Redox Reactions: Crash Course Chemistry #10

The Importance of Electrons in Chemistry

This section introduces the significance of electrons in chemistry and their role in bonding atoms to form molecules. It also mentions redox reactions as the most important chemical reactions on Earth.

Electrons and Bonding

• Electrons are crucial for atoms to bond with other atoms and form molecules.

• The transfer of one or more electrons from one atom to another is involved in important reactions on Earth.

Redox Reactions

• Redox reactions, short for reduction-oxidation reactions, involve the transfer of electrons between atoms.

• Reduction refers to a substance gaining electrons, while oxidation involves losing electrons.

• These terms may seem inaccurate, but they have been widely used and are difficult to change.

Acid Base Reactions vs. Redox Reactions

This section explains that while acid-base reactions involve swapping protons, redox reactions focus on electron transfers.

Acid Base Reactions

• Acid-base reactions primarily involve swapping protons (H+ ions) between substances.

Redox Reactions

• Redox reactions are centered around the transfer of one or more electrons from one atom to another.

• Reduction refers to gaining electrons, while oxidation involves losing electrons.

Understanding Reduction and Oxidation

This section clarifies the meaning of reduction and oxidation in redox reactions.

Terrible Terminology Choices

• Reduction is when a substance gains electrons, contrary to its usual meaning of making something less.

• Oxidation does not always involve oxygen; it refers to a molecule losing electrons.

Historical Context

• Early chemists observed that substances became lighter during smelting due to oxygen gas leaving compounds.

Oxygen's Role

• Oxygen is an oxidizer that pulls electrons off other molecules to stabilize itself.

• When heated, oxygen becomes more energetic.

The Persistence of Reduction and Oxidation Terminology

This section explains why the terms reduction and oxidation persist in redox reactions despite their inaccuracies.

Attempts to Rename

• Some chemists have proposed alternative terms like electronation and de-electronation for reduction and oxidation.

• However, once a set of terms is established and widely used, it becomes challenging to change them.

OILRIG: A Mnemonic Device

• To remember the meaning of reduction and oxidation, the phrase "OILRIG" can be used:

• Oxidation is loss of electrons.

• Reduction is gain of electrons.

The Concept of Oxidation States

This section introduces the concept of oxidation states or oxidation numbers as a way to track electron transfers in redox reactions.

Proto Chemists' Observations

• Early chemists noticed that substances became lighter during smelting, leading them to describe this process as reduction.

Antoine Lavoisier's Discovery

• Antoine Lavoisier discovered that the weight loss during smelting was due to oxygen gas leaving compounds.

Understanding Electron Transfers

• Today, we understand that redox reactions involve electron transfers between atoms.

Importance of Tracking Electrons

• It is crucial to keep track of electrons in redox reactions, similar to keeping track of currency in transactions.

Electron Sharing in Covalent Bonds

This section explains how covalent compounds share electrons and introduces the concept of oxidation numbers for tracking electron ownership.

Covalent Bonding Analogy

• Covalent compounds can be compared to marriages where multiple individuals are involved.

Electron Sharing in Covalent Bonds

• Covalent compounds share electrons between atoms.

Tracking Electron Ownership

• To simplify the tracking of electron transfers, oxidation numbers are assigned to atoms.

• Oxidation number represents what an atom's charge would be if it exclusively owned all its electrons.

Rules for Assigning Oxidation Numbers

This section provides rules for assigning oxidation numbers to different elements in compounds.

Rule 1: Elements by Themselves

• The oxidation number of any element by itself (monatomic, diatomic, or polyatomic) is always 0 since atoms do not have a charge when they are not ions.

Rule 2: Monatomic Ions

• For monatomic ions, the oxidation number is equal to the size or number of its charge. For example, Fe2+ has an oxidation state of +2, while chloride ion (Cl-) has an oxidation state of -1.

Rule 3: Oxygen's Oxidation Number

• Oxygen typically has an oxidation number of -2 in most compounds except peroxide molecules like hydrogen peroxide (H2O2).

Rule 4: Hydrogen's Oxidation Number

• Hydrogen usually has an oxidation number of +1.

Rule 5: Halogens' Oxidation Number

• Fluorine and other halogens generally have an oxidation number of -1 unless bonded to fluorine or oxygen.

New Section

This section discusses the concept of oxidation states and how they are determined for polyatomic ions. It also introduces the idea that oxidation states can change depending on the elements they are bonded with.

Determining Oxidation States for Polyatomic Ions

• A polyatomic ion must have an oxidation state that matches its charge.

• The example of the sulfate ion (SO42-) is given, where sulfur's oxidation state is determined to be +6 based on the overall charge of -2 for the compound.

• The oxidation states of most elements can change depending on what they are bonded with.

New Section

This section explains why certain elements, like sulfur, do not have specific rules for determining their oxidation states due to their variability in different compounds.

Variability of Oxidation States

• The oxidation states of most elements can change depending on what they are bonded with.

• Sulfur is used as an example of an element that does not have a specific rule for determining its oxidation state because it can vary in different compounds.

New Section

This section introduces the concept of redox reactions and compares them to "molecular divorce courts" where electrons are exchanged between compounds.

Redox Reactions and Electron Exchange

• Redox reactions involve the exchange of electrons between compounds.

• The analogy of "molecular divorce courts" is used to describe how electrons change hands, being haggled over and traded, resulting in some compounds gaining or losing electrons.

• Some compounds benefit from these electron exchanges while others lose out.

New Section

This section presents an example of a redox reaction called the Haber process, which involves converting elemental nitrogen into ammonia.

The Haber Process

• The Haber process is a redox reaction that converts elemental nitrogen from the air into ammonia (NH3).

• Elemental nitrogen and hydrogen both have an oxidation state of zero initially.

• In the product, ammonia, nitrogen has an oxidation state of -3, indicating that it gained electrons and was reduced.

New Section

This section explains how to determine oxidation states in a redox reaction using the example of the Haber process.

Determining Oxidation States in Redox Reactions

• In the product, ammonia (NH3), each hydrogen atom has an oxidation state of +1.

• Since there are three hydrogens, nitrogen must have an oxidation state of -3 to balance out the overall charge of zero for the compound.

• Nitrogen gaining electrons and having its oxidation state decrease is considered reduction.

• Hydrogen losing electrons and having its oxidation state increase is considered oxidation.

New Section

This section discusses balancing redox equations by breaking them down into half-reactions.

Balancing Redox Equations with Half-Reactions

• Redox equations can be balanced by breaking them down into half-reactions.

• Although it may not be necessary for simple equations like the one in the example, balancing half-reactions can help with more complex equations.

• The example equation for the reduction half-reaction is N2 (oxidation state 0) becoming NH3 (oxidation state -3).

• The number of nitrogens and electrons are balanced first before combining with the oxidation half-reaction.

New Section

This section presents a more complicated example involving silver diamine to demonstrate why balancing redox equations with half-reactions can be necessary.

Balancing Complex Redox Equations

• The example involves the reaction between silver diamine and an organic aldehyde to produce a carboxylic acid, ammonia, and water.

• Oxidation states are assigned to each element in the reactants and products.

• The oxidation state of silver is +1 in the reactants and 0 in the products.

• The carbon in the aldehyde has an oxidation state of +1.

• Balancing half-reactions helps ensure that the number of electrons is equal on both sides of the equation.

New Section

This section concludes by summarizing the process of balancing redox equations using half-reactions.

Summary: Balancing Redox Equations with Half-Reactions

• Balancing redox equations can be achieved by breaking them down into half-reactions for reduction and oxidation.

• The number of electrons transferred must be balanced by adjusting coefficients in front of each half-reaction.

• Combining the balanced half-reactions gives a perfectly balanced redox equation.

New Section Understanding Oxidation States and Balancing Oxidation Reactions

In this episode of Crash Course Chemistry, the focus is on understanding oxidation states and balancing oxidation reactions. The transcript was written by Kim Krieger and myself, with script editing by Blake de Pastino, and chemistry consultation by Dr. Heiko Langner.

Assigning Oxidation States

• Assigning oxidation states is an important step in understanding chemical reactions.

• Oxidation state refers to the charge that an atom would have if all its bonds were 100% ionic.

• Rules are followed to assign oxidation states:

• Elements in their elemental form have an oxidation state of zero.

• Group 1 elements have an oxidation state of +1, while Group 2 elements have an oxidation state of +2.

• Oxygen usually has an oxidation state of -2, except in peroxides where it has a -1 oxidation state.

• Hydrogen usually has an oxidation state of +1, except when bonded to metals where it has a -1 oxidation state.

Balancing Oxidation Reactions

• Balancing redox reactions involves assigning oxidation states and ensuring that the number of electrons gained equals the number lost.

• Steps for balancing redox reactions:

• Identify the reactants and products involved in the reaction.

• Assign appropriate oxidation states to each element in the reaction.

• Determine which elements are being oxidized (losing electrons) and which are being reduced (gaining electrons).

• Balance the atoms other than hydrogen and oxygen first by adding coefficients as needed.

• Balance oxygen atoms by adding water molecules (H2O).

• Balance hydrogen atoms by adding hydrogen ions (H+).

• Balance the charge by adding electrons.

• Ensure that the number of electrons gained equals the number lost.

• Check that all elements and charges are balanced.

Examples

• Example 1: Balancing a simple oxidation reaction

• Given reaction: Mg + O2 -> MgO

• Assigning oxidation states: Mg(0), O(-2)

• Magnesium is oxidized from an oxidation state of zero to +2, while oxygen is reduced from an oxidation state of -2 to zero.

• Balancing the reaction:

• Mg + O2 -> MgO

• 2Mg + O2 -> 2MgO

• Example 2: Balancing a more complex oxidation reaction

• Given reaction: FeS2 + O2 -> Fe2O3 + SO2

• Assigning oxidation states: Fe(+3), S(+4), O(-2)

• Iron is oxidized from an oxidation state of zero to +3, sulfur is oxidized from an oxidation state of zero to +4, and oxygen is reduced from an oxidation state of -2 to zero.

• Balancing the reaction:

• FeS2 + O2 -> Fe2O3 + SO2

• 4FeS2 +11O->Fe8O12+8SO