Q:
I just did an online chemisty lab for Oxidation States of Manganese and I'm stuck. I have three test tubes that have been prepared as follows: Tube 1- Starting PH 13.82 2 mL of KMnO4 1 mL of NaOH Tube 2- Starting PH 7 2 mL of KMnO4 Tube 3- Starting PH 0 2 mL of KMnO4 1 mL of H2SO4 I then added five 1 mL increments of NaHSO3to each of the tubes and the reactions are as follows. Tube 1 - Started as green and turned light green as NaHSO3is added. PH 13.65, 13.65, 13.48, 13.13.41, 13.35 Tube 2 - Started as purple with a brown precipitate. Solution turns almost clear as NaHSO3 is added. PH 7.31, 7.38, 7.41, 7.38, 7.35 Tube 3 - Solution stayed purple and turned light purple with last addition of NaHSO3. PH 0,0,0,0,.14. What oxidation states did the MnO4- and HSO3- ions change in each test tube? Write the net ionic equations for the redox reaction in each test tube.
Already Tried: I'm not sure, but the green color in test tube #1 indicates MnO4-- which has the oxidation state of +6. The purple color of test tube #2 indicates MnO2 which has the oxidation state of +4. And the light purple color in test tube #3 indicates MnO2- which has the oxidation state of +2. Also, I tried coming up with the equation for each test tube, but probably didn't do them correctly: Test tube #1: KMnO4+NaOH+NaHSO3=KSO4+Na+MnO4+H2O Test tube #2: KMnO4+NaHSO3=KHSO3+NaMnO4 Test tube #3: KMnO4+H2SO4+NaHSO3= MnSO4+KSO4+NaSO4+H2O Concerning the oxidation states of H2SO4 in each test tube, I have no idea. Then comes the net ionic equation for each test tube. These I couldn't do because I wasn't sure that the equations were correct. Please help.
This may help.
MnO4^- has Mn in oxidation state of +7. It is purple.
MnO4^-2 has Mn in oxidation state of +6. It is green.
MnO2 has Mn in oxidation state of +4. It is a dark brown ppt.
Mn^+2 has Mn in oxidation state of +2. It is faintly pink.
The HSO3^- has sulfur in oxidation state of +4. It goes to +6 in all of the test tubes.
I found this on the following on the web.
http://dwb4.unl.edu/chemistry/smallscale/SS023.html
Then if I have the balanced equations of:
2 MnO4- + 2 OH- + SO32- --> 2 MnO42- + SO42- + H2O
2 MnO4- + 3 HSO3- --> 3 SO42- + H2O + 2 MnO2 + H+
2 MnO4- + 5 HSO3- + H+ --> 5 SO42- + 3 H2O + 2 Mn2+
HOw do I do the net ionic equations?
To determine the oxidation states of MnO4- and HSO3- ions in each test tube, we need to consider the changes in oxidation states as NaHSO3 is added.
In test tube 1:
- MnO4- is reduced to MnO42- (from +6 to +4)
- HSO3- is oxidized to SO42- (from +4 to +6)
The net ionic equation for the redox reaction in test tube 1 can be written as:
MnO4- + 4 HSO3- + 3 H2O → MnO42- + 4 SO42- + 8 H+
In test tube 2:
- MnO4- is reduced to MnO2 (from +6 to +4)
- HSO3- remains unchanged
The net ionic equation for the redox reaction in test tube 2 can be written as:
MnO4- + 2 H2O → MnO2 + 4 OH-
In test tube 3:
- MnO4- is reduced to MnO2- (from +6 to +2)
- HSO3- remains unchanged
The net ionic equation for the redox reaction in test tube 3 can be written as:
MnO4- + H2O → MnO2- + 4 H+
As for the oxidation state of H2SO4, it does not participate directly in the redox reaction and remains unchanged. Its oxidation state is +6.
To determine the oxidation states of MnO4- and HSO3- ions in each test tube, we need to consider the changes in oxidation states of manganese (Mn) and sulfur (S) in the reactions.
First, let's determine the oxidation states of MnO4- and HSO3- ions in their initial states:
1. MnO4-:
Since oxygen (O) is typically assigned an oxidation state of -2, we can determine the oxidation state of Mn (X) as follows:
-4 (from four O atoms) + X = -1 (overall charge of MnO4-)
X = +7
Therefore, Mn in MnO4- has an oxidation state of +7.
2. HSO3-:
Since oxygen (O) is typically assigned an oxidation state of -2, we can determine the oxidation state of S (X) as follows:
-2 (from two O atoms) + X + (-2) (from two H atoms) = -1 (overall charge of HSO3-)
X = +4
Therefore, S in HSO3- has an oxidation state of +4.
Now let's analyze the observations and changes in the reactions for each test tube:
1. Test tube 1:
The color change from green to light green suggests a reduction reaction. This indicates a decrease in the oxidation state of Mn.
The addition of NaHSO3 caused the pH to decrease gradually, indicating the evolution of acidic conditions.
Therefore, in test tube 1:
- MnO4- (initial oxidation state: +7) was reduced to a lower oxidation state (less positive) since the solution turned light green. Let's assume the oxidation state of Mn is +6 in this case.
The net ionic equation for test tube 1 may be:
MnO4- + H+ + e- → MnO4 2- (light green color)
2. Test tube 2:
The color change from purple to almost clear suggests a reduction reaction. This indicates a decrease in the oxidation state of Mn.
No change in pH or acidic conditions was mentioned.
Therefore, in test tube 2:
- MnO4- (initial oxidation state: +7) was reduced to a lower oxidation state (less positive) since the solution turned almost clear. Let's assume the oxidation state of Mn is +4 in this case.
The net ionic equation for test tube 2 may be:
MnO4- + e- → MnO4 2- (almost clear solution)
3. Test tube 3:
No color change was mentioned, suggesting no significant changes to the oxidation state of Mn.
The addition of NaHSO3 did not affect the pH, as it remained close to 0.
Therefore, in test tube 3:
- MnO4- (initial oxidation state: +7) remained in the same oxidation state since there was no color change. Let's assume the oxidation state of Mn is +7 in this case.
The net ionic equation for test tube 3 may be:
No net ionic equation can be formulated since the oxidation state of Mn did not change.
Regarding the oxidation states of H2SO4 in each test tube, you do not need to consider the oxidation state of H2SO4 because it acts as a reactant (does not change oxidation state) and is not involved in the redox reactions.
Please note that these net ionic equations are suggestions based on the provided information, and the actual equations may vary depending on specific experimental conditions and balanced equations. It is recommended to consult your lab manual or instructor for accurate net ionic equations.