R3.2 Electron Transfer Reactions (except 3.2.9-11) + S3.1.6 + R1

Unit H: Redox

Redox Reactions: The Balance of Change

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Introduction

Open Ended

Understanding Balance
What does it mean for something to be "in balance"? Can you think of examples in everyday life where balance is important?

Open Ended

Perspective on Change
How do you personally define gain and loss in your life? Does your perspective affect how you see changes?

Open Ended

Redox in Technology

Batteries power our phones and other devices. What might be happening at the chemical level to make that possible?

Open Ended

Perspective on Change

In a reaction where one substance loses something (like electrons) and another gains, how do we decide what’s being “lost” or “gained”?

Open Ended

Transformation and Duality

Oxidation and reduction are always linked—one can’t happen without the other. What does this teach us about the nature of interactions in the world?

Electron transfer, or redox reactions form one of the broadest and most important classes of reactions in chemistry. All reactions involving oxygen, such as combustion and corrosion, are redox reactions.

Redox definition in terms of :

-> Electron transfer
-> Oxygen/Hydrogen
-> Redox agents

Electron Transfer

​Electron Transfer

Example of electron transfer:


Na ⟶ Na⁺ +e^-

- sodium is oxidized to the +1 oxidation state.

 

Cl₂ + 2e⁻⟶ 2Cl⁻ 

- chlorine is reduced to the -1 oxidation state.

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Students use this mnemonic to help remembering which is which: LEO the lion says GER.







Another well-known mnemonic is this: OIL RIG

​Another way is to simply remember that reduction is to reduce the oxidation number. Therefore, oxidation must increase the value.

A reducing agent and the oxidizing agent that it becomes when it is oxidized form a redox couple, which we will represent with the reduced form of the couple followed by a diagonal line then the oxidized form as in Cu/Cu²⁺.

Oxygen/Hydrogen Transfer

​Oxygen/Hydrogen Transfer

The terms oxidation and reduction can be defined in terms of adding or removing oxygen or hydrogen to a compound.

​​Oxygen/Hydrogen Transfer

Notice that these are exactly the opposite

Redox agents

Redox agents

• Oxidizing agents add oxygen to another substance or remove hydrogen from it.

• Reducing agents remove oxygen from another substance or add hydrogen to it.

Dropdown

 

An ​ ​

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agent ​oxidizes something else.

Therefore, the ​ oxidizing agent must be ​

Select answer

.

ü  ​

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is gain of electrons (OIL RIG).

ü  Therefore, an oxidizing agent must ​

Select answer

electrons.

Rules to determine oxidation states

The oxidation state of an uncombined element is zero. This applies regardless of the structure of the element: Xe, Cl_2, S₈, and large structures of carbon or silicon each have an oxidation state of zero.

• The sum of the oxidation states of all the atoms or ions in a neutral compound is zero.

• The sum of the oxidation states of all the atoms in an ion is equal to the charge on the ion.

• The more electronegative element in a substance is assigned a negative oxidation state. The less electronegative element is assigned a positive oxidation state. Remember that electronegativity is greatest at the top-right of the periodic table and decreases toward the bottom-left.

• Some elements almost always have the same oxidation states in their compounds:

The Periodic Table of Oxidation States​
https://i0.wp.com/www.compoundchem.com/wp-content/uploads/2015/11/The-Periodic-Table-Of-Oxidation-States-2016.png?ssl=1

Exceptions

 

1. Hydrogen in the metal hydrides: Metal hydrides include compounds like sodium hydride, NaH. Here the hydrogen exists as a hydride ion, H^-. The oxidation state of a simple ion like hydride is equal to the charge on the ion—in this case, -1.
   
   
   2. Oxygen in peroxides: Peroxides include hydrogen peroxide, H₂O₂. This is an electrically neutral compound, so the sum of the oxidation states of the hydrogen and oxygen must be zero. Because each hydrogen has an oxidation state of +1, each oxygen must have an oxidation state of -1 to balance it.

Exceptions

 

3. Oxygen in F₂O: The deviation here stems from the fact that oxygen is less electronegative than fluorine; the fluorine takes priority with an oxidation state of -1. Because the compound is neutral, the oxygen has an oxidation state of +2.

4. Chlorine in compounds with fluorine or oxygen: Because chlorine adopts such a wide variety of oxidation states in these compounds, it is safer to simply remember that its oxidation state is not -1, and work the correct state out using fluorine or oxygen as a reference. An example of this situation is given below.

Drag and Drop

Assign oxidation states to all atoms of barium fluoride (BaF₂):
Ba: ​

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F:​

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Drag these tiles and drop them in the correct blank above

+2

-1

Drag and Drop

Assign oxidation states to all atoms of formaldehyde (Cu₂O):
Cu: ​

Placeholder

O:​

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Drag these tiles and drop them in the correct blank above

+1

-2

​Example:
Assign oxidation states to elements in KMnO4 (Potassium permanganate).

Solution Steps:
1. Use the rules for assigning oxidation states:
o Group 1 metals (e.g., K) always have an oxidation state of +1.
o Oxygen typically has an oxidation state of -2 (except in peroxides or with fluorine).
o The sum of oxidation states in a compound equals the overall charge.

2. Break down each compound:
o KMnO4 : K=+1, O=-2, K: 1 atom (+1) = +1 O: 4 atoms (-2) = -8

Then, solve for Mn such that the total = 0 (+1)+ (-8) + Mn = 0 Mn = +7

Oxidation number = (7+) / (1 atom) = 7+

​You Try


a)     Fe₂O₃ (Iron(III) oxide)

b)     H₂SO₄​ (Sulfuric acid) 

c)     What is the oxidation state of chromium in Cr²⁺

d)     What is the oxidation state of chromium in CrCl₃?

e)     What is the oxidation state of chromium in Cr(H₂O)₆³⁺

f)       What is the oxidation state of chromium in the dichromate ion, Cr₂O₇^2-?

 

​Oxidation states In naming compounds:

You will have come across names like iron(II) sulfate and iron(III) chloride. The (II) and (III) are the oxidation states of the iron in the two compounds: +2 and +3 respectively. That tells you that they contain Fe²⁺ and Fe³⁺ ions.

Oxidation Naming Practice: Workbook

Multiple Choice

Applying IUPAC rules, what is the name of Cu₂SO₄?

• copper(II) sulfate(VI)

• copper(I) sulfite(VI)

• copper(I) sulfate(VI)

• copper(II) sulfate(IV)

Multiple Choice

What is the oxidation state of chromium in the Cr₂O₇²⁻ ion?

• 6+

• 6

• VI

• +6

Multiple Choice

Which is a correct statement about reducing agents?

• The oxidation state of at least one element in the reducing agent must be able to decrease.

• They readily lose electrons.

• They only work in acid solution.

  
  

• They must contain at least one hydrogen atom.

Multiple Choice

Which substance cannot act as a reducing agent?

• SO₂

• H₂S

• SO₃

• S

Multiple Choice

In which compound does the metal ion have an oxidation state of +2?

• CO

• [Fe(CN)₆]³⁻

• Na₃VO₄

• [CuCl₄]²⁻

Multiple Choice

What is the sum of all the coefficients in the correctly balanced half-equation for the reduction of the dichromate(VI) ion in acidic solution?

Cr₂O₇²⁻(aq) + H⁺(aq) + e−  → Cr³⁺(aq) + _ H₂O(l)

• 30

• 23

• 24

• 29

Multiple Choice

In which reaction is iodine simultaneously oxidized and reduced?

• Br₂(aq) + 2KI(aq) → 2KBr(aq) + I₂(s)

• Ag⁺(aq) + I⁻(aq) → AgI(s)

• I₂(aq) + 2Na₂S₂O₃(aq) → Na₂S₄O₆(aq) + 2NaI(aq)

• IO₃⁻(aq) + 5I⁻(aq) + 6H⁺(aq) → 3H₂O(l) + 3I₂(s)

​Example: Given a molecular or empirical formula

Assign oxidation states to all atoms in
a) CH₃OH

According to the rules, hydrogen and oxygen have oxidation states of +1 and −2, respectively. Because methanol has no net charge, carbon must have an oxidation state of −2:

[(4 H atoms)(+1)] + [(1 O atom)(−2)] + [(1 C atom)(−2)] = 0


b)Fe₃O₄
Oxygen has an oxidation state of −2 , giving an overall charge of −8 per formula unit. This must be balanced by the positive charge on three iron atoms, giving an oxidation state of +8/3 for iron:

[(4 O atoms)(−2)]+[(3 Fe atoms)(+8/3)]= 0

c) CH₃CO₂H
1. Hydrogen (H) and Oxygen (O) Oxidation States:

H: +1, O: -2

Total charge for H and O: [(4 H) x +1] + [(2 O) x -2] = 0

2. Methyl Group (CH₃):

Carbon bonded to 3 H (+1 each) and 1 C.

C–C bond ignored.
Carbon oxidation state: -3

3. Carboxylic Acid Group (CO₂H):

C bonded to 2 O (-2 each) and 1 H (+1).

Net charge: [(2 O) x -2] + [(1 H) x +1] = -3
Carbon oxidation state: +3.

Final Oxidation States in CH₃CO₂H:

CH₃: C = -3, H = +1

CO₂H: C = +3, O = -2, H = +1

Total: (-3) + (3 x +1) + (+3) + (2 x -2) + (+1) = 0

Practice Problems:

a) 1. Assign oxidation states in:

CH₄, C2H₆, HCOOH, CO₂
C=−4, H=+1 : C=−3, H=+1: C=+2,H=+1,O=−2: C=+4, O=−2

b) Explain differences in oxidation states between CH₄ and CO₂.
The oxidation state of carbon changes from −4 to +4 in CO2​ due to differences in the electronegativity of the bonded atoms.

c) Find oxidation state of C in CH₃CHO (acetaldehyde).
CH3​: C=−3, H=+1

CHO: C=+1,H=+1,O=−2

Drag and Drop

Assign oxidation states to all atoms of ethanol CH₃CH₂OH:
C: ​

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H:​

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C:​

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O:​

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Drag these tiles and drop them in the correct blank above

-3

+1

+3

-2