Question

1.
Tube # 3: (a) It started out as purple, and turned into green.
(b) No precipitate was noticably formed.

Tube #4: (a) This tube was purple, and turned into white.
(b) A brown precipitate was formed after adding 1 mL of NaHSO3.

Tube #5: (a) This tube was also purple, but then changed to very light translucent purple.
(b) No prepcipitate was noticably formed.

2. (a) Green is MnO4-, manganate, which is manganese +6.
(b)Brown precipitate is MnO2, manganese dioxide, which is manganese +4, and water-white solutions are Mn2+, manganous.
(c) Purple is MnO4-, permanganate, which is manganese +7.
The HSO3- has sulfur in oxidation state of +4, therefore, it goes to +6 in all of the test tubes.

3. Write the net ionic equations for the redox reactions in each test tube???

I have no idea how to do net ionic equations. All I have are the balanced chemical equations:
(a)2 MnO4- + 2 OH- + SO32- --> 2 MnO42- + SO42- + H2O
(b) 2 MnO4- + 3 HSO3- --> 3 SO42- + H2O + 2 MnO2 + H+
(c) 2 MnO4- + 5 HSO3- + H+ --> 5 SO42- + 3 H2O + 2 Mn2+

Please Help!

Answer 1

a,b, and c are net ionic redox equations. You can't make them any more net ionic than they are.

Answer 2

THANK YOU!!