Electrochemistry Part II: The Galvanic Cell Jespersen Chap. 20 Sec 1 Dr. C. Yau Fall 2014 1 1

What is a Galvanic Cell? A galvanic cell is a spontaneous electrochemical cell in which electricity is produced by a spontaneous redox reaction. The resulting electron transfer is forced to take place through a wire. It is also known as a voltaic cell. Do not confuse it with the electrolytic cell (discussed later) where an electrochemical reaction is forced to take place by passing electricity through the cell.

Comparison of Types of Cell Galvanic or Voltaic Cell: Spontaneous electrochemical rxn Electricity produced Electrolytic Cell: Non-spontaneous electrochemical rxn Electric current is passed thru a wire to force the reaction to take place.

Anatomy Of A Galvanic Cell Note: Textbk is inconsistent with which is on the left (cathode or anode). There is no convention which is on the left side in a sketch of this sort. cathode anode Zn  Zn2+ + 2 e− Cu2+ (aq) + 2e− Cu (s) Half-cells (compartments containing reactants for each half- reaction) Electrodes to conduct current through the solution. Salt bridge to allow ion movement to keep solns neutral. Supporting electrolyte (spectator ions: NO3-) Connecting external circuit (wire and voltmeter) Figure 19.2 A galvanic cell. The cell consists of two half-cells. Oxidation takes place in one half cell and reduction in the other as indicated by the half reactions. 4

Electrochemical Cells In all cells, electrons transfer between the cathode (the reduction half-cell) and the anode (the oxidation half-cell) REMEMBER! "Red-Cat and An-Ox” Reduction at the Cathode & Oxidation at the Anode Figure 19.3 Expanded view of Figure 19.2 to show changes that take place at the anode and cathode in the copper-silver galvanic cell. (Not drawn to scale.) At the anode, Cu2+ ions enter the solution when copper atoms are oxidized, leaving electrons behind on the electrode. Unless Cu2+ ions move away from the electrode or NO3– ions move toward it, the solution around the electrode will become positively charged. At the cathode, Ag+ ions leave the solution and become silver atoms by acquiring electrons from the electrode surface. Unless more silver ions move toward the cathode or negative ions move away, the solution around the electrode will become negatively charged. 5

A closer look at the electrodes: Cu2+ Cu cathode Zn2+ Zn anode Zn Cu Reduction of Cu2+ to Cu e- extracted from electrode (soln becomes more negative) Oxidation of Zn to Zn2+ Leaves e- behind on the electrode (soln becomes more positive) 6

Electrochemical Cells Electrical current is conducted via the movement of electrons and ions. To prevent charge buildup, a salt bridge allows ions to move between the cells. REMEMBER! Electrons flow from anode to cathode through the wire. a to c Cations move towards the cathode. Anions move towards the anode. Red-Cat and An-Ox. 7

There is a buildup of what charges at each cell? Zn2+ Zn Cu KCl or KNO3 There is a buildup of what charges at each cell? The salt bridge often made of KCl or KNO3 (unreactive ions – spectator ions) What ions in the salt bridge move to which cell? 8

C. It depends on the reaction Towards which compartment will electrons flow in an electrochemical cell? A. Toward the cathode B. Toward the anode C. It depends on the reaction Through which components of the cell will ions not flow? A. The electrodes B. The solution C. The salt bridge http://www.mhhe.com/physsci/chemistry/essentialchemistry/flash/galvan5.swf anode to cathode 9

Standard Cell Notation (Line Cell Notation) salt bridge Rxn at anode Rxn at cathode Zn (s) | Zn2+ (aq) || Cu2+ (aq) | Cu (s) anode electrode cathode electrode anode electrolyte cathode Cell reactions separated by || that represents the salt bridge with ANODE on left, CATHODE on right. Electrodes appear at the outsides Reaction electrolytes in inner section Phases (phys. States) separated with | Species in the same state separated with ; Concentrations shown in ( ) If reaction includes no conductive metal, Pt or C(gr) is used as the electrode ChemFAQ: Describe a galvanic cell using standard cell notation. 10

Standard Cell Notation (Line Cell Notation) Rxn at Write the half reactions for the galvanic cell shown above. Cu (s) Cu2+ (aq) + 2e- Ag+ (aq) + e- Ag (s) Make a sketch of the galvanic cell and label it fully. If reaction includes no conductive metal, Pt or C(gr) is used as the electrode ChemFAQ: Describe a galvanic cell using standard cell notation. 11

Now, consider the reaction of Al3+(aq) + Zn (s) Al(s) + Zn2+(aq) Write the half-reactions. Balance the electrons and write the balanced net ionic equation. Sketch the galvanic cell (electrochemical cell). Label it fully. Write the standard cell notation. Do Prac Exer 1 & 2 on p. 924 12

Given: Mg(s) | Mg2+(aq) || Sn2+(aq) | Sn(s) Sketch the galvanic cell corresponding to this standard cell notation. Label it fully. Practice with p. 969 #20.50, 20.51 13

Now, consider the reaction of Fe3+ + Zn Fe2+ + Zn2+ Write the half-reactions. Balance the electrons and write the balanced net ionic equation. Write the standard cell notation. Sketch the galvanic cell (electrochemical cell). Label it fully. How can you have an electrode that is an ion (such as Fe3+)? 14

Where there are no conductive metals involved in a process, an inert electrode is used. C(gr) and Pt are often used. 2Fe3+ + Zn 2Fe2+ + Zn2+ Zn (s) Zn2+(aq) Fe3+(aq) Fe2+(aq) Zn(s) |Zn2+(aq) || Fe3+(aq); Fe2+(aq)|Pt(s) Zn inert anode cathode (where Fe3+ reduces to Fe2+ at the surface of the Pt electrode)

Balance and identify the cathode and anode H2O2(aq) + CO2(g) → H2C2O4(aq) + O2(g) (acidic) H2O2(aq) → O2(g) +2H+ + 2e- 2H+ + 2e- + 2CO2(g) → H2C2O4(aq) + (acidic) oxid reduc H2O2(aq) + 2CO2(g) → H2C2O4(aq) + O2(g) (acidic) Which is at the cathode? At the anode? 16

Write Line Notation for the cell: H2O2(aq) + CO2(g) → H2C2O4(aq) + O2(g) (acidic) Standard Cell notation for the reaction: oxidation reduction C(gr)| H2O2(aq) ;H+|O2(g)||CO2(g)|H2C2O4(aq); H+|C(gr) 17

Balance and identify the cathode and anode CrO3(s) + MnO2(s)→MnO4-(aq) + Cr3+(aq) (basic) CrO3(s) + 3H2O(l) + 3e- → Cr3+(aq) +6OH-(aq) MnO2(s) + 4OH- → MnO4-(aq) + 2H2O + 3e- CrO3(s) + MnO2(s) + H2O(l) →MnO4-(aq) + Cr3+(aq) + 2OH-(aq) Balancing redox equations by half-reaction method is given in Sec 6.2 (p. 222) 18

Galvanic Cells without Metal Electrodes Equation from previous slide: CrO3(s)+ MnO2(s) + H2O(l) →MnO4-(aq)+Cr3+(aq) + 2OH-(aq) Write the Standard Cell Notation: C(gr);MnO2(s)|MnO4-(aq)||CrO3(s)|Cr3+(aq);OH-|C(gr) p. 970 #20.52