Q1 : Week 4 – 9/25 – 9/29

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LAB 1 – Gerber Voltaic Cell is LATE!!!!!  

LAB 2 – Voltaic Cell design is due Tuesday 9/26

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9/26 – Tuesday – B Day – 2/3b Lab, 4 

Main focus –   

a) To Review the concepts of the electrolytic cell with an electroplating demo and electrolysis demo.

b) To take Notes on the 2 types of Electrolytic Cells.      

Textbook: 17.7  – Electrolysis

Period 2,4:

Period 3b:

1) Setup the Electrolysis of copper (II) chloride lab – 

2) Michael Faraday and his experimental finding with electrolytic cells:

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1.  Complete Stoichiometry 8 redox electrolytic – 1999.pdf worksheet SIDE 1 One ONLY with

        the lecture below that reviews question below!

 

      Please complete side 1 of this worksheet the following ways:

 

         a) You need to be able to draw and label the electrolytic cell that the question describes just like you                    did from the diagram of the demo’s today.  3 hints below will help!

 

        * Please remember that you are heating the AlCl₃ (s) until the ions are free.  Based on the chemical              formula you should be able to figure out what ions are present.  Also this goal of the electrolytic cell is make the elemental form of each ion.  We are looking to make pure Al  and Chlorine gas (Cl₂).

 

         *Remember that you will include a battery (batt) in your diagram and depending  where you draw the anode/cathode of the Battery (voltaic cell) and how you connect it to the electrodes (that do not react) in the electrolytic cell will determine which  electrode is the anode/cathode in the electrolytic cell. There are many correct ways to diagram this.

 

         *But the ions will move only to the correct electrode based on their charges!

 

          b) The second part, determining how many grams was deposited on the cathode is tonight’s  

               stretch. Please read the Notes on Electrolytic Cell below.

 

          c) Try making the calculation using the notes, the key, or the lecture.  I suggest that you try it  

               after reading the notes and then view my lecture below to see how I did the question.

                You could use the key below or just let me walk you through the problem. 

 

Do not worry about the calculation in the second part if it gives you trouble. I will go over and reteach that in class tomorrow.                                                                                                                                                                                                    

Stoichiometry 8 redox electrolytic – 1999.pdf

View Download

 

Stoichiometry 8 redox electrolytic KEY – 1999.pdf
View Download

   

NOTES on the Electrolytic cells:

These electrolytic problems are centered around a couple of concepts:

 

1. Coulomb = SI unit of charge

2. Faraday’s Constant

 

If the current in a circuit is 1 Ampere then 1 coulomb of charge passes a point in the circuit every second.

 

1.0 A = 1  C/sec   (1 Amp = 1 Coulomb per second)

 

Abbreviation:  C as in Amperes above

Equation symbol: q or Q (as in Coulombs Law)

 

q = -1.60 x 10^-19 C in a single electron (the charge in a single electron = -1.60 x 10^-19 C )

 

There are also 1.60 x 10¹⁸ electrons in 1 Coulomb

 

We use the concept of Coulomb to connect Amperes which measures current (or measure amount of the flow of electrons) that is being pushed through the conductor (or wire) by a voltage (Energy). A measures amperage over a measures amount of time will give us the AMOUNT OF CHARGE that passes over a point in the wire IN THAT TIME INTERVAL.

 

So Amperage (A = C/sec ) describes TOTAL CHARGE IN A TIME INTERVAL!  So the key is the TIME and  measured Amperage will give us total charge.

 

Now once we have TOTAL CHARGE we can relate Faraday’s Constant which came from Faraday’s laws of electrolysis that *basically states:

 

1: The mass of the deposited metal is directly proportional to the quantity of electrical charge that passes through the electrolyte.

 

2: For the same quantity of electricity or charge passes through different electrolytes, the mass of the deposited chemical is directly proportional to its chemical equivalent and inversely proportional to its valency.

 

Now at the time of Michael Faraday there was very little support for the idea of atoms , atomic or moles.  Just chemical properties. One property that was well know was the valency of elements.  Copper has a valency 2 because they knew 2 Cl ions would bind with Cu.  This valency of course is valence electrons and is related to the 

ELECTRONS IN THE HALF REACTION!!!!

 

                                                                                 Cu⁺² + 2e^– —->   Cu (s)

 

There was no understanding of mole concept yet so Faraday could not relate these observations in an electrolytic cell to moles.  Once the mole equivalent was determined in the turn of the Century (1905) could a constant be described that relates charge of electrons with a mole of electrons.  Faraday set the stage and the mole concept closed the door.  Faraday would dance aggressively happy is he saw what we can predict with this constant given in his name:

 

                                   F (Fancy) = 96,500 C / 1 mole of electrons

 

SO we can Now link the Total charge we get from measures Amperes and Time to the Quantity of electrons that are being forced through conductor. Since these electrons are being used at each electrode to “Force” an non-spontaneous reaction we can use the ratio of electrons that are responsible for the Coulombs to determine the number of atoms that are being produced at either electrode.  

 

In the case of Copper above there are 2e^– per 1 Cu⁺² in the half reaction above so 

if there will be twice as many electrons needed as Cu atoms deposited.  

 

Faraday’s constant the total charge (we get from Amperage and time) to total number of electrons in half reactions occurring at the electrodes. Using the ration of electrons to atoms produced in the half reaction we can determine the amount of atoms.                                                                                                                                                

Example:  How many grams of Cu are deposited on the cathode in an electrolytic cell if the cell runs for 90 minutes at an amperage of 3.0 A?

 

                          

                                                                                                                                                                                                                               Jump to:  Wednesday Homework / top ______________________________________________                                                                                             9/27 – Wednesday – A Day – 2/3a Lab, 4 

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Main focus –

a) To experimentally determine the mass of copper deposited at the cathode of an electrolytic cell from an aqueous solution of copper (II) Chloride.

b) To calculate the theoretical mass of copper that would be expected.

Period 2 – Lab today

1.  Complete the experimental data gathering by performing a hot filtration of the copper metal from your electrolytic cell solution.

a) mass out your filter paper with name of group on the paper.

b) Hot filtration –

c) Place in oven – We will measure the dry copper and filter paper Next Wednesday!

d) Calculate the expected or theoretical mass of copper that should be expected,

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