4B4 Oxidation and Reduction

Question 1

Define:

oxidation number

Answer 1

The charge an atom would have if the compound was made up of ions.

Oxidation numbers indicate the degree of oxidation or reduction of an atom.

Question 2

What is the oxidation number of a free element?

Answer 2

Zero

Free elements like oxygen, nitrogen, and sulphur have an oxidation number of zero.

Question 3

Fill in the blank:

The sum of oxidation numbers in a neutral compound is _____.

Answer 3

Zero

The sum equals the net charge of the compound.

Question 4

Which rule determines the oxidation number of Group IA elements?

Answer 4

Group IA elements always have an oxidation number of +1.

This rule applies unless combined with a more electronegative element.

Question 5

List three uses of oxidation numbers.

Answer 5

1. Naming compounds.
2. Balancing equations.
3. Determining valency.

Oxidation numbers are critical in redox chemistry.

Question 6

How are oxidation numbers calculated in a binary ionic compound?

Answer 6

1. Identify the elements in the compound.
2. Identify their charge by looking at the periodic table.
3. Pay attention to the ratio of elements.
4. Solve for any unknown charges.

For example, the compound FeCl₃ is contains 3 chlorine, each with a charge of -1 for a total value of -3. Therefore, the iron (Fe) must equal to +3 since there is no overall charge on this compound.

Question 7

Calculate the oxidation number of chlorine in ClO₃⁻.

Answer 7

5

Cl + 3 (-2) = −1

Solving gives Cl = +5

Question 8

True or False:

The oxidation number of a monatomic ion equals the charge of the ion.

Answer 8

True

For example, Na⁺ has an oxidation number of +1 and S²⁻ has an oxidation number of -2.

Question 9

What is the oxidation number of fluorine in compounds?

Answer 9

-1

Fluorine is always -1 because it is the most electronegative element.

Question 10

Define:

redox reaction

Answer 10

A reaction involving the transfer of electrons where one species is oxidized and another is reduced.

Oxidation and reduction always occur together in redox reactions.

Question 11

What is an oxidizing agent?

Answer 11

A compound that easily gains electrons, decreases its oxidation state, and gets reduced.

Oxidizing agents accept electrons, becoming more negative, and are reduced in the process.

Question 12

What is a reducing agent?

Answer 12

A compound that easily loses electrons, increases its oxidation state, and gets oxidized.

Reducing agents donate electrons, becoming more positive, and are oxidized in the process.

Question 13

How can you identify an oxidizing agent in a reaction?

Answer 13

Look for a compound whose oxidation state decreases during the reaction.

Oxidizing agents lower their oxidation state as they gain electrons.

Question 14

True or False:

Hydrogen gas can act as both an oxidizing and a reducing agent.

Answer 14

True

The role of hydrogen gas depends on the context of the reaction.

Question 15

Which element is oxidized in the reaction Zn → Zn²⁺ + 2e⁻?

Answer 15

Zinc

(Zn)

Zinc’s oxidation number increases from 0 to +2.

Question 16

Which element is reduced in the reaction 2Fe³⁺ + 3Mg → 2Fe⁺ + 3Mg²⁺?

Answer 16

Iron

(Fe)

Fe³⁺ is reduced to Fe by gaining electrons; Its oxidation number decreases from 3 to o.

Question 17

Which element acts as the oxidizing agent in the reaction: Fe²⁺ + Se⁴⁺ → Fe³⁺ + Se³⁺?

Answer 17

Selenium

(Se⁴⁺)

Selenium’s oxidation state decreases from +4 to +3, making it the oxidizing agent.

Question 18

How can half-reactions be used to represent redox processes?

Answer 18

By isolating either oxidation or reduction processes in a reaction.

Half-reactions focus on either electron gain (reduction) or electron loss (oxidation) without showing the full reaction.

Question 19

List two ways to recognize a redox reaction.

Answer 19

1. Changes in oxidation states.
2. Simultaneous occurrence of reduction and oxidation.

Redox reactions are identified by analyzing the oxidation state changes of reactants and products.

Question 20

Fill in the blank:

During a displacement reaction, the element that is displaced is _______.

Answer 20

reduced

In displacement reactions, the displaced element gains electrons and is reduced.

Question 21

Fill in the blank:

Combustion reactions use _______ to oxidize organic compounds.

Answer 21

oxygen

Combustion reactions involve oxygen as the oxidizing agent to break down organic compounds.

Question 22

What is electrochemistry?

Answer 22

The study of the transformation of chemical energy into electrical energy and vice versa.

Electrochemistry involves redox reactions and is vital for technologies like batteries and fuel cells.

Question 23

Define:

an electrochemical cell

Answer 23

A device that converts chemical energy into electrical energy through a redox reaction.

Electrochemical cells are the basis of batteries and involve electron flow through a circuit.

Question 24

What are the two main parts of an electrochemical cell?

Answer 24

1. Electrodes
2. An electrolyte

Electrodes are conductors where reactions occur, and the electrolyte conducts ions.

Question 25

Which **electrode is positive** in an electrochemical cell?

Answer 25

The cathode. ## Footnote The cathode is where reduction occurs, and electrons are gained.

Question 26

# Fill in the blank: The \_\_\_\_\_\_ is the negative electrode **where oxidation occurs**.

Answer 26

anode ## Footnote The anode loses electrons and is the source of electron flow in the circuit.

Question 27

What is the **function of the electrolyte** in an electrochemical cell?

Answer 27

* To conduct ions. * To prevent charge buildup on the electrodes. ## Footnote The electrolyte completes the internal circuit by allowing ion flow between compartments.

Question 28

How are **electrons transferred** in an electrochemical cell?

Answer 28

From the **anode to the cathode** through a wire. ## Footnote Electrons flow from the anode (oxidation site) to the cathode (reduction site). Electrons flow due to the redox reaction, generating an electrical current.

Question 29

List *two* **materials** commonly **used as electrodes** in electrochemical cells.

Answer 29

1. Zinc (anode) 1. Copper (cathode) ## Footnote These metals are common due to their distinct electrochemical properties.

Question 30

Which component **prevents charge buildup** in the electrolyte solutions?

Answer 30

The salt bridge or porous barrier. ## Footnote It allows ions to move between compartments to balance the charge.

Question 31

# Fill in the blank: A **lemon** in a homemade battery acts as the \_\_\_\_\_\_\_\_.

Answer 31

electrolyte ## Footnote The lemon juice contains ions that conduct electricity.

Question 32

What are the **electrodes** in a homemade lemon battery?

Answer 32

A **copper penny** (cathode) and a **galvanized nail** (anode). ## Footnote These electrodes are connected by a wire to complete the circuit.

Question 33

What happens if the **electrodes** in a lemon battery **touch**?

Answer 33

The circuit **short-circuits**, and no useful current is produced. ## Footnote Proper separation is essential for the battery to function.

Question 34

How can you **increase the voltage** of a homemade lemon battery?

Answer 34

Connect **multiple lemons** in series. ## Footnote Connecting lemons in series adds their voltages together. Connecting cells in series combines their voltages to power devices requiring more energy.

Question 35

What are the *two* **types of electrochemical cells**?

Answer 35

1. Galvanic (Voltaic) cells 1. Electrolytic cells ## Footnote Galvanic cells generate electricity through a spontaneous redox reaction, while electrolytic cells use external energy to drive reactions.

Question 36

Which electrode is **negatively charged** in a Galvanic cell?

Answer 36

The anode. ## Footnote In Galvanic cells, the anode is negative because it releases electrons.

Question 37

Which electrode is **positively charged** in an electrolytic cell?

Answer 37

The anode. ## Footnote In electrolytic cells, the anode is positive because it attracts anions.

Question 38

What is **cell notation**, and how is it used?

Answer 38

A **shorthand to represent the reactions** in an electrochemical cell, with the anode on the left and cathode on the right. ## Footnote For example, Zn(s) | Zn²⁺ (aq) || Cu²⁺ (aq) | Cu(s) represents the half-reactions of Zinc and Copper.

Question 39

What is the relationship between **E°cell, E°cathode, and E°anode**?

Answer 39

E°cell = E°cathode − E°anode ## Footnote The electrochemical potential of the cell is the difference in reduction potentials of the cathode and anode. Reduction potentials can be found in standard tables, and the anode potential must be reversed for oxidation.

Question 40

How do the **charges** of the anode and cathode **differ in Galvanic and electrolytic** cells?

Answer 40

* Galvanic cells: The **anode is negative** and the cathode is positive. * Electrolytic cells: The **anode is positive** and the cathode is negative. ## Footnote The roles are reversed because of the nature of the energy source driving the reactions.

Question 41

What is the general equation for the **anode reaction**?

Answer 41

X(s) → X⁺(aq) + e⁻ ## Footnote This represents the loss of electrons during oxidation.

Question 42

What is the general equation for the **cathode reaction**?

Answer 42

Y⁺(aq) + e⁻ → Y(s) ## Footnote This represents the gain of electrons during reduction.