What if you do not have a net ionic equation?
What if it is a complete molecular equation?

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If you don't have a net ionic equation to balance, that means you have what is usually called a molecular equation (or a complete molecular equation). These are harder to balance because the elements being oxidized and reduced are mixed in with substances that are are not being reduced or oxidized.

The solution is to, as much as possible, remove the "extraneous" substances: the ones that are neither oxidized or reduced. This gives you an unbalanced net ionic equation, which you proceed to balance. You then add back in the substances you had removed.

The problem? There is no explicit set of steps for how to do this. So, I will just let the technique evolve through some examples.

Comment: one thing that frustrates students learning the technique discussed in this tutorial is knowing how to decide what to eliminate and what not to. The easy answer to that is to advise you to eliminate everything but that which gets oxidized and reduced. For example, this equation:

CuCl2 + KI --> CuI + I2 + KCl

You must be able to see that the Cu is reduced from +2 to +1 and that the I goes from -1 to zero. That means that the chloride and the potassium ion are spectator ions and we arrive at this:

Cu2+ + I¯ ---> Cu+ + I2

and eventually this:

2Cu2+ + 2I¯ ---> 2Cu+ + I2
The problem then is that, to a large degree, knowing what to do to take it back to the complete molecular equation is experience-based. You have to see (and do) a number of examples before the ideas begin to set in.

That is why it must seem to a student that the teacher just magically knows what to do and the proper order of steps to follow. This is the experience factor. In the ChemTeam's case, high school chemistry was taken in the late 60's, well before many of you were born. Much experience has transpired since then!


Example #1: NaCl + H2SO4 + MnO2 ---> Na2SO4 + MnCl2 + H2O + Cl2

Solution:

1) What to remove?

NaCl: sodium ion is a classic spectator ion and it will be removed.
H2SO4; notice how only sulfate is present on the right-hand side. It will be removed.
MnO2: this stays since Mn is reduced in the reaction.
Cl2: it stays since the chloride from the left-hand side is oxidized.

Here is what remains: Cl¯ + H+ + MnO2 ---> Mn2+ + H2O + Cl2

That's the net ionic.

Notice I kept the MnO2 intact.

2) Now for the half-reactions:

Cl¯ ---> Cl2
MnO2 ---> Mn2+

3) Balance them:

2Cl¯ ---> Cl2 + 2e¯
2e¯ + 4H+ + MnO2 ---> Mn2+ + 2H2O

4) Since the electrons are already balanced, we add:

2Cl¯ + 4H+ + MnO2 ---> Cl2 + Mn2+ + 2H2O

5) We add in two more chloride to make MnCl2:

4Cl¯ + 4H+ + MnO2 ---> Cl2 + MnCl2 + 2H2O

6) We add in 4 sodium ion to make NaCl:

4NaCl + 4H+ + MnO2 ---> 4Na+ + Cl2 + MnCl2 + 2H2O

7) Add in four sulfate ions:

4NaCl + 2H2SO4 + MnO2 ---> 2Na2SO4 + Cl2 + MnCl2 + 2H2O

Example #2: SeCl2 ---> H2SeO3 + Se + HCl

Solution:

1) Here is what can be removed:

a) the SeCl2 can be changed to Se2+
b) the H2SeO3 can be changed to SeO32¯
c) the HCl.

I dropped the chlorine and hydrogen because neither was oxidized or reduced. The prime skill needed in changing a molecular equation to a net ionic is recognizing which elements have had their oxidation states changed.

I will discuss the SeO32¯ a bit more below. Also, the HCl tells me that this solution is to be balanced in acidic solution.

2) We now have this unbalanced net ionic equation:

Se2+ ---> SeO32¯ + Se

3) Here are the two half-reactions:

Se2+ ---> SeO32¯
Se2+ ---> Se

4) Balance the two half-reactions (in acidic solution):

3H2O + Se2+ ---> SeO32¯ + 6H+ + 2e¯
Se2+ + 2e¯ ---> Se

5) Add the two half-reactions and eliminate electrons:

3H2O + 2Se2+ ---> Se + SeO32¯ + 6H+

6) Restore items that were removed (shown in two steps). This gives the final answer:

a) 3H2O + 2SeCl2 ---> Se + SeO32¯ + 4HCl + 2H+
b) 3H2O + 2SeCl2 ---> Se + H2SeO3 + 4HCl

Why not drop the oxygen and get this half-reaction?

Se2+ ---> Se4+ + 2e¯

It does work because you get this balanced net ionic equation as an answer:

2Se2+ ---> Se4+ + Se

However, it seems to me to be a bit of a hassle to get back to the final answer. You might wish to try that on your own. Also, H2SeO3 is a weak acid. You might want to keep the full formula in its half-reaction and then balance. You can also get to the correct answer that way.


Example #3: H2SO3 + KMnO4 ---> MnSO4 + H2SO4 + K2SO4 + H2O

Solution:

You must be able to see that it is the Mn that gets reduced (from +7 to +2) and that the S gets oxidized (from +4 to +6).

1) Here is what I will remove:

a) the H2O (it will get back in via both half-reactions)
b) the H in H2SO3 and H2SO4
c) the potassium ion

Here is the unbalanced net-ionic equation that results:

SO32¯ + MnO4¯ ---> MnSO4

In order to make the half-reactions, I going to allow the MnSO4 to ionize. Why?

I know that the sulfate will play no role in balancing the reduction half-reaction, so I don't need it. How did I know? I know because the sulfur in the sulfate is being oxidized. I can't have an oxidation occuring in a reduction half-reaction. I will recombine the Mn2+ and the sulfate when I recover the balanced, molecular equation.

Also, suppose I keep the MnSO4 unionized. That means I would have to include the permanganate in the equation, to keep Mn present on both sides. That means I have to balance the net-ionic equation. Ah, no, I won't do that!

2) Here are the half-reactions that result:

SO32¯ ---> SO42¯
MnO4¯ ---> Mn2+

3) Here are the balanced half-reactions (note that it is acidic solution and that I have equalized the electrons):

5 [H2O + SO32¯ ---> SO42¯ + 2H+ + 2e¯]
2 [5e¯ + 8H+ + MnO4¯ ---> Mn2+ + 4H2O]

4) The final, balanced net ionic equation (extra water and hydrogen ion have been removed):

5SO32¯ + 2MnO4¯ + 6H+ ---> 5SO42¯ + 2Mn2+ + 3H2O

5) Make some MnSO4:

5SO32¯ + 2MnO4¯ + 6H+ ---> 3SO42¯ + 2MnSO4 + 3H2O

6) Make some potassium permanganate:

5SO32¯ + 2KMnO4 + 6H+ ---> 2SO42¯ + 2MnSO4 + 3H2O + K2SO4

Notice how I made some potassium sulfate at the same time.

7) Make some H2SO3:

2SO32¯ + 2KMnO4 + 3H2SO3 ---> 2SO42¯ + 2MnSO4 + 3H2O + K2SO4

8) Now, I will add 4 hydrogens on each side:

2KMnO4 + 5H2SO3 ---> 2H2SO4 + 2MnSO4 + 3H2O + K2SO4

Example #4a: HCl + KMnO4 ---> H2O + KCl + MnCl2 + Cl2

Example #4b: KMnO4 + KCl + H2SO4 ---> K2SO4 + MnSO4 + Cl2 + H2O

What follows is a discussion of 4a. Please attempt 4b on your own after study of 4a. A link to an answer for 4b is provided at the end of the discussion of 4a.

Solution:

We can eliminate the following:

H+ from the HCl
H2O
K+
Cl¯, but only from the right-hand side

Why only from the right-hand side? It is because the Cl¯ will be oxidized. I know this because there is Cl2 on the right-hand side as well. This tells me the chlorine is going from a -1 oxidation state to zero. It is being oxidized. So, I need Cl¯ on the left-hand side, but not the right-hand side.

Here is the net-ionic equation:

Cl¯ + MnO4¯ ---> Mn2+ + Cl2

From this, we get the half-reactions:

MnO4¯ ---> Mn2+
Cl¯ ---> Cl2

The balanced half-reactions:

2 [5e¯ + 8H+ + MnO4¯ ---> Mn2+ + 4H2O]
5 [2Cl¯ ---> Cl2 + 2e¯]

Note the addition of the factors 2 and 5, which will balance the electrons.

The balanced equation, in net-ionic form:

16H+ + 2MnO4¯ + 10Cl¯ ---> 2Mn2+ + 5Cl2 + 8H2O

Now, we must re-create the molecular equation, but in balanced form. First, make some HCl:

6H+ + 2MnO4¯ + 10HCl ---> 2Mn2+ + 5Cl2 + 8H2O

Next, add potassium ion back in:

6H+ + 2KMnO4 + 10HCl ---> 2Mn2+ + 2K+ + 5Cl2 + 8H2O

To both sides, add in six chlorine ions:

2KMnO4 + 16HCl ---> 2MnCl2 + 2KCl + 5Cl2 + 8H2O

Where did they go? On the left-hand side, they hooked up with the hydrogen ion to make 6 HCl, giving a total of 16HCl. On the right-hand side, two went with the potassium ion to make KCl and the other 4 made 2MnCl2.

I can add Cl¯ back in since we know, from the starting equation, that it was always present. Some of the Cl¯ reacted, to form Cl2, and some remained as spectator ions.

That last equation is your answer.

Link to solution for Example 4b


Example #5: Fe(s) + NO3¯(aq) ---> Fe3+(aq) + NO2(g)

Solution:

Notice that that reaction is already in net ionic form. I propose to balance it and then put it into molecular form.

Balance by half-reactions in acidic solution:

Fe(s) ---> Fe3+(aq) + 3e¯
e¯ + 2H+ + NO3¯(aq) ---> NO2(g) + H2O

Fe(s) ---> Fe3+(aq) + 3e¯
3e¯ + 6H+ + 3NO3¯(aq) ---> 3NO2(g) + 3H2O

6H+ + Fe(s) + 3NO3¯(aq) ---> Fe3+(aq) + 3NO2(g) + 3H2O

I knew to balance in acidic solution because of the nitrate. I needed positive ions coupled with the nitrate to give me a compound. Hydroxide did not fit my needs. Also, note that the iron(III) ion is in solution. In the presence of hydroxide, Fe(OH)3 would precipitate.

To make a molecular equation, join up some hydrogen ion with some nitrate ion:

3H+ + Fe(s) + 3HNO3(aq) ---> Fe3+(aq) + 3NO2(g) + 3H2O

We need more nitrate ion to combine with the rest of the hydrogen ion AND the ferric ion:

Fe(s) + 6HNO3(aq) ---> Fe(NO3)3(aq) + 3NO2(g) + 3H2O

It's now a fully-balanced molecular equation. Note that iron(III) nitrate is soluble, making the three nitrate ions added in this last step as the only spectator ions in the reaction.


Go to problems #1 - 10 Here's a list of the examples and problems (without any solutions)
Go to problems #11 - 25 Return to Redox menu
Go to problems #26 - 50