17.1 Galvanic Cells (Batteries) Ch. 17: Electrochemistry 17.1 Galvanic Cells (Batteries)

Galvanic Cells device that changes chemical energy into electrical energy uses a reduction/oxidation reaction reducing agent transfers electrons to the oxidizing agent oxidation: loss of electrons (  ox. state) reduction: gain of electrons (  ox state)

Galvanic Cells if the two half-reactions are combined in the same container, the electron exchange occurs directly as work is released as heat= no electricity to harness the energy, we keep each half-reaction in a separate container so the electron transfer occurs through a wire = Electricity

Galvanic Cell salt bridge or porous disk used to allow for unrelated ions to move to allow for balance of charge

Parts of Galvanic Cell anode: oxidation cathode: reduction Names for Q. In which direction will the electrons flow? A.from reducing agent to oxidizing agent Names for locations of each half-rxn anode: oxidation cathode: reduction

Car Battery: Lead storage battery

Cell Potential cell driving force on electron to move them through the wire also called Electromotive Force (emf) units are Volts (V) 1 V = 1 Joule/Coulombs of charge skip the last paragraph in 17.1- too much physics!

17.2 Standard Reduction Potentials Ch. 17: Electrochemistry 17.2 Standard Reduction Potentials

Standard Reduction Potentials we assign  values to each half-reactions we can find the total cell by summing the individual potentials for the combination of half-reactions they are always written as a reduction process so must switch one of them Always change the sign of ° when you reverse the direction do not multiply the ° by an integer used to balance the equation- it is intensive

Standard Reduction Potentials ° values for reduction half-reactions with all solutions having 1 M conc and all gases 1 atm

Standard Hydrogen Electrode Pt electrode in contact with 1 M H+ and H2(g) at 1 atm assigned an  of zero can calculate others by pairing with this and measuring total cell 2H+(aq) + 2e-(aq)  H2(g)

Standard Hydrogen Electrode

Example Find the °cell for the reaction Fe3+(aq) + Cu(s)  Cu2+(aq) + Fe2+(aq) half reactions: Fe3+(aq) + e-  Fe2+(aq) °=0.77 V Cu2+(aq) + 2e-  Cu(s) °=0.34 V 2nd one must be reversed °cell = -0.34V + 0.77V = 0.43V

Line Notation short hand for describing cells anode is on the L and cathode is on R separate anode and cathode with || separate phases in one half-rxn with | if no part of a half-rxn is solid, use Pt electrode electrodes go on far ends of notation Mg(s) | Mg2+(aq) || Al3+(aq) | Al(s) Pt(s) | ClO3-(aq), ClO4-(aq) || MnO4-(aq), Mn2+(aq) | Pt(s)

Galvanic Cells run spontaneously in the direction where °cell is + Describe a Galvanic Cell: balanced chemical eq. (make sure °cell +) give the direction of electron flow assign the anode and cathode give line notation

Example Write balanced equation Describe the Galvanic Cell based on the following half-reactions: Ag+ + e-  Ag °=0.80 V Fe3+ + e-  Fe2+ °=0.77 V Write balanced equation Fe eq. must be reversed Ag+ + Fe2+  Ag + Fe3+ °cell = 0.03V

Example Assign cathode and anode Give the direction of electron flow oxidation: Fe2+  Fe3+ +e- reduction: Ag+ +e-  Ag electrons flow from Fe2+ compartment to Ag+ compartment Assign cathode and anode anode: oxidation: Fe2+  Fe3+ +e- cathode: reduction: Ag+ +e-  Ag

Line Notation Pt(s) | Fe2+(aq), Fe3+(aq) || Ag+(aq) | Ag(s)

Alkaline Battery

Sacrificial Metal to prevent rust on large structure

Fuel Cell (Hydrogen fuel)