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Galvanic Cell Notation • Half-Cell Notation • Types of Electrodes • Cell

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Galvanic Cell Notation • Half-Cell Notation • Types of Electrodes • Cell Galvanic Cell Notation ¾Inactive (inert) electrodes – not involved in the electrode half-reaction (inert solid conductors; • Half-cell notation serve as a contact between the – Different phases are separated by vertical lines solution and the external el. circuit) 3+ 2+ – Species in the same phase are separated by Example: Pt electrode in Fe /Fe soln. commas Fe3+ + e- → Fe2+ (as reduction) • Types of electrodes Notation: Fe3+, Fe2+Pt(s) ¾Active electrodes – involved in the electrode ¾Electrodes involving metals and their half-reaction (most metal electrodes) slightly soluble salts Example: Zn2+/Zn metal electrode Example: Ag/AgCl electrode Zn(s) → Zn2+ + 2e- (as oxidation) AgCl(s) + e- → Ag(s) + Cl- (as reduction) Notation: Zn(s)Zn2+ Notation: Cl-AgCl(s)Ag(s) ¾Electrodes involving gases – a gas is bubbled Example: A combination of the Zn(s)Zn2+ and over an inert electrode Fe3+, Fe2+Pt(s) half-cells leads to: Example: H2 gas over Pt electrode + - H2(g) → 2H + 2e (as oxidation) + Notation: Pt(s)H2(g)H • Cell notation – The anode half-cell is written on the left of the cathode half-cell Zn(s) → Zn2+ + 2e- (anode, oxidation) + – The electrodes appear on the far left (anode) and Fe3+ + e- → Fe2+ (×2) (cathode, reduction) far right (cathode) of the notation Zn(s) + 2Fe3+ → Zn2+ + 2Fe2+ – Salt bridges are represented by double vertical lines ⇒ Zn(s)Zn2+ || Fe3+, Fe2+Pt(s) 1 + Example: A combination of the Pt(s)H2(g)H Example: Write the cell reaction and the cell and Cl-AgCl(s)Ag(s) half-cells leads to: notation for a cell consisting of a graphite cathode - 2+ Note: The immersed in an acidic solution of MnO4 and Mn 4+ reactants in the and a graphite anode immersed in a solution of Sn 2+ overall reaction are and Sn . in different phases →Write the half reactions (a list of the most common (no physical half-reactions is given in Appendix D) contact) ⇒ no need Sn2+ → Sn4+ + 2e- ×5 (oxidation) + - + - 2+ of a salt bridge MnO4 + 8H + 5e → Mn + 4H2O(l) ×2 (reduction) + - H2(g) → 2H + 2e (anode, oxidation) 2+ - + - 4+ - + 5Sn + 2MnO4 + 16H + 10e → 5Sn + 10e + AgCl(s) + e- → Ag(s) + Cl- (×2) (cathode, reduction) 2+ + 2Mn + 8H2O(l) + - 2AgCl(s) + H2(g) → 2Ag(s) + 2H + 2Cl →The graphite (C) electrodes are inactive + - 2+ 4+ + - 2+ ⇒ Pt(s)H2(g)H ,ClAgCl(s)Ag(s) ⇒ C(s)Sn ,Sn || H ,MnO4 ,Mn C(s) Why Do Galvanic Cells Work? 21.3 Cell Potentials • Consider a cell made of two active metal • Electromotive force (emf) – drives the electrodes, M and M , and their ions. 1 2 electrons in the el. circuit ¾ If the cell circuit is open, the two metals are in equilibrium with their ions – emf is the difference between the electrical + - + - potentials of the two electrodes (voltage) 1) M1 ↔ M1 + e 2) M2 ↔ M2 + e ¾ The produced electrons accumulate in the metal • Cell potential (Ecell) → Ecell = emf electrodes and produce electrical potentials – Units → volts (V) → (1 V = 1 J/C since the ¾ If M1 has a greater tendency to give out its electrons, electrical work is equal to the applied voltage st the 1 equilibrium is shifted further to the right and times the charge moving between the electrodes) the potential of M is more negative 1 o ¾ When the circuit is closed, electrons flow from the • Standard cell potential (E cell) – the cell more negative M1 (anode) toward the less negative potential at standard-state conditions (gases → M2 (cathode) 1 atm, solutions → 1 M, liquids & solids → pure) 2 ¾Ecell is measured with a voltmeter • Electrode potentials (E) – characterize the ¾If the (+) terminal of the voltmeter is connected individual electrodes (half-reactions) to the (+) electrode (cathode), the voltmeter – The cell potential is the difference between the shows a positive reading electrode potentials of the cathode and anode ¾Ecell characterizes the overall cell reaction Ecell = Ecathode – Eanode ¾If E > 0, the cell reaction is spontaneous cell • Standard electrode potentials (Eo) – ¾ If Ecell < 0, the cell reaction is non-spontaneous electrode potentials at the standard-state ¾If Ecell = 0, the cell reaction is at equilibrium o o o E cell = E cathode – E anode 2+ 2+ Example: Zn(s)Zn (1M)|| Cu (1M)Cu(s) – Eo values are reported for the half-reaction +1.10 V written as reduction (standard reduction potentials) → listed in Appendix D Zn(s) + Cu2+ → Zn2+ + Cu(s) o E cell = 1.10 V > 0 → spontaneous reaction ¾Absolute values for E and Eo can’t be measured – If the unknown electrode is the cathode of the cell o o o ⇒ A reference electrode (half-cell) is needed → E cell = E unkn – E ref o o o o o • The potentials of all electrodes are measured relative → E unkn = E cell + E ref = E cell +0= E cell > 0 to the reference electrode – If the unknown electrode is the anode of the cell o o o • Standard hydrogen electrode – used as a → E cell = E ref – E unkn o o o o o o reference electrode → E ref = 0 V (assumed) → E unkn = E ref – E cell = 0–E cell = –E cell < 0 + H (1M)H2(g, 1atm)Pt(s) Example: + - + - Pt(s)H2(g, 1atm)H (1M),Cl(1M)AgCl(s)Ag(s) 2H (1M) + 2e → H2(g, 1atm) + – To find the potential of any electrode, a cell is H /H2 → anode constructed between the unknown electrode and Ag/AgCl → cathode o o the reference electrode E cell = E Ag/AgCl –Eref = Eo – The cell potential is directly related to the Ag/AgCl o unknown electrode potential E Ag/AgCl = +0.22 V 3.
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