Redox Test Topics
Redox Test Topics - Chemistry Honors
Oxidation Numbers
Assigning oxidation numbers to elements in a chemical species (atom, ion, molecule).
Half-Reactions
Separating a redox reaction into two half-reactions:
Oxidation half-reaction: Shows the loss of electrons.
Reduction half-reaction: Shows the gain of electrons.
Oxidation and Reduction Processes
Identifying characteristics of oxidation and reduction:
Oxidation: Increase in oxidation number, loss of electrons.
Reduction: Decrease in oxidation number, gain of electrons.
Oxidizing and Reducing Agents
Labeling oxidizing and reducing agents in a reaction equation:
Oxidizing agent: The substance that causes oxidation (and is itself reduced).
Reducing agent: The substance that causes reduction (and is itself oxidized).
Redox Reaction Identification
Indicating whether a reaction is redox or not:
Redox reactions involve a change in oxidation numbers.
Balancing Redox Reactions (Acidic Solutions)
Balancing redox reactions in acidic solutions using the half-reaction method:
Separate the reaction into half-reactions (oxidation and reduction).
Balance atoms other than O and H in each half-reaction.
Balance O by adding H₂O.
Balance H by adding H⁺.
Balance charge by adding electrons (e⁻).
Make the number of electrons equal in both half-reactions by multiplying by appropriate coefficients.
Add the half-reactions together, canceling out electrons and any common species (H₂O, H⁺).
Electrolytic Cell
Explaining the redox process occurring in an electrolytic cell:
Electrolytic cell: Uses an external power source to drive a non-spontaneous redox reaction.
Voltaic/Galvanic Cell
Explaining the redox process occurring in a voltaic/galvanic cell:
Voltaic/Galvanic cell: Uses a spontaneous redox reaction to generate electrical energy.
Cell Diagram
Labeling a cell diagram, including:
Anode: Electrode where oxidation occurs (electrons are released).
Cathode: Electrode where reduction occurs (electrons are consumed).
Direction of electron flow: Electrons flow from the anode to the cathode.
Half-reactions occurring: Oxidation at the anode, reduction at the cathode.
Total cell reaction: Overall redox reaction occurring in the cell.
Ion flow through the salt bridge: Maintains charge balance in the half-cells.
Activity Series
Using the activity series to create a voltaic cell and identify the anode and cathode:
Activity series: A list of metals (or other substances) ranked in order of their ease of oxidation.
Voltaic Cell Sketch
Sketching a voltaic cell, labeling:
Cathode
Anode
Direction of electron flow
Batteries
Describing how a battery produces electrochemical energy:
Batteries use spontaneous redox reactions to generate electricity.
Cell Potential
Defining cell potential and describing how it is determined:
Cell potential (E_{cell}): The difference in potential between the cathode and anode, measured in volts (V).
E{cell} = E{cathode} - E_{anode}
It indicates the spontaneity of the redox reaction.
Standard Hydrogen Electrode (SHE)
Defining the Standard Hydrogen Electrode (SHE) potential of an electrode:
SHE: A reference electrode used to measure the standard electrode potential of other electrodes.
By convention, the standard electrode potential of SHE is defined as 0 V at 298 K (25°C) and 1 atm pressure.
Electrolytic vs. Voltaic Cells
Distinguishing between electrolytic and voltaic cells:
Feature
Electrolytic Cell
Voltaic Cell
Spontaneity
Non-spontaneous
Spontaneous
Energy Input
Requires external power source
Generates electrical energy
Redox Reaction
Driven by electricity
Drives electric current
Anode Charge
Positive
Negative
Cathode Charge
Negative
Positive
Uses of Electrolytic Cells
Listing some possible uses of an electrolytic cell:
Electroplating
Electrolysis of water
Production of metals (e.g., aluminum)
Oxidation Numbers
Assigning oxidation numbers to elements in a chemical species (atom, ion, molecule). Oxidation numbers are assigned based on a set of rules to help track the flow of electrons in a redox reaction. These numbers can be positive, negative, or zero, depending on the electronegativity of the atoms involved.
Half-Reactions
Separating a redox reaction into two half-reactions:
Oxidation half-reaction: Shows the loss of electrons. It represents the process where a species increases its oxidation number.
Reduction half-reaction: Shows the gain of electrons. It represents the process where a species decreases its oxidation number.
Oxidation and Reduction Processes
Identifying characteristics of oxidation and reduction:
Oxidation: Increase in oxidation number, loss of electrons. Occurs when an atom, ion, or molecule loses electrons, resulting in a more positive oxidation state.
Reduction: Decrease in oxidation number, gain of electrons. Occurs when an atom, ion, or molecule gains electrons, resulting in a more negative oxidation state.
Oxidizing and Reducing Agents
Labeling oxidizing and reducing agents in a reaction equation:
Oxidizing agent: The substance that causes oxidation (and is itself reduced). It accepts electrons from another species, leading to its own reduction.
Reducing agent: The substance that causes reduction (and is itself oxidized). It donates electrons to another species, leading to its own oxidation.
Redox Reaction Identification
Indicating whether a reaction is redox or not:
Redox reactions involve a change in oxidation numbers. If the oxidation number of any element changes during a reaction, it is a redox reaction.
Balancing Redox Reactions (Acidic Solutions)
Balancing redox reactions in acidic solutions using the half-reaction method:
Separate the reaction into half-reactions (oxidation and reduction).
Balance atoms other than O and H in each half-reaction.
Balance O by adding H₂O.
Balance H by adding H⁺.
Balance charge by adding electrons (e⁻).
Make the number of electrons equal in both half-reactions by multiplying by appropriate coefficients.
Add the half-reactions together, canceling out electrons and any common species (H₂O, H⁺).
Electrolytic Cell
Explaining the redox process occurring in an electrolytic cell:
Electrolytic cell: Uses an external power source to drive a non-spontaneous redox reaction. Electrical energy is converted into chemical energy.
Voltaic/Galvanic Cell
Explaining the redox process occurring in a voltaic/galvanic cell:
Voltaic/Galvanic cell: Uses a spontaneous redox reaction to generate electrical energy. Chemical energy is converted into electrical energy.
Cell Diagram
Labeling a cell diagram, including:
Anode: Electrode where oxidation occurs (electrons are released). It is the negative electrode in a voltaic cell and the positive electrode in an electrolytic cell.
Cathode: Electrode where reduction occurs (electrons are consumed). It is the positive electrode in a voltaic cell and the negative electrode in an electrolytic cell.
Direction of electron flow: Electrons flow from the anode to the cathode.
Half-reactions occurring: Oxidation at the anode, reduction at the cathode.
Total cell reaction: Overall redox reaction occurring in the cell.
Ion flow through the salt bridge: Maintains charge balance in the half-cells by allowing the migration of ions.
Activity Series
Using the activity series to create a voltaic cell and identify the anode and cathode:
Activity series: A list of metals (or other substances) ranked in order of their ease of oxidation. Metals higher in the series are more easily oxidized and will serve as the anode in a voltaic cell.
Voltaic Cell Sketch
Sketching a voltaic cell, labeling:
Cathode
Anode
Direction of electron flow
Batteries
Describing how a battery produces electrochemical energy:
Batteries use spontaneous redox reactions to generate electricity. They consist of one or more voltaic cells connected in series.
Cell Potential
Defining cell potential and describing how it is determined:
Cell potential (E arrcell): The difference in potential between the cathode and anode, measured in volts (V).
E{cell} = E{cathode} - E
arrcell*{anode}It indicates the spontaneity of the redox reaction. A positive E
arrcell*{cell} indicates a spontaneous reaction.
Standard Hydrogen Electrode (SHE)
Defining the Standard Hydrogen Electrode (SHE) potential of an electrode:
SHE: A reference electrode used to measure the standard electrode potential of other electrodes.
By convention, the standard electrode potential of SHE is defined as 0 V at 298 K (25°C) and 1 atm pressure.
Electrolytic vs. Voltaic Cells
Distinguishing between electrolytic and voltaic cells:
Feature | Electrolytic Cell | Voltaic Cell |
|---|---|---|
Spontaneity | Non-spontaneous | Spontaneous |
Energy Input | Requires external power source | Generates electrical energy |
Redox Reaction | Driven by electricity | Drives electric current |
Anode Charge | Positive | Negative |
Cathode Charge | Negative | Positive |
Listing some possible uses of an electrolytic cell:
Electroplating
Electrolysis of water
Production of metals (e.g., aluminum)