Class 33: Voltaic Cells Announcements Voltaic (Galvanic) Cells ...
Class 33: Voltaic Cells Announcements Voltaic (Galvanic) Cells ...
Class 33: Voltaic Cells Announcements Voltaic (Galvanic) Cells ...
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Chem 1011 – Intersession 2011<br />
<strong>Class</strong> #<strong>33</strong><br />
<strong>Class</strong> <strong>33</strong>: <strong>Voltaic</strong> <strong>Cells</strong><br />
• Sec 18.3 – <strong>Voltaic</strong> (or <strong>Galvanic</strong>) <strong>Cells</strong>:<br />
Generating Electricity from Spontaneous<br />
Chemical Reactions<br />
▫ Electrochemical Cell Notation<br />
<strong>Voltaic</strong> (<strong>Galvanic</strong>) <strong>Cells</strong><br />
• Electrical Current<br />
▫ flow of electrical charge.<br />
• Example: Electrons flowing through a wire.<br />
• Since redox reactions involve the transfer of electrons<br />
from one substance to another, they have the potential<br />
to generate electrical current.<br />
• Electrical current is measured in units of ampere (A),<br />
where:<br />
1 A = 1 C/s<br />
• And the charge on a single electron is 1.602 x 10 -19 C, so<br />
1 A corresponds to the flow of 6.424 x 10 18 electrons per<br />
second!<br />
<strong>Voltaic</strong> (<strong>Galvanic</strong>) <strong>Cells</strong><br />
• <strong>Voltaic</strong>* Cell<br />
▫ an electrochemical cell that produces electrical<br />
current from a spontaneous chemical reaction.<br />
• Electrolytic Cell<br />
▫ consumes electrical current to drive a<br />
nonspontaneous chemical reaction.<br />
• *<strong>Voltaic</strong> cells are also called <strong>Galvanic</strong> cells.<br />
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<strong>Announcements</strong><br />
• Review Sessions:<br />
▫ Thursday, June 16 th : 12pm – 2pm in C-30<strong>33</strong><br />
▫ Monday, June 20 th : 12pm – 2pm in C-4002<br />
• On the Web<br />
▫ Term Grades (tests only right now)<br />
▫ Answers to Tests #1 and #2<br />
• Final Exam<br />
▫ Tuesday, June 21 st : 9am – 11:30am in SN-4063<br />
<strong>Voltaic</strong> (<strong>Galvanic</strong>) <strong>Cells</strong><br />
<strong>Voltaic</strong> (<strong>Galvanic</strong>) <strong>Cells</strong><br />
Oxidation half-cell Reduction half-cell<br />
Oxidation half-reaction Reduction half-reaction<br />
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<strong>Voltaic</strong> <strong>Cells</strong> 1<br />
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Chem 1011 – Intersession 2011<br />
<strong>Class</strong> #<strong>33</strong><br />
<strong>Voltaic</strong> (<strong>Galvanic</strong>) <strong>Cells</strong><br />
Zn electrode<br />
(anode)<br />
Cu electrode<br />
(cathode)<br />
Electrode: conductive surfaces though which electrons can enter or leave half-cells<br />
<strong>Voltaic</strong> (<strong>Galvanic</strong>) <strong>Cells</strong><br />
Anions migrate<br />
toward anode<br />
<strong>Voltaic</strong> (<strong>Galvanic</strong>) <strong>Cells</strong><br />
Cations migrate<br />
toward cathode<br />
• The flow of electrons provides a usable source of<br />
electrical energy. If it were feasible to use this<br />
galvanic cell commercially, we would call it a<br />
battery.<br />
• As the galvanic cell operates, the redox reaction of<br />
the cell approaches equilibrium and the capacity of<br />
the cell to deliver useful electrical energy decreases.<br />
• At equilibrium, the cell ceases to function (think<br />
"dead battery"!)<br />
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<strong>Voltaic</strong> (<strong>Galvanic</strong>) <strong>Cells</strong><br />
• Anode<br />
▫ electrode where oxidation occurs. Since electrons flow out<br />
of the anode and into the external circuit, the anode is<br />
designated as negative (–).<br />
• Cathode<br />
▫ electrode where reduction occurs. Since electrons flow<br />
from the external circuit into the cathode, the cathode is<br />
designated as positive (+).<br />
• Electrons flow from the anode, to the cathode (from<br />
negative to positive) through the wires connecting the<br />
electrodes.<br />
<strong>Voltaic</strong> (<strong>Galvanic</strong>) <strong>Cells</strong><br />
• Salt Bridge<br />
▫ The function of the salt bridge is to allow the<br />
migration of positive and negative ions between the<br />
half-cells, thus completing the electrical circuit..<br />
• An example is an inverted, U-shaped tube,<br />
containing a strong electrolyte, suspended in gel,<br />
held in by permeable stoppers.<br />
• The salt bridge serves to complete the circuit,<br />
allowing electrical current to flow.<br />
Shorthand Cell Notation<br />
• Shorthand description of <strong>Voltaic</strong> cell:<br />
electrode | electrolyte || electrolyte | electrode<br />
Oxidation (Anode) Reduction (Cathode)<br />
• oxidation half-cell on the left; reduction half-cell on the<br />
right<br />
• single | = phase barrier<br />
▫ if multiple electrolytes in same phase, a comma is used<br />
rather than |<br />
▫ can use an inert electrode<br />
• double line || = salt bridge<br />
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Chem 1011 – Intersession 2011<br />
<strong>Class</strong> #<strong>33</strong><br />
Shorthand Cell Notation<br />
Anode = Zn (s)<br />
The anode is<br />
oxidized to Zn 2+<br />
Zn (s) | Zn 2+ (aq) || Cu 2+ (aq) | Cu (s)<br />
Inert Electrode<br />
Another Shocking <strong>Class</strong> at 11!<br />
Cathode = Cu (s)<br />
Cu 2+ ions are<br />
reduced at the<br />
cathode<br />
• Sec 18.4 – Standard Electrode Potentials<br />
▫ Predicting the Spontaneous Direction of an<br />
Oxidation Reduction Reaction<br />
▫ Predicting Whether a Metal will Dissolve in Acid<br />
• Sec 18.8 – Electrolysis: Driving Nonspontaneous<br />
Chemical Reactions with Electricity<br />
▫ Stoichiometry of Electrolysis<br />
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Inert Electrodes<br />
• Some electrochemical cells do not use metal strips as<br />
electrodes; instead, they use gaseous or aqueous<br />
reactants and platinum or graphite rods which function<br />
as inert electrodes.<br />
• When the participants in a half-reaction are all in the<br />
aqueous phase, a conductive surface is needed for<br />
electron transfer to take place. In such cases an inert<br />
electrode of graphite or platinum is often used.<br />
▫ In the following electrochemical cell, an iron strip acts as<br />
the anode and a platinum strip acts as the cathode. Iron is<br />
oxidized at the anode and MnO 4 – is reduced at the cathode.<br />
Problem<br />
• For the following voltaic cell, label each electrode by<br />
metal type, indicate whether it is the anode or cathode,<br />
and show the sign of each electrode. Within each halfcell<br />
write the correct balanced half-reaction. Indicate<br />
the direction of electron flow through the wire. Label the<br />
ions of the salt bridge (KNO 3) and indicate their<br />
direction of movement.<br />
Cr (s) | Cr 3+ (aq) || Ag + (aq) | Ag (s)<br />
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