ElectroChemistry

Objectives Galvanic versus voltaic Standard potential Use of the Nernst equation Water electrolysis

Background Electrochemistry is all around us- Batteries Understand how they work Principles to increase the technology

batteries Conversion of the chemical energy in the solution to useful work-electrical energy e- are transferred from one chemical species to another via the oxidation reduction process

Galvanic cell Simple battery A battery that leaked –appears a salt on the outside of the battery- This is the solution that makes the battery Driven by redox reaction, in a closed system, when the e- are all converted then the battery dies.

Galvanic cell Cu+2 + Zn  Zn+2 +Cu This is the overall reaction Zn  Zn e occurs on the anode Cu+2 + 2e  Cu occurs on the cathode These are the two half cell reaction

The process Zinc losses 2 electrons (oxidized) which are attracted to the Cu ions and converts the Cu+2 to copper wire (reduced). If the e- is captured before they can reduce the Cu then work can be done Captured in a salt bridge

Circuits Salt bridges are Na and other negative ions NO3-. When present the ions enter the solution to balance the charges as the Zn ions are formed and the Cu ions are depleted.

Galvanic cells 200 BC to plate silver with gold Credited to Alessandro Volta in 1800 first commercial battery Daniel Cell –Used today in door bells

Standard cells Potential of the cell is compared to the half cell reaction H+ + 2e  H2 This cell is potential is 0.00V = E˚ Standard Hydrogen electrode = SHE To compare both solution are at 1 M All half cells are written as reductions. Table of reduction potentials are in many texts

Half-reaction E° (V) Half-reaction Mg e− Mg(s) −2.38Mg Al e− Al(s) −1.66Al Mn e− Mn(s) −1.18Mn Zn e− Zn(s) −0.76Zn Fe e− Fe(s) −0.44Fe Pb e− Pb(s) −0.13Pb 2 H+ + 2 e− H2(g) ≡ 0 Cu2+ + e− Cu Cu Cu e− Cu(s) Cu+ + e− Cu(s) Fe3+ + e− Fe Pb e− Pb Au+ + e− Au(s) Ag2+ + e− Ag

Half cells Zn e− Zn(s) −0.76Zn Cu e− Cu(s) Cu e−  Cu(s)reduced Zn(s)  Zn e− oxidized +0.76Zn Rapid spontaneous reaction

Nernst Equation Walter Nernst in 1800 worked on cell Thermodynamics E = E˚ - RT ln Q nF E˚ = standard cell potential - from the table E = actual cell potential- this what you are to calculate T = temperature (Kelvin) – R = constant 8.314J/Moldeg F = Faraday constant (96,483 coulombs per mol) N = number of electrons transferred in the balanced equations. Q reaction quotient (products over reactants)

Electrolysis cell Electrolysis of water Uses a battery to move the electrons in opposite directions Charles Hall discovered this while in college and set out to use it for Al ore processing.

Electrolyisis In the water at the negatively charged cathode, a reduction reaction takes place, with electrons (e − ) from the cathode being given to hydrogen cations to form hydrogen gas (the half reaction balanced with acid):reduction Cathode (reduction): 2H + (aq) + 2e − → H 2 (g) H At the positively charged anode, an oxidation reaction occurs, generating oxygen gas and giving electrons to the cathode to complete the circuit:oxidation Anode (oxidation): 2H 2 O(l) → O 2 (g) + 4H + (aq) + 4e − H 2 O H The same half reactions can also be balanced with base as listed below. Not all half reactions must be balanced with acid or base. Many do like the oxidation or reduction of water listed here. To add half reactions they must both be balanced with either acid or base. Cathode (reduction): 2H 2 O(l) + 2e − → H 2 (g) + 2OH − (aq) H 2 OHOH Anode (oxidation): 4OH − (aq) → O 2 (g) + 2H 2 O(l) + 4e − OHOH 2 O Combining either half reaction pair yields the same overall decomposition of water into oxygen and hydrogen: Overall reaction: 2H 2 O(l) → 2H 2 (g) + O 2 (g)H 2 OHO The number of hydrogen molecules produced is thus twice the number of oxygen molecules. Assuming equal temperature and pressure for both gases, the produced hydrogen gas has therefore twice the volume of the produced oxygen gas. The number of electrons pushed through the water is twice the number of generated hydrogen molecules and four times the number of generated oxygen molecules.

Procedure Will need filter papers cut 5 cm by 1 cm 2 pieces of Cu One each al and zn These are your electrodes Make sure they are clean and free of oxidation

Prepare the Couples

Salts Add the appropriate salt under each electrode Add the salt bridge last- make sure the salt bridge reaches the electrode. Measure the potential of the both cell

labworks Open and click on calibrate Scroll to V- the units are mV Then design First line as is Second line timer Third line scroll to V Click aquire and then start

Lab works After clicking acquire plug in the probes as shown by the instructor Red is red and purple is for negative – black probe.

Measuring the V

Precautions Do not allow the electrodes to come in contact with each other Keep even constant pressure on the electrodes.

Salt solutions THESE ARE THE CELLS IN YOUR INSTRUCTION Create them on one(1) filer paper on the document protector, to keep the salt off the desk. Total of 4 cells 0.1 M Cu and 0.1M Zn 0.1 M Cu and 0.1 M Al Then ALSO do 1.0 M Cu and 0.1 M Zn 1.0 m Cu and 1.0 M Zn

Do this first Electrolysis of water Graphite sample cup Alligator clips with 9 V cap and tape 9 v battery NaNO3 solution 2 drops Water DI Use Yamada indicator 2 drop

procedure Put Di water in the cup ¾ full 2 drops indicator 5 drops NaNO3 Attach the graphite to the clips and tape them to the side of the cup as shown Do not let the clips enter the water and the electrodes touch the sides of the cup

Procedure Should see bubbles at each electrode Color at each electrode If this is done first Then left to one side while the other experiment is done The graphite will be shredded, Why

Report Your post lab report, prelab and the extra information on the two extra cells. Correction on the pre lab this is what it should read Cu/Zn Cu/Mg Measured potential V Oxidation half cell calculated potential V Reduction half cell