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Voltaic Cell Lab Assignment

Presentation on theme: "Chapter A2 2.4 – Voltaic Cells."— Presentation transcript:

1 Chapter A22.4 – Voltaic Cells

2 Voltaic cells The rest of the chapter focuses on WHY this occurs.
the focus on metals so far been been what happening on the surface of the metaltarnishing/ rusting, (aka color change), precipitate formation or bubbles/ gas formationThe rest of the chapter focuses on WHY this occurs.a focus on the movement of electrons between the two metals

3 Voltaic cellsthough commonly a voltaic cell is referred to as a “battery”, technically cells are only referred to as a battery when several are togetherwhen an electronic device is operating, voltaic cells provide a continuous flow (current) of electrons, which is converted into current to power the device

4 Voltaic cells the voltaic cell that you know looks like this:
this is the version of the voltaic cell we will make in the lab.

5 Voltaic cellsan electrode is a solid piece of metal that is suspended in a solution (of the same metal ions as the electrode) and connected to an external circuit.

6 Voltaic cellsthe electrode zinc, is immersed into an electrolyte solution, where the zinc electrode acquires an excess of electrons, becoming negatively charged

7 Voltaic cellsthe other electrode is usually composed of a different material (copper) and will become positively charged

8 Voltaic cellsonce a circuit is closed between the two electrodes, the electrons will repel from the negative zinc electrode, pass through the circuit and flow through to the positive electrode

9 Voltaic cellsthe reaction will continue until the negative electrode can no longer be supply electrons.

10 Voltaic cells

11 Voltaic cellsa salt bridge is a glass U-shaped tube that is filled with an ionic solutionthis is to allow for free flow of electrons from one solution to the other

12 How the cell works:because it is the more reactive of the two metals, the zinc electrode will become oxidized, (lose electrons)these electrons will travel from the electrode, through a metal wire, and then into an electronic devicethe device; a voltmeter, measures the quantity of electrons passing through it (= amount of electricity)

13 How the cell works:the electrons will pass through the device, back through another wire, into the copper electrodethese electrons will be attracted to the Cu2+(aq) ions in the solution and they will be reduced.over time, the zinc electrode shrinks in size (as Zn  Zn2+) and the copper electrode grows (Cu2+  Cu)

14 How the cell works:if the two solutions were NOT connected, the zinc would run out of electrons and the cell would stop workingthe solution in the salt bridge allows a continuous flow of electrons back into the zinc solution

15 How the cell works:

16 Analyzing a voltaic cell:
Step #1: identify the electrode where oxidation occurslocate the two metals on the activity series (right side)the metal closer to the BOTTOM will be OXIDIZED = reducing agentthe electrode that is oxidized is called the anodethe other electrode is reduced, and is called the cathode

17 Analyzing a voltaic cell:
Step #2: describe the oxidation process in the anodewrite the oxidation half-reactionEg. Pb(s)  Pb2+(aq) +2e-electrons leave the anode and travel to the external circuit running the electronic devicethe voltmeter measures the quantity of electrons (amount of electricity) being producedbecause the anode is the electrode where the electrons originate, it is considered the negative electrode

18 Analyzing a voltaic cell:
Step #3: describe the reduction process in the anodethe electrons travel through the voltmeter and into the cathodethe electrons are attracted to the positively-charge metal ions in the cathode solutionthe cathode ions will unite with the electrons and form a solid metal, which is deposited on the electrodeEg. Ag+ (aq) + e-  Ag (s)

19 Analyzing a voltaic cell:
Step #4: describe how the salt bridge completes the circuitall electrical circuits require a complete circuit in order to function.the salt bridge connects the cathode back to the anode to replenish the electrons on the anode sidethe salt bridge contains a third ionic solutionthe positive ions from the salt bridge solution will be attracted to the cathode, while the negative ions from the salt bridge solution will migrate toward the anode.

20 Voltaic Cell- Example in this voltaic cell: anode cathode KCl
zinc is the _________ – it is oxidizedcopper is the __________– it is reducedthe solution in the salt bridge is _____________ (aq)chloride ions are a spectator ion – their job is to replenish the electron supply at the anodeanodecathodeKCl

21 Cell Notation anode salt bridge cathode
voltaic cells can also be represented using short hand cell notationZn(s) | Zn2+(aq) || Cu2+(aq) | Cu(s)anode salt bridge cathode oxidation reductionthe anode is listed on the left, the cathode on the right (think alphabetical order)the vertical line | represents a boundary between a metal and its solutionthe double line || represents the salt bridge

22 Practice Problem #1:a) Draw a voltaic cell using the following supplies:two beakersU-tube & cotton ballswire & voltmetertin and magnesium stripssolutions of SnSO4(aq), MgSO4(aq), and NaNO3(aq)b) Label the direction of e- flow, the anode, cathode, OA, RA, - and + electrodes, voltmeter and salt bridgec) Write the short hand cell notation

23 Practice Problem #1 (Solution):

24 Assignment: Practice problem (page 87) Practice problems (page 91)
34Practice problems (page 91)37 & 39Practice problem (page 92)402.4 Summary 9page 94)Q’s 2, 3 & 6 pg 94

25 Homework Check-up Zn(s) / Zn2+(aq) // Ni2+ (aq) / Ni(s)
Use the cell notation below to answer the following questions. Assume the salt bridge contains a solution of potassium nitrate (KNO3(aq) )Zn(s) / Zn2+(aq) // Ni2+ (aq) / Ni(s)Identify which metal will be oxidized and which would be reducedIdentify the anode and the cathodeWrite the half reactions that occur at each electrodeDraw the voltaic cell. Label the direction of the electron flow and the anions of the salt bridge.Homework Check-up

26 Zn(s) / Zn2+(aq) // Ni2+ (aq) / Ni(s)
According to the activity series, zinc is the more reactive metal- so Zn(s) is oxidized and the Ni2+(aq) are reduced.Anode = Oxidation= zinc metal Cathode = Reduction = nickel metalOxidation: Zn(s) Zn2+(aq) + 2e–Reduction: Ni2+(aq) + 2e–  Ni(s)Identify which metal will be oxidized and which would be reduced (1 mark)Identify the anode and the cathode (1 mark)Write the half reactions that occur at each electrode

27 Zn(s) / Zn2+(aq) // Ni2+ (aq) / Ni(s)
d) Draw the voltaic cell (2 marks) Label the direction of the electron flow (1 mark) and the anions of the salt bridge. (1 mark)Cell must have 2 beakers, (one with Zn(s) and Zn2+(aq), one with Ni(s) and Ni 2+ (aq)), connected by wires to a voltmeter, and a salt bridgeElectrons must leave the Zn anode and flow towards the Ni cathodeThe negative NO3- anions are attracted to the anode. The positive K+ cations are attracted to the cathode(think alphabetical…Anion-Anode. Cation= Cathode)

28 Assignment:Complete the pre-lab assignment for the Voltaic Cells LabYour pre-lab MUST BE done, in order to participate in the lab!??

29 Chapter A22.5 – Electrolytic Cells

30 Electrolytic vs. Voltaic
an electrolytic cell is a system where a non-spontaneous redox reaction is forced to occura reaction that is non-spontaneous will only occur if energy is addedin an electrolytic cell, energy is added in the form of electricityVoltaicElectrolyticspontaneous?yesnorequires energy?produces voltage?useenergy sourceelectroplatingchange in energyexothermicendothermic

31 Electroplatingmetals, like gold and silver, that are the most stable and corrosion-resistant are also the most expensiveto manufacture a metal object that is resistant to corrosion it would NOT be cost-effective to make the whole thing out of goldinstead, a thin coating of gold is applied to the surface of a more affordable metal

32 Electroplatingthe object to be coated is submerged in a solution of the metal ions (e.g. silver ions for objects that are to be coated in silver metal)an external energy source (a battery) supplies energy forcing electrons to flow into the objectthe negatively charged electrons will attract the positively charged metal ions from the solution, and turn them back into metal atoms, which will accumulate on the surface of the object to be plated.

33 Electrolytic cells Step #1:
electrons from the plating (the expensive) metal cathode are attracted to the + electrode of the power sourceby removing electrons from the metal atoms, ions are formed and added into the solution

34 Electrolytic cells Step #2: Step #3:
once removed from the metal, the free electrons flow through the power source.Step #3:electrons are forced out of the - end of the power source and accumulate on the surface of the object to be plated

35 Electrolytic cells Step #4:
positive Au+(aq) from the solution are attracted to the negative electrons in the object to be platedthe Au+(aq)ions gain the electrons, and turn back into solid gold coating the object.

36 Electroplatingelectroplating is a good way to protect metals that are easily oxidized, like ironmetals that work as good electroplaters (coatings) are chromium (aka chrome), platinum, silver and gold

37 Gold jewelry solid gold gold plated
two types of gold jewelry exist - that which is made out of solid gold, and that which is gold platedsolid goldkarats - pure gold is 24Kgold is a soft metal, so it is often combined with other metals like brass (copper and zinc) and nickel to make it more durablethe number of karats in the gold refers to how many 1/24th of gold it containsgold platedif you have a piece of gold plated jewelry, care must be taken to avoid any deep scratchesdeep scratches will expose the oxidizable metal underneath

38 Other uses for electrolytic cells
refining metalsa sample of impure metal (anode), pure metal (cathode)ions of the pure metal will travel from the anode to the cathode to build up the atoms of pure metalelectrolysisdecomposition of a compound by means of an electric currente.g. electrolysis of water makes it decompose into O2 and H2

39 Other uses for electrolytic cells
producing non-metalsnon-metals, especially the halogens, are difficult to obtain in pure form because they are so reactivenon-metal atoms will accumulate around the anode of an electrolytic cellrecharging voltaic cellswhen you use a battery recharger, you are using an electrolytic cell to reverse the process that occurs normally in the voltaic cellyou are literally re-charging the voltaic cell with a new supply of electrons

40 Assignment:Complete the Voltaic & Electrolytic cells Worksheet

41 Assignment:Prepare for your Chapter A2 Exam

 

Abstract:

In this experiment, a standard table of reduction potentials of a series of metal ions is constructed usingcopper, iron, lead, magnesium, silver, and zinc. These half cells are are connected by a salt bridge and allpotentials are measured with respect to the zinc electrode. Also, the measured voltage of a nonstandardcopper cell is calculated through the Nernst equation. The solubility product constant of AgCl is alsodetermined through the Nernst equation. The

sp

 value for AgCl was determined to be 7

.

33

 ×

 10

11

,yielding a percent error of 59.3%. The voltage for the cell reaction was experimentally determined to be 0.81V.

Theory:

An

 electrochemical cell

 is produced when a redox reaction occurs. The resulting electron transferbetween the reaction runs through an external wire. Because the oxidation and reduction reactions arephysically separated from each other, these are called

 half-cell reactions

. A half cell is prepared fromcontact with the metal with its solution of ions. Each element’s unique electron configuration allows each todevelop a different electrical potential.The

 standard reduction potential

 is the voltage that a half-cell, under standard conditions (1 M, 1atm, and 25

C), develops when combined with the standard hydrogen electrode, that is arbitrarily assignedto a potential of zero volts. A positive

 E 

cell

 value indicates that the redox reaction in that particular cell is

spontaneous

.Calculations of nonstandard potentials can be made using the

 Nernst Equation

:

=

RT nF 

ln(

Q

) (1)where

 E 

 is the measured cell potential,

 E 

is the standard cell potential,

 R

 is the gas constant (8.314 J/mol

·

K),

, is the temperature in K,

n

 is the number of moles of electrons transferred as shown by the redoxreaction, and

 is the

 Faraday constant

 (9

.

65

 ×

 10

4

C/mol).At STP, the Nernst equation can be simplified to

=

 0

.

0592

n

 log(

Q

) (2)

Procedure:

A wellplate is set up such that the first row contains approximately 2 mL of 1.0 M Zn(NO

3

)

2

 solutionin each well. In the second row, two mL of Cu(NO

3

)

2

, AgNO

3

, Fe(NO

3

)

3

, Mg(NO

3

)

2

, and Pb(NO

3

)

2

areadded in their respective wells. A salt bridge is used to connect to the two adjacent wells (made from filterpaper soaked in KNO

3

 solution. A voltmeter is used to measure the potential difference for each of the 5half cells. After, measure the potential difference between at least six combinations of various electrodes.Again, use the voltmeter to measure the potential difference.In part 2, Cu(NO

3

)

2

 is diluted to 0.0010 M. It is added onto a wellplate and measured against thestandard zinc half-cell.In part 3, 10 mL of 1.0 M NaCl solution is mixed with one drop of 1.0 M AgNO

3

. After precipitationoccurs, some of the solution is poured into the well plate and measured against the standard zinc half-cell.1

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