# Net ionic equation

1. Nov 7, 2004

### needhelpperson

What is the net ionic equation for the reaction of 0.5M solution of sodium hypochlorite and 0.5M solution of sodium thiosulfate? Calculate the enthalpy change.

I have a very terrible chemistry teacher. I haven't learned how to write ionic equations yet, but he wants this assignment for marks.

Besides that, enthalpy change, we've done somewhat of it. But we don't have the numbers for the bond energies. I don't even know the equation yet so i really don't know how to solve this at all. And is there a difference between the enthalpy change and standard heat of formation? I looked up a chart for Sodium hypochlorite, but it doesn't say the Hf of the it. I am really stuck here. Please help me with this whole problem.

2. Nov 8, 2004

### chem_tr

I think hypochlorite (bleach) will oxidize the sulfide sulfur in thiosulfate to give sulfate ion, along with the reduced form of hypochlorite, that is, chloride:

$$4Cl^+ +8e^- \longrightarrow 4Cl^-$$
$$S^{2-}\longrightarrow S^{6+}+8e^-$$

Now, as one sulfur is removed and one atom (oxygen) has to come there, we should put water and hydroxide to either sides.

$$4ClO^- + \underbrace{^-S-S[(=O)]_2-O^-}_{S_2O_3^{2-}} + 2OH^- \longrightarrow 4Cl^- + 2SO_4^{2-} + H_2O$$

About entropy change, you should count the molecules in both sides and apply the relevant formula (a 7-7 tie is present, so what the answer might be?).

3. Nov 8, 2004

### needhelpperson

Thanks a lot. But how could you tell that hypochlorite will oxidize the sulfide from thiosulfate? It's stuff like this that we don't learn. He just gives us the equations and we don't understand the stuff. I hate this teacher so much...
But thanks for your help though. I appreciate it.

As well i'm trying to find the heat of formation for hypochlorite and thiosulfate but i can't find it anywhere. Would you know what it is or where to find it? thanks

If it's not any problem, can you show a step by step approach to how you reached the equation? I'm trying to learn these stuff by myself and this would help a lot. Again, thanks alot.

Last edited: Nov 8, 2004
4. Nov 9, 2004

### chem_tr

1. Write the half-redox reactions, assuming that the sulfur is to be oxidized from sulfide (2-) to sulfate (6+); eight electrons are released in this half reaction. So, if hypochlorite provides only two, since Cl(+) to Cl(-) requires two electrons, multiply this with 4 to balance the electrons. Remember that no electron can be free in redox reactions.
2. Combine these two half-redox reactions to obtain a balanced reaction.
3. Convert them into more useful ones like ClO- for Cl+. For sulfide, you'll do the conversion as S2- to SO42-.
4. Now, why did I choose sulfide? Oxygen in sulfate never wants to be oxidized, and as sulfide is bigger, it has less control over outer shell electrons; so they can be easily modified.
5. Since hypochlorite is a good oxidizer, I presumed that a full oxidation scheme is to be followed; however, oxidation states lower than 6+ may also be possible; it may only involve sulfur, I mean, just remove two electrons from sulfide to obtain elemental sulfur, this is another alternative. Another stable oxidation state for sulfur is 4+ (SO32-)
6. If you write the reaction, you'll see that after converting the ions into their respective ones, there will be some shortages of small atoms; here, oxygen and hydrogen counts are unbalanced; we easily conclude that a water (or hydroxide) molecule should be added to the proper side to balance both electron and atom counts.

About Physical Chemistry-related issues, I must admit that this is my weak point. But you may try to find formation reactions of hypochlorite from chloride and thiosulfate from sulfate (or vice versa)

5. Nov 9, 2004

### needhelpperson

You're awesome at this stuff, can i add you to my msn, so that i can email some questions when i need help on them?

Last edited: Nov 9, 2004