Oxidation, Reduction, Redox reactions.

Question 1

Oxidation/ reduction in terms of electrons.

Answer 1

OIL, RIG
Oxidation is loss
Reduction is gain

Question 2

Ca + 1/2O2 –> CaO
What is being oxidised and what is being reduced in this reaction and why?
What is the reducing AGENT and what is the OXIDISING AGENT in this reaction?

Answer 2

• Ca goes from oxidation state of 0 to 2+, it loses electrons so, it is oxidised.
• 1/2O2 goes from oxidation state of 0 to 2-, it gains electrons so, it is reduced.
• Reducing agent: Ca
• Oxidising agent: O2

Question 3

What is a reduing/ oxidising agent?

Answer 3

• Reduing agent: electron donors, lose electrons - are oxidised themselves.
• Oxidising agent: electron acceptors. gain electrons - are reduced themselves.

Question 4

What oxidation state do elements (on their own) have?
ie. O2/ Ca/ Cl2

Answer 4

• 0.

Question 5

What oxidation state does H have - including the exception?

Answer 5

• +1.
• Exception: when the H is in hydrides ie. NaH (hydrides are when H is bonded to any metal), charge will be -1 here!

Question 6

True or False

Group 1 and group 2 elements in compounds will always have oxidation states of +1 and +2.

Answer 6

• True.
• Group 1 and group 2 elements in compounds will always have oxidation states of +1 and +2.

Question 7

What is a hydride?

Answer 7

• Hydride is H bonded to any metal (ie. NaH.)

Question 8

What is the oxidation state of chlorine, when it is in a compound? Including the exception.

Answer 8

-1
- Exception: except in compound with F and O, it would have a positive value.

Question 9

In ClF3, what is the oxidation state of the Cl and why?

Answer 9

• 3+.
• Fluorine = more electronegative than chlorine so takes negative charge rather than the chlorine.

Question 10

True or False

Fluorine’s charge/ oxidation state will change depending on what compound it is in.

Answer 10

• False.
• Fluorine’s charge/ oxidation state is always -1.

Question 11

What oxidation state does oxygen have in compounds? Including the exceptionSSS.

Answer 11

• 2-
• ExceptionSS : -1 in peroxides and +2 in OF2.

Question 12

Why does O have oxidation state of +2 in OF2?

Answer 12

• Because fluorine is more electronegative than oxygen, so takes the negative charge/ oxidation state rather than oxygen.

Question 13

What oxidation state does O have in H2O2?

Answer 13

-1.

Question 14

What oxidation state/ charge will aluminium ALWAYS have when it is in a compound?

Answer 14

+3

Question 15

What oxidation state does S have in H₂SO₄? Show workings.

Answer 15

• O (most electronegative) has oxidation state of 2-. 4 x 2- = -8.
• H = +1 charge. (2x +1) = +2.
• Need +6 for charges to balance so, S = oxidation state of +6.

Question 16

What oxidation state does S have in SO₄²⁻ ? Show working.

Answer 16

• O = 2- charge. (4 x 2-) = -8.
• Whole molecule is 2- so S must have charge/ oxidation state of +6.

Question 17

What must all the oxidation states in a compund add to?

Answer 17

• All the oxidation states in compound must add up to the overall charge on the molecule.

Question 18

What oxidation state does V have in VO²⁺?

Answer 18

• O has oxidation state/ charge of 2-.
• V must have oxidation state/ charge of 4+, to make the overall charge 2+.

Question 19

What are the steps for balancing half equations?

Answer 19

• Balance any atoms (other than O and H.)
• Balance any oxygens using H2O.
• Balance any hydrogens with H+ ions.
• Balance charge using electrons (e-)

Question 20

Write the half equation to show the conversion of MnO₄⁻ to Mn²⁺. Show all the steps.What does this equation show?

Answer 20

• Balance oxygens using water:
  MnO₄⁻ –> Mn²⁺ + 4H₂O.
• Balance any hydrogens with H+ ions.
  MnO₄⁻ + 8H⁺ –> Mn²⁺+ 4H₂O
• Balance the charge.
  MnO₄⁻ + 8H⁺ + 5e⁻–> Mn²⁺+ 4H₂O
  SHOWS REDUCTION (don’t think that it is oxidation because the charge goes from -1 to +2, it’s not comaparable because they are two completely different species.)

Question 21

These are two half equations in redox reaction.
Fe²⁺ –> Fe³⁺ + e⁻
MnO₄⁻ + 8H⁺ + 5e⁻ –> Mn²⁺ + 4H₂O
Write the overall ionic equation for this redox reaction/ steps to get to this stage?

Answer 21

• Make sure electrons are balanced.
• 5e- on bottom equation only 1 e- on top equation, so multiply top equation by 5.
• 5Fe²⁺ –> 5Fe³⁺ + 5e⁻
• Cancel out the electrons.
• Combine the two equations together:
  5Fe²⁺ + MnO₄⁻ + 8H⁺ –> 5Fe³⁺ + Mn²⁺ + 4H₂O

Question 22

What are redox reactions?

Answer 22

• When oxidation/ reduction are happening at the same time ie. in the same equation.

Question 23

True or False

Ionic equations can contain electrons, the electrons just have to be balanced on either side of the equation

Answer 23

• False!!
• Ionic equations DO NOT contain electrons.