Chris Cialek & Mark Kotz Standards Committee January 11, 2011.
Chemistry and Chemical Reactivity 6th Edition John C. Kotz Paul M. Treichel Gabriela C. Weaver...
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Transcript of Chemistry and Chemical Reactivity 6th Edition John C. Kotz Paul M. Treichel Gabriela C. Weaver...
Chemistry and Chemical Reactivity 6th Edition
John C. Kotz Paul M. Treichel
Gabriela C. Weaver
Principles of Reactivity: Electron Transfer Reactions
© 2006 Brooks/Cole Thomson
Lectures written by John Kotz
TRANSFER REACTIONS
Atom/Group transfer
HCl + H2O ---> Cl- + H3O+
Electron transfer
Cu(s) + 2 Ag+(aq) ---> Cu2+(aq) + 2 Ag(s)
Electron Transfer Reactions
• Electron transfer reactions are oxidation-reduction or redox
reactions.
• Redox reactions can result in the generation of an electric
current or be caused by imposing an electric current.
• Therefore, this field of chemistry is often called
ELECTROCHEMISTRY.
Review of Terminology for Redox Reactions
• OXIDATION—loss of electron(s) by a species; increase in oxidation number.
• REDUCTION—gain of electron(s); decrease in oxidation number.
• OXIDIZING AGENT—electron acceptor; species is reduced.
• REDUCING AGENT—electron donor; species is oxidized.
OXIDATION-REDUCTION REACTIONS
Direct Redox ReactionOxidizing and reducing agents in direct contact.
Cu(s) + 2 Ag+(aq) ---> Cu2+(aq) + 2 Ag(s)
Balancing EquationsStep 1: Divide the reaction into half-reactions, one for
oxidation and the other for reduction.
Ox Cu ---> Cu2+
Red Ag+ ---> Ag
Step 2: Balance each for mass. Already done in this case.
Step 3: Balance each half-reaction for charge by adding electrons.
Ox Cu ---> Cu2+ + 2e-
Red Ag+ + e- ---> Ag
Balancing EquationsStep 4: Multiply each half-reaction by a factor so that
the reducing agent supplies as many electrons as the oxidizing agent requires.
Reducing agent Cu ---> Cu2+ + 2e-
Oxidizing agent 2 Ag+ + 2 e- ---> 2 Ag
Step 5: Add half-reactions to give the overall equation.
Cu + 2 Ag+ ---> Cu2+ + 2Ag
The equation is now balanced for both charge and mass.
Balancing Equations for Redox Reactions
Some redox reactions have equations that must be balanced by special techniques.
MnO4- + 5 Fe2+ + 8 H+---> Mn2+ + 5 Fe3+ + 4 H2O
Mn = +7 Fe = +2 Fe = +3Mn = +2
Balancing EquationsBalance the following in acid solution— VO2
+ + Zn ---> VO2+ + Zn2+
Step 1: Write the half-reactions
Ox Zn ---> Zn2+
Red VO2+ ---> VO2+
Step 2: Balance each half-reaction for mass.
Ox Zn ---> Zn2+
Red VO2+ ---> VO2+ + H2O2 H+ +
Add H2O on O-deficient side and add H+ on other side for H-balance.
Balancing EquationsStep 3: Balance half-reactions for charge.
Ox Zn ---> Zn2+ + 2e-
Red e- + 2 H+ + VO2+ ---> VO2+ + H2O
Step 4: Multiply by an appropriate factor.
Ox Zn ---> Zn2+ + 2e-
Red 2e- + 4 H+ + 2 VO2+
---> 2 VO2+ + 2 H2O
Step 5: Add balanced half-reactions
Zn + 4 H+ + 2 VO2+
---> Zn2+ + 2 VO2+ + 2 H2O
Tips on Balancing Equations
• Never add O2, O atoms, or O2- to balance oxygen ONLY add H2O or OH-.
• Never add H2 or H atoms to balance hydrogen ONLY add H+ or H2O.
• Be sure to write the correct charges on all the ions.
• Check your work at the end to make sure mass and charge are balanced.
• PRACTICE!
Potential Ladder for Reduction Half-Reactions
Best oxidizing agents
Best reducing agents
Figure 20.14
TABLE OF STANDARD REDUCTION POTENTIALS
2
Eo (V)
Cu2+ + 2e- Cu +0.34
2 H+ + 2e- H 0.00
Zn2+ + 2e- Zn -0.76
oxidizingability of ion
reducing abilityof element
Using Standard Potentials, Eo
Table 20.1
• Which is the best oxidizing agent:
O2, H2O2, or Cl2? _________________
• Which is the best reducing agent:
Hg, Al, or Sn? ____________________
Standard Redox Potentials, Eo
Any substance on the right will reduce any substance higher than it on the left.
• Zn can reduce H+ and Cu2+.
• H2 can reduce Cu2+ but not
Zn2+
• Cu cannot reduce H+ or Zn2+.
Standard Redox Potentials, Eo
Cu2+ + 2e- --> Cu +0.34
+2 H + 2e- --> H2 0.00
Zn2+ + 2e- --> Zn -0.76
Northwest-southeast rule: product-favored reactions occur between • reducing agent at southeast corner • oxidizing agent at northwest corner
Any substance on the right will reduce any substance higher than it on the left.
Ox. agent
Red. agent
CELL POTENTIALS, Eo
Can’t measure 1/2 reaction Eo directly. Therefore, measure it relative to a STANDARD HYDROGEN CELL
2 H+(aq, 1 M) + 2e- <----> H2(g, 1 atm)
Eo = 0.0 V
Calculating Cell Voltage
• Balanced half-reactions can be added together to get overall, balanced equation.
Zn(s) ---> Zn2+(aq) + 2e-Cu2+(aq) + 2e- ---> Cu(s)--------------------------------------------Cu2+(aq) + Zn(s) ---> Zn2+(aq) + Cu(s)
If we know Eo for each half-reaction, we could get Eo for net reaction.
Uses of Eo Values
Organize half-reactions by relative ability to act as oxidizing agents
• Use this to predict direction of redox reactions and cell potentials.
Cu2+(aq) + 2e- ---> Cu(s) Eo = +0.34 VZn2+(aq) + 2e- ---> Zn(s) Eo = –0.76 V
Note that when a reaction is reversed the sign of E˚ is reversed!
Using Standard Potentials, Eo
• In which direction do the following reactions
go?
• Cu(s) + 2 Ag+(aq) ---> Cu2+(aq) + 2 Ag(s)
– Goes right as written
• 2 Fe2+(aq) + Sn2+(aq) ---> 2 Fe3+(aq) + Sn(s)
– Goes LEFT opposite to direction written
• What is Eonet for the overall reaction?
Eo and Thermodynamics
• Eo is related to ∆Go, the free energy change for the reaction.
• ∆G˚ is proportional to –nE˚
∆Go = -nFEo where F = Faraday constant
= 9.6485 x 104 J/V•mol of e-
(or 9.6485 x 104 coulombs/mol)and n is the number of moles of electrons
transferred
Eo and ∆Go
∆Go = - n F Eo For a product-favored reaction Reactants ----> Products
∆Go < 0 and so Eo > 0Eo is positive
For a reactant-favored reaction Reactants <---- Products
∆Go > 0 and so Eo < 0Eo is negative