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CHEM 1133- Midterm Exam Guide - Comprehensive Notes for the exam ( 44 pages long!)


Department
Chemistry
Course Code
CHEM 1133
Professor
Ruth Heisler
Study Guide
Midterm

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CU-Boulder
CHEM 1133
MIDTERM EXAM
STUDY GUIDE

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Exam 3 Chemistry Short Study Guide
Balancing Redox Reactions (21.1)
Steps for balancing redox reactions
1. Split into two half-reactions
2. Balance atoms and charges
a. Balance atoms other than O and H
b. Balance O with H2O
c. Balance H with H+
d. Balance charge with e-s (here pay attention to charge, not O.N)
3. Multiply each half-rxn by an integer so e-s cancel
4. Add half-rxns together and cancel any common species
5. If the solution is BASIC, add OH- to each side to react with H+s
Electrochemical Cells (21.2)
Redox reactions have the potential to generate electrical current, the flow of electric charge
Two types of electrochemical cells:
o Voltaic cells (galvanic cells): electrochemical cell that produces electrical current from a spontaneous chemical rxn (∆G < 0)
Ex: batteries
o Electrolytic cells: electrochemical cell that consumes electrical current to drive a nonspontaneous chemical rxn (∆G > 0)
Ex: recharging batteries
Components of voltaic and electrolytic cells
o Each half-rxn takes place in its own half-cell, so the the reactions are physically separate
o Each half-cell consists of an electrode in an electrolyte solution
Electrode: solid metal pieces that conduct electricity between the cell and the
surroundings which are dipped into an electrolyte
Electrolyte: mixture of ions involved in the reaction or carry charge
o The half-cells are connected by an external circuit
o A salt bridge completes the electrical circuit
Operation of the voltaic cell
o Anode: oxidation occurs
Oxidation (loss of e- occurs at the anode, the source of e-
Overtime, the metal anode decreases in mass and the ion in the electrolyte
solution increases
Negative electrode
o Cathode: reduction occurs
Reduction occurs at the cathode, where the e-s are used up
Overtime, the ion concentration in this half-cell decreases and the mass of the metal cathode increases
Positive electrode
o A salt bridge allows ions to flow through bot compartments to maintain charge balance
Usually contains non-reacting cations and anions
Active and Inactive electrodes
o An active electrode is an active component in its half-cell and is a reactant or product in the overall rxn
o An inactive electrode provides a surface for the reaction and completes the circuit. It does not participate actively in the overall
rxn.
Inactive electrodes are necessary when none of the rxn components can be used as an electrode.
o Inactive electrodes are usually unreactive substances such as graphite or platinum.
Clicker: True statements of a voltaic cell:
o True: oxidation occurs at the anode and reduction occurs at the cathode
o True: electrons travel through a wire from the anode to the cathode
o True: it contains a spontaneous reaction
o True: the anode compartments build up with positive charge and thus, anions travel from the cathode to the anode to maintain
charge neutrality
Cell Potentials (21.3)
Cell potential (Ecell)
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Exam 3 Chemistry Short Study Guide
o Potential differences gives rise to force that results in motion of electrons and is referred to as electromotive force (emf)
Electrical current is driven by a difference in potential energy, caused by an electrical field resulting from the charge
difference on the two electrodes, called potential difference
The SI unit of potential difference is called a volt (V): 1 V =
A large potential difference corresponds to a large difference in charge between the two electrodes and therefore, a
strong tendency for electron to flow
o The potential difference between the 2 electrodes in a voltaic cell is called the cell potential (Ecell) or cell emf
o The electrodes in each half-cell have their own individual potential called the standard electrode potential, designated as either
cathode or E°anode
o When the half-cells are connected, e flow from the electrode with the more negative charge (greater potential energy) to the
electrode with the more positive charge (less potential energy)
o A voltaic cell converts free energy change of a spontaneous rxn into electrical energy
o When all species are in their standard state:
cell= E°cathode - anode
o The spontaneity of a cell can be predicted from E°cell or from E°cathode and E°anode
If cell > 0, rxn is spontaneous
Standard Hydrogen Electrode (SHE)
o All electrode potentials are measured relative to SHE which is arbitrarily assigned an electrode potential of zero
o The half-cell for the SHE consists of an inert platinum electrode immersed in 1 M HCl with H2 gas at 1 atm bubbling through
the solution
The Pt electrode has two functions
Provides a surface for the dissociation and oxidation of H2 molecule
H2 2H++ 2e - E°=0.00V
Serves as an electrical conductor to the external circuit
o By connecting the SHE to an electrode of another half-cell, the potential difference (or V) between the two electrodes can be
measured and the electrode potential of the other half-cell can be determined
Like ∆G°, we can use the sign of E° to predict if a rxn is spontaneous
o If cell > 0, rxn is spontaneous
o If cell < 0, rxn is not spontaneous
Free Energy and Electrical Work (21.4)
Predicting if reactions will occur:
o Positive E°cell will occur
o Negative E°cell will NOT occur
The relationship between ∆G° and E°cell
o Spontaneous reactions:
From thermodynamics: G° < 0 (-)
From electrochemistry: cell > 0 (+)
o cell (V) is a measure of the difference of potential energy per unit of charge
E° (V) = 

Potential energy difference = Wmax
Charge = Q
o Can quantify the charge (Q) that flows in an electrochemical reaction by using Faraday’s constant (F)
G° = -nF E°cell where F = 9.65 x 104 J/Vmole
G° represents the max amount of work that can be done by a reaction
Connecting E°cell and K
cell
K
spontaneous
+
-
> 1
Products favored
Not spontaneous
-
+
< 1
Reactants favored
o cell =
 
At standard conditions: cell =

Non-standard conditions:
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