Electrochemistry Chapter 4 4.4 and 4.8 Chapter 19 19.1-19.5 and 19.8.
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Transcript of Electrochemistry Chapter 4 4.4 and 4.8 Chapter 19 19.1-19.5 and 19.8.
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Electrochemistry
Chapter 44.4 and 4.8Chapter 1919.1-19.5 and 19.8
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Oxidation-Reduction Reactions
Electron Transfer Reactions Many redox reactions take place in water. Half-reactions OILRIG
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Oxidation-Reduction Reactions
2Mg (s) + O2 (g) 2MgO (s)
2Mg 2Mg2+ + 4e- Oxidation half-reaction (lose e-)
O2 + 4e- 2O2- Reduction half-reaction (gain e-)
2Mg + O2 + 4e- 2Mg2+ + 2O2- + 4e-
2Mg + O2 2MgO
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Oxidation-Reduction Reactions
Oxidizing Agent- a substance that accepts electrons from another substance, causing the other substance to be oxidized.
Reducing Agent- a substance that donates electrons to another substance, causing it to become reduced.
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Oxidation-Reduction Reactions
Zn (s) + CuSO4 (aq) ZnSO4 (aq) + Cu (s)
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Oxidation-Reduction Reactions
Zn (s) + CuSO4 (aq) ZnSO4 (aq) + Cu (s)
Zn Zn2+ + 2e- Zn is oxidized
Zn is the reducing agent
Cu2+ + 2e- Cu Cu2+ is reduced
Cu2+ is the oxidizing agent
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Oxidation-Reduction Reactions
Copper wire reacts with silver nitrate to form silver metal.What is the oxidizing agent in the reaction?
Cu (s) + 2AgNO3 (aq) Cu(NO3)2 (aq) + 2Ag (s)
Cu Cu2+ + 2e-
Ag+ + 1e- Ag Ag+ is reduced
Ag+ is the oxidizing agent
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Oxidation Number
Oxidation Number- the number of charges the atom would have in a molecule (or an ionic compound) if electrons were transferred completely.
H2(g) + Cl2(g) → 2HCl(g)
0 0 +1 -1
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Assigning Oxidation Numbers
1. Free elements (uncombined state) have an oxidation number of zero.
Na, Be, K, Pb, H2, O2, P4 = 0
2. In monatomic ions, the oxidation number is equal to the charge on the ion.
Li+, Li = +1; Fe3+, Fe = +3; O2-, O = -2
3. The oxidation number of oxygen is usually –2. In H2O2
and O22- it is –1.
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Assigning Oxidation Numbers
4. The oxidation number of hydrogen is +1 except when it is bonded to metals in binary compounds. In these cases, its oxidation number is –1.
5. Group IA metals are +1, IIA metals are +2 and fluorine is always –1.
6. The sum of the oxidation numbers of all the atoms in a molecule or ion is equal to the charge on the molecule or ion.
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Assigning Oxidation Numbers
Oxidation numbers of all the elements in HCO3
- ?
HCO3-
O = -2 H = +1
3x(-2) + 1 + ? = -1
C = +4
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Assigning Oxidation Numbers
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Redox Titrations
Oxidizing agent/Reducing Agent Equivalence point reached when reducing
agent is completely oxidized by the oxidizing agent.
Indicator is needed
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Redox Titrations
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Electrochemistry
Branch of Chemistry that deals with interconversion of electrical and chemical energy.
ImportanceUse of electricity in everyday life Interconversion of energy Study of electrochemical processes
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ElectrochemistryElectrochemical processes are oxidation-reduction reactions
in which:
• the energy released by a spontaneous reaction is converted to electricity or
• electrical energy is used to cause a nonspontaneous reaction to occur
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Balancing Redox Equations
Simple reactions easy to balance Complex reactions
Usually encountered in laboratoryCan include CrO4
2-, Cr2O72-, MnO4
-, NO3- and
SO42-
Follow balancing guidelines
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Balancing Redox Reactions
Guidelines Step 1: Write the unbalanced equation for the reaction in ionic
form. Step 2: Separate the equation into two half-reactions Step 3: Balance each half-reaction for number and type of atoms
and charges. (Look at medium) Step 4: Add the two half-reactions together and balance the final
equation by inspection. The electrons on both sides must cancel. (Be sure they are equal)
Step 5: Verify that the equation contains the same type and numbers of atoms and the same charges on both sides of the equation.
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Balancing Redox Equations
1. Write the unbalanced equation for the reaction ion ionic form.
The oxidation of Fe2+ to Fe3+ by Cr2O72- in acid solution?
Fe2+ + Cr2O72- Fe3+ + Cr3+
2. Separate the equation into two half-reactions.
Oxidation:
Cr2O72- Cr3+
+6 +3
Reduction:
Fe2+ Fe3++2 +3
3. Balance the atoms other than O and H in each half-reaction.
Cr2O72- 2Cr3+
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Balancing Redox Reactions4. For reactions in acid, add H2O to balance O atoms and H+ to
balance H atoms.Cr2O7
2- 2Cr3+ + 7H2O14H+ + Cr2O7
2- 2Cr3+ + 7H2O
5. Add electrons to one side of each half-reaction to balance the charges on the half-reaction.
Fe2+ Fe3+ + 1e-
6e- + 14H+ + Cr2O72- 2Cr3+ + 7H2O
6. If necessary, equalize the number of electrons in the two half-reactions by multiplying the half-reactions by appropriate coefficients.
6Fe2+ 6Fe3+ + 6e-
6e- + 14H+ + Cr2O72- 2Cr3+ + 7H2O
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Balancing Redox Reactions7. Add the two half-reactions together and balance the final
equation by inspection. The number of electrons on both sides must cancel.
6e- + 14H+ + Cr2O72- 2Cr3+ + 7H2O
6Fe2+ 6Fe3+ + 6e-Oxidation:
Reduction:
14H+ + Cr2O72- + 6Fe2+ 6Fe3+ + 2Cr3+ + 7H2O
8. Verify that the number of atoms and the charges are balanced.
14x1 – 2 + 6x2 = 24 = 6x3 + 2x3
9. For reactions in basic solutions, add OH- to both sides of the equation for every H+ that appears in the final equation.
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Galvanic Cells
Oxidation-Reduction Reaction Oxidizing agent and Reducing agent are not in
contact External conducting medium Galvanic Cell/Volatic Cell- the experimental
apparatus for generating electricity through the use of a spontaneous reaction.
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Copyright © The McGraw-Hill Companies, Inc. Permission required for reproduction or display.
e– e–
Zincanode
Coppercathode
Saltbridge
Cottonplugs
ZnSO4
solution
Zn is oxidized toZn2+ at anode.
2e– + Cu2+ (aq) Cu(s)
CuSO4
solution
Zn2+
SO 42–
Cl– K+
Voltmeter
(aq) + Cu(s)
Zn2+
2e–Cu2+
Cu
SO42–
Zn(s) Zn2+ (aq) + 2 e–
Net reaction
Cu
Cu2+ is reducedto Cu at cathode.
Zn(s) + Cu2+ (aq) Zn2+
2+
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Galvanic Cells
Cell Voltage- the difference in electrical potential between the anode and the cathode. Measured by a voltmeter.
Electromotive force (emf) and Cell Potential Voltage of a cell depends on:
Composition of electrode and ions Concentration of ions Temperature
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Galvanic Cells
Cell Diagram
Zn (s) + Cu2+ (aq) Cu (s) + Zn2+ (aq)
Zn (s) | Zn2+ (1 M) || Cu2+ (1 M) | Cu (s)
anode cathode
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Standard Reduction Potentials
Used to calculate the emf of a galvanic cell.
Use the values from anode and cathode
Standard reduction potential (E0) is the voltage associated with a reduction reaction at an electrode when all solutes are 1 M and all gases are at 1 atm.
Standard emf (E0 )cell
E0 = Ecathode - Eanodecell0 0
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Standard Reduction Potentials
Zn (s) | Zn2+ (1 M) || H+ (1 M) | H2 (1 atm) | Pt (s)
Anode (oxidation): Zn (s) Zn2+ (1 M) + 2e-
Cathode (reduction): 2e- + 2H+ (1 M) H2 (1 atm)
Zn (s) + 2H+ (1 M) Zn2+ + H2 (1 atm)
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Standard Reduction PotentialsE0 = 0.76 Vcell
Standard emf (E0 )cell
0.76 V = 0 - EZn /Zn 0
2+
EZn /Zn = -0.76 V02+
Zn2+ (1 M) + 2e- Zn E0 = -0.76 V
E0 = EH /H - EZn /Zn cell0 0
+ 2+2
E0 = Ecathode - Eanodecell0 0
Zn (s) | Zn2+ (1 M) || H+ (1 M) | H2 (1 atm) | Pt (s)
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Standard Reduction Potentials
19.3
Pt (s) | H2 (1 atm) | H+ (1 M) || Cu2+ (1 M) | Cu (s)
2e- + Cu2+ (1 M) Cu (s)
H2 (1 atm) 2H+ (1 M) + 2e-Anode (oxidation):
Cathode (reduction):
H2 (1 atm) + Cu2+ (1 M) Cu (s) + 2H+ (1 M)
E0 = Ecathode - Eanodecell0 0
E0 = 0.34 Vcell
Ecell = ECu /Cu – EH /H 2+ +2
0 0 0
0.34 = ECu /Cu - 00 2+
ECu /Cu = 0.34 V2+0
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Tips
E0 is for the reaction as written The more positive E0 the greater the tendency
for the substance to be reduced The half-cell reactions are reversible The sign of E0 changes when the reaction is
reversed Changing the stoichiometric coefficients of a
half-cell reaction does not change the value of E0
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Spontaneity of Redox Reactions
Relationship between Eºcell, ΔGº and K.
G0 = -nFEcell0 n = number of moles of electrons in reaction
F = 96,500J
V • mol = 96,500 C/mol
=0.0257 V
nln KEcell
0 =0.0592 V
nlog KEcell
0
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Spontaneity of Redox Reactions
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Relationships between Eºcell, ΔGº and K
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The Effect of Concentration on Cell Emf Not all reactions in galvanic cells can
occur under standard state conditions. Nernst Equation- equation used to
calculate cell potential under nonstandard state conditions.
-0.0257 V
nln QE0E = -
0.0592 Vn
log QE0E =
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Will the following reaction occur spontaneously at 250C if [Fe2+] = 0.60 M and [Cd2+] = 0.010 M? Fe2+ (aq) + Cd (s) Fe (s) + Cd2+ (aq)
2e- + Fe2+ 2Fe
Cd Cd2+ + 2e-Oxidation:
Reduction: n = 2
E0 = -0.44 – (-0.40)
E0 = -0.04 V
E0 = EFe /Fe – ECd /Cd0 0
2+ 2+
-0.0257 V
nln QE0E =
-0.0257 V
2ln -0.04 VE =
0.0100.60
E = 0.013
E > 0 Spontaneous
19.5
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Concentration Cells
Galvanic cells consisting of the same type of electrodes in the same solution. Each solution is a different concentration.
Zn(s) | Zn2+(0.10M) || Zn2+(1.0M) | Zn(s)
E = Eº - 0.0257V/ 2 x ln [Zn2+] dil / [Zn2+]conc
E = 0- 0.0257V/2 ln (0.10/1.0) E = 0.0296 V
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Concentration Cells
ImportanceUnequal ion concentrationsMembrane potentialPropagation
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Electrolysis
Electrolysis- is the process in which electrical energy is used to cause a nonspontaneous chemical reaction to occur.
Electrolytic Cell- an apparatus for carrying out electrolysis.