies, as well as in the study of chemical reactions and their rates.

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