chemistry ch 13
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CHEM1212Module13.docxName: __________ _________________________
CHEM 1212 Module 13 Reading Guide
Redox reactions
___________ is the branch of chemistry that studies the relationship between chemistry and electricity. Electrochemistry involves using either spontaneous redox reactions to create ___________ or using electricity to cause ___________ electron transfer reactions to occur. By definition, a redox reaction is one that entails changes in ___________ state for one or more of the elements involved. You can separate a redox reaction into two ___________ , one of which depicts the reduction and the other of which depicts the ___________ .
Oxidation = Increase in oxidation state = ___________ of electrons
Reduction = Decrease in oxidation state = ___________ of electrons
Reducing agent=species that is ___________ , provides the electrons to the other species
Oxidizing agent=species that is ___________ , takes on the electrons
The charge is balanced in each half-reaction using an appropriate number of ___________ . To combine these half-reactions into the overall reaction, the number of electrons gained and lost must be ___________ . ___________ is a reaction in which one reactant acts as both the oxidizing and ___________ agent.
Rules for assigning oxidation numbers
The oxidation number (or oxidation state) of an element in a compound is the ___________ its atoms would possess if the compound was ionic. The oxidation number of an element in a compound is essentially an assessment of how the electronic environment of its atoms is different in comparison to atoms of the pure ___________ .
The following guidelines are used to assign oxidation numbers to each element in a molecule or ion.
1. The oxidation number of an atom in an elemental substance is ___________ .
2. The oxidation number of a ___________ ion is equal to the ion’s charge.
3. Oxidation numbers for common nonmetals are usually assigned as follows:
· Hydrogen: +1 when combined with ___________ , −1 when combined with metals
· Oxygen: −2 in most compounds, sometimes −1 (peroxides, O22−), very rarely −1/2 (superoxides, O2−), positive values when combined with F (values vary)
· Halogens: −1 for F always, −1 for other halogens except when combined with oxygen or other halogens (positive oxidation numbers in these cases, varying values)
4. The sum of oxidation numbers for all atoms in a molecule or polyatomic ion equals the ___________ on the molecule or ion.
Practice Problems
2. For each of the following equations, identify the oxidizing agent, the reducing agent, the element oxidized, and the element reduced.
a. 4 H+(aq) + 2 NO3−(aq) + Cu(s) → Cu2+(aq) + 2 NO2(g) + 2 H2O(l)
b. Cu2O(s) + 2 H+(aq) → Cu(s) + Cu2+(aq) + H2O(l)
c. Zn(s) + CuCl2(s) → ZnCl2(s) + Cu(s)
Balancing Redox reactions
Equations representing redox reactions can be very challenging to balance by inspection, and the use of a systematic approach called the half-reaction method is helpful. This approach involves the following steps:
1. Write skeletal equations for the oxidation and reduction ___________ .
2. Balance each half-reaction for all elements except ____ and ____.
3. Balance each half-reaction for O by adding _____.
4. Balance each half-reaction for H by adding ____.
5. Balance charge in each half-reaction by adding ___________ to one side.
6. If necessary, ___________ one or both half-reactions so that the number of electrons consumed in one is equal to the number produced in the other.
7. Add the two half-reactions and simplify by removing species that appear on both sides of the equation.
8. For reactions occurring in basic media, carry out these additional steps:
a. Add _____ ions to both sides of the equation in numbers equal to the number of H+ ions.
b. On the side of the equation containing both H+ and OH− ions, combine these ions to yield ___________ molecules.
c. Simplify the equation by removing any redundant water molecules.
9. Check to see that both the number of atoms and the total charges are balanced.
Practice Problems
6. Balance the following oxidation–reduction equation in acidic solution
HAsO2(s) + Cl2(aq) → Cl−(aq) + H3AsO4(aq)
8. Balance the following oxidation–reduction equation in basic solution.
ClO−(aq) + Cr2+(aq) → Cr2O72−(aq) + Cl2(g)
Voltaic cells
A voltaic cell, also called a galvanic cell, uses a ___________ chemical reaction to create a flow of electrons that can be harnessed as electrical energy to do work.
Voltaic cells have the following important features:
1. Both oxidation and reduction are taking place at the same time, though not in the same ___________ .
2. The same ___________ reaction is taking place as if the two reactants were in direct contact.
3. Electrons move from the ___________ to the ___________ spontaneously.
In an electrochemical cell the half-reactions are split into separate compartments called half-cells. In each half-cell there are metal bars where the redox reactions take place called ___________ . Electrodes can be involved in the chemical reaction or they can be ___________ materials that conduct ___________ but remain unchanged, such as a platinum wire or a carbon rod. The electrodes are connected by a wire to carry the electrons. A negative electrode is the ___________ of electrons and is called the anode. The positive electrode is called the ___________ . Reduction always occurs at the ___________ in an electrochemical cell, and oxidation always occurs at the ___________ . The physical connection between half-cells that allows for movement of ions is called a ___________ and is necessary to complete the electrical circuit.
Standard cell notation summarizes the flow of electrons in a complete electrochemical circuit. This is a symbolic representation of electron flow in an electrochemical cell using a single vertical line to designate phase separation and a double vertical line to indicate a salt bridge. It shows the direction of electron movement, from left (anode) to right (cathode). Cell notation is written following these guidelines:
· The relevant components of each half-cell are represented by their chemical formulas or element symbols
· All interfaces between component phases are represented by vertical lines; if two or more components are present in the same phase, their formulas are separated by commas
· By convention, the schematic begins with the ___________ and proceeds left-to-right identifying phases and interfaces encountered within the cell, ending with the cathode
· Half-cells are separated by a double vertical line.
Practice Problems
21. Consider the voltaic cell represented by the following standard cell notation, where X and Y are generic metals:
X(s) ∣ X2+(aq) ∥ Y+(aq) ∣ Y(s)
a. Write the net ionic equation for the overall reaction.
b. Sketch the voltaic cell. Indicate the direction of electron flow, positive ion flow, and negative ion flow.
c. Which electrode gains mass?
d. Which electrode loses mass?
Cell Potentials
The driving force for the flow of electrons is called the electrical ___________ , E, or electromotive force, emf. This is measured in volts. The standard hydrogen electrode is a half-cell composed of hydrogen gas and a platinum electrode in 1 M hydrogen ion solution that is assigned a voltage of ___________ ; it is used to measure the standard reduction potentials of other half-reactions as relative values. Each half-reaction has its own standard reduction potential, E°, measured at the standard conditions of 1.0 M solutions, 1.0 atm gas pressure, and ___________ liquids and solids if they are present. Reactants in half-reactions that have large positive E values will be good ___________ agents. Once E° values are known for various half-cell reactions, any two of them can be combined to calculate the potential of a cell formed by combining those two half-cells.
___________ cell potentials are associated with a spontaneous redox reaction. Negative cell potentials indicate that the reverse direction is spontaneous. E can be related to delta G using the number of ___________ and a constant that accounts for charge of electrons:
where F = 96485 C/mol.
Practice Problems
Calculate the standard potential of each of the following cell:
Cu(s) + Pb2+(aq) → Cu2+(aq) + Pb(s)
Calculate the standard free-energy change of the reaction at 25°C.
Calculate the equilibrium constant.
Text segments from OpenStax, Chemistry. OpenStax CNX. Sep 15, 2020 http://cnx.org/contents/[email protected]. Section URL: https://openstax.org/books/chemistry-atoms-first-2e/pages/16-1-review-of-redox-chemistry ; White, J. et al. Interactive General Chemistry. MacMillan, NY. 2019.; Chemistry LibreTexts, Libretexts, [online] Available from: https://chem.libretexts.org/Courses/Heartland_Community_College/HCC:_Chem_161/5:_Thermochemistry/5.7:_Enthalpy_of_Formation. Reading guide style Adapted from "Ionic equations Study Guide" by Montgomery College is licensed under CC BY 4.0 Document not to be reposted on the internet without express permission.
