SAHAR ADHAM LECTUERE OF PHYSICAL CHEMISTRY

galvanicelectrolytic need power source two electrodes produces electrical current anode (-) cathode (+) anode (+) cathode (-) salt bridge vessel conductive medium Comparison of Electrochemical Cells E ° cell > 0. E ° cell < 0.

 Electron transfer reactions are oxidation-reduction or redox reactions.  Results in the generation of an electric current (electricity) or be caused by imposing an electric current.  Therefore, this field of chemistry is often called ELECTROCHEMISTRY.

 oxidation: loss of electrons  reduction: gain of electrons LEO the lion says GER! GER!

 Batteries  Corrosion  Industrial production of chemicals such as Cl 2, NaOH, F 2 and Al  Biological redox reactions The heme group

 One ½ cell rxn. occurs in each compartment.  Zn  Zn e – in the anode.  Cu e –  Cu in cathode.  But not without a connection. Zn Zn 2+ Cu 2+ Zn + Cu 2+  Zn 2+ + Cu SO 4 2– Cu Anode=Oxidation Cathode=Reduction

 But even with a connection of the electrodes, no current flows.  We need to allow neutrality in the solutions with a salt bridge to shift counterions. Zn Zn 2+ Cu 2+ Zn + Cu 2+  Zn 2+ + Cu SO 4 2– Cu 2e –

Cell Potential Cell Potential or Electromotive Force (emf): The “pull” or driving force on the electrons.

 The voltage generated by the Zn/Cu galvanic cell is +1.1V under standard conditions.  Standard conditions are:  T = 25°C and P = 1 bar for gases.  Solids and liquids are pure.  Solutions are 1 M in all species.  E ° cell is sum of ½ cell E ° values.

CELL POTENTIALS AND REDUCTION POTENTIALS E°cell = E°reduced - E°oxidized E°cell = E°cathode - E°anode

 All ½ cells are catalogued as reduction reactions & assigned reduction potentials, E °.  The lower reduction potential ½ rxn is reversed to become the oxidation. E ° oxidation = – E ° reduction  That makes spontaneous E ° cell > 0.  But E ° red can’t be found w/o E ° ox !

 We need a standard electrode to make measurements against!  The Standard Hydrogen Electrode (SHE)  2H + (aq) + 2e –  H 2 (1 bar) E °  0 V  1 bar H 2 flows over a Pt electrode, and the full E ° cell is assigned to the other electrode. E ° SHE = 0 V.  E.g., standard calomel electrode:  Hg 2 Cl 2 (s) + 2e –  2 Hg( l ) + Cl – E ° SCE = +0.27V  a more physically convenient reference.

 Shorthand for a complete redox cell is of the form:  Anode | anodic soln. || cathodic soln. | Cathode  So making a cell of Cu corrosion,  Cu | Cu 2+ || NO 3 –, NO(g), H + |Pt  where all ions should be suffixed (aq) and both metals should have (s).

Primary Battery : can not be recharged e.g. Mercury Battery Secondary Battery: rechargeable (storage batteries) e.g. Ni-Cad Battery Fuel Cell: reactants supplied from an external source e.g. H2/O2 fuel cells.

MERCURY BATTERY Anode: Zn is reducing agent under basic conditions : Cathode : HgO + H 2 O + 2e- ---> Hg + 2 OH - can not be recharged

NI-CAD BATTERY Anode (-) Cd + 2 OH - ---> Cd(OH) 2 + 2e- Cathode (+) NiO(OH) + H 2 O + e- ---> Ni(OH) 2 + OH - rechargeable

It is because the products of the reaction are solids that the Ni-Cd battery can be recharged The solid hydroxides are sticky, and remain in place. If current is applied, the reaction can be driven backwards !

When you charge a battery, you are forcing the electrons backwards (from the + to the -). To do this, you will need a higher voltage backwards than forwards.

But in mercury battery the ZnO is not sticky, and doesn’t remain attached to the electrode. This battery is not rechargeable

H 2 AS A FUEL Cars can use electricity generated by H 2 /O 2 fuel cells. H 2 carried in tanks or generated from hydrocarbons

17-44 Galvanic cells for which the reactants are continuously supplied. anode: 2H 2 + 4OH   4H 2 O + 4e  cathode : 4e  + O 2 + 2H 2 O  4OH  2H 2 (g) + O 2 (g)  2H 2 O(l) Fuel Cells

Mercury batteries take advantage of the high density of Hg to be quite small: used in watches, hearing aids, calculators, etc. Lithium-iodine batteries are particularly small and lightweight, but also very long-lived Often used in pacemakers, where they can last for 10 years

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V1.06

For a galvanic cell, the electrode at which reduction occurs is called the: A: AnodeB: Cathode final 5/50 Dr. Keck poll

For a galvanic cell, the electrode at which reduction occurs is called the: B: Cathode

For a galvanic cell, the electrode with negative polarity is called the: A: AnodeB: Cathode final 50/50 Dr. Keck poll

For a galvanic cell, the electrode with negative polarity is called the: A: Anode

Which of the following statements is incorrect a. In a galvanic cell, reduction occurs at the anode. b. The cathode is labeled "+" in a voltaic cell. c. Oxidation occurs at the anode in a voltaic cell. d. Electrons flow from the anode to the cathode in all electrochemical cells. 32

a. In a galvanic cell, reduction occurs at the anode.

Consider the following notation for an electrochemical cell Zn|Zn 2+ (1M)||Fe 3+ (1M), Fe 2+ (1M)|Pt What is the balanced equation for the cell reaction? a. Zn(s) + 2Fe 3+ (aq) → 2Fe 2+ (aq) + Zn 2+ (aq) b. Zn 2+ (aq) + 2Fe 2+ (aq) → Zn(s) + 2Fe 3+ (aq) c. Zn(s) + 2Fe 2+ (aq) → 2Fe 3+ (aq) + Zn 2+ (aq) d. Zn(s) + Fe 3+ (aq) → Fe 2+ (aq) + Zn 2+ (aq) e. Zn(s) + Fe 2+ (aq) → Fe(s) + Zn 2+ (aq) 34

 Zn(s) + 2Fe 3+ (aq) → 2Fe 2+ (aq) + Zn 2+ (aq)

What is the oxidation state of nitrogen in HNO 3 ? A: +3B: +4 C: +5D: -5 final50/50 Dr. Keck poll

What is the oxidation state of nitrogen in HNO 3 ? C: +5

38 Consider the following electrode potentials: Mg e –  Mg E° = –2.37 V V e –  V E° = –1.18 V Cu 2+ + e –  Cu + E° = 0.15 V Which one of the reactions below will proceed spontaneously from left to right? a. Mg 2+ + V  V 2+ + Mg b. Mg Cu +  2Cu 2+ + Mg c. V Cu +  V +2 + Cu 2+ d. V + 2Cu 2+  V Cu + e. none of these

 d. V + 2Cu 2+  V Cu +

What is the oxidative state of iodine in IO 3 - ? A: +7B: +6 C: +5D: +4 final50/50 Dr. Keck poll

What is the oxidative state of iodine in IO 3 - ? C: +5