Electrochemistry Lab Report

Measurement of Physical Properties

 

In any measurement it is important to know the precision of the measurement and also its accuracy. All physical measurements should be made as precisely and accurately as possible. Maximizing the precision of a measurement is accomplished by using the most precise equipment available, and using it properly. If possible it is also wise to compare your result with either a known or theoretical value.

 

Significant Figures

 

When calculating a result from more than one measurement it important to retain the uncertainty information from all the measurements. There is an entire field of mathematics devoted to this topic. In this course we use the relatively simple method of significant figures. A summary of the rules with examples:

 

Addition and subtraction: line up the numbers to be added or subtracted; the answer is truncated to the decimal place of the least precise number.

 

Ex. 12.1 + 2.345 = 14.4

 

15.678 – 2.2 = 13.5 (notice I rounded up)

 

Multiplication and Division: Significant Figures in the answer are equal to the number of significant figures in the least precise number.

 

15.6 x 2.1 = 31

 

16.789 ⎟ 25.67432 = 0.65392

 

25.1 x 3.00 = 75.0

 

Note zeroes before another number as in 0.65392 do not count. In the middle and the end they count.

 

 

Laboratory Notebooks

 

In this course you will be required to keep a laboratory notebook. A good laboratory notebook is an accurate record of everything, which occurred in the lab. In patent disputes a good lab book versus an inaccurate lab book can mean millions of dollars. In this course it may mean hundreds of points. Before the lab you will be required to prepare a lab report outline to be completed during the lab session. Each lab report will contain:

 

Title and Purpose

  1. Procedure and Observations
  2. Data and Calculations
  3. Results and Conclusions
  4. Answers to questions in manual

 

 

An example of a lab report is shown below:

 

1. Title: Density of liquid and a solid.

Purpose to measure the density of a liquid and an unknown solid.

 

2. Procedure:

 

Part 1 Liquid

 

Weigh an empty 10.0 mL volumetric flask Fill with unknown liquid.

Weigh filled volumetric flask

 

Part 2 Solid

 

Fill a graduated cylinder with about 25 mL of water

Measure precisely volume of water.

 

weigh dry solid sample

Observations

 

Unknown # 5 smells like gasoline

 

Mass of empty flask = 12.032 grams

 

Mass of full flask = 18.685 grams

 

Part 2 Unknown #12 Shiny orange color

 

 

 

Volume of water = 24.83 mL

Volume of water + metal = 28.53 mL

Mass of Dry metal = 46.409 grams

 

 

Data and Calculations:

 

Part 1 Liquid:

 

Density = Mass / volume

Mass of liquid = mass of liquid + flask – mass of flask = 18.685 grams – 12.032 grams = 6.652 grams.

Density = 6.652 grams / 10.000 mL = 0.6652 g/mL

 

Part 2 solid

volume of solid = volume of solid + water – volume water = 28.53 mL – 24.83 mL = 3.70 mL density = 46.409 g/3.70 mL = 12.5 g/mL

 

Results and Conclusions:

 

The density of the liquid was determined to be 0.6652 g/mL by comparison with the density table in the CRC it appears the sample could be hexane, which has a density of 0.660 g/mL

 

The density of solid was 12.5 grams / mL. The solid looked like copper, but the density of copper from the CRC is: 8.94 g/mL, which is significantly less than my unknown sample. Therefore although the sample looks like copper it must be something else.

 

 

In this example, the data is recorded in the section with the observations, and the procedure is recorded in one column and the observations are recorded in an adjoining column. This allows you to record your observations with the correct section of the procedure. In some experiments, the type and volume of data is better recorded in a table. In this case it should follow the procedure section. You should still leave room in the procedure section for observations. One of the objectives of this course is for students to learn how to determine what data they need to collect, and how to organize it. For some experiments explicit instructions for organizing the data and calculations will be given, but for other experiments you will need to determine this for yourself before class. In the case of repetitive calculations tables are necessary. A spreadsheet such as Excel can be used, and instructions are included for the Reaction Rate experiment. All your calculations must follow the rules for significant figures and every value must have a correct unit. A spreadsheet or calculator will not determine the correct number of significant figures; it is up to you.

When determining the results and conclusions, there are some things to keep in mind. The results should relate back to the purpose. Address directly if the purpose was fulfilled. If the result is a number clearly restate what it is and the unit for the number. If possible compare your result with a literature value. If you received no result or an unexpected result, give some scientific explanation of this. Human error is not a good explanation, because the experiment or section, which was in error, should be repeated. Thoroughness is important but it is not necessary to write everything you know about density or volume etc.

To be ready to use all the lab time efficiently, before lab class you should have completed the purpose, procedure and arranged the data table or written down what you need to measure.

 

Lab Instructors may have additional report requirements.

 

 

Lab 12 Introduction to Electrochemistry

 

Introduction

Electrochemistry refers to a branch of chemistry that measures oxidationreduction reactions at the interface of a solution and an electrode (an electrified metal). Electrochemical measurements can provide a wealth of information including stoichiometry, rates of reaction and mass transfer, and even equilibrium constants. To establish electrical current an electrochemical cell must have three main components, shown in the image below:

  1. Electrodes that are connected externally by means of a metal conductor
  2. Two electrolyte solutions with electrical contact that permit movements of ions from one to the other often through the use of a salt bridge
  3. An electron transfer reaction (anode – oxidation and cathode – reduction, respectively) that occur at each of the two electrodes

 

word image 34

 

Today’s experiment has been broken down into two demonstrative experiments. In Part 1 you will characterize a very simple electrochemical system, in which you will observe the electrolysis (electrical break down) of water at the interface of graphite electrodes.

 

𝐻%𝑂→𝐻l+𝑂𝐻9

 

In Part 2 you will create electrochemical cells to measure the electrical potential junctions (Ecell) of copper and zinc electrodes within their respective salt solutions to determine E0 using the Nernst Equation. A concentration cell will then be formed, in which the electrochemical cell makes use of an imbalance of ionized species on different sides of a cell to provide an electromotive force. That force (measured in volts) is an attempt to equalize the concentration, which means is will get smaller as the system gets closer and closer to equilibrium.

 

o 0.0592V logQ Ecell =E

n

 

The equation above is the Nernst equation, which is used to determine the Ecell for a system that is not in equilibrium. Here (n) represents the number of electrons being exchanged in a reaction.

 

Ag s( )→Ag+ +1e n = 1

Cu( )s Cu2+ +2e n = 2

 

Additionally Q is the reaction quotient, which is represented by

 

Q=[Ag++]2

[Ag ]1

 

If [Ag+]2 > [Ag+]1 then the log Q term will be negative, giving us a positive value for Ecell. In cases when an insoluble solid is produced within the electrochemical cell, the Q can be informative of the solubility constant, Ksp. You then measure the solubility constant of an insoluble salt within your electrochemical cell.

Procedure Part 1

  1. Obtain two pieces of pencil lead from your instructor. Each piece should be roughly 3 cm in length.
  2. Add 20 ml of water to the petri dish
  3. Add 2 ml (~40 drops) of universal indicator to the 20 ml of water.
  4. Add 10 ml of 1 M Sodium Sulfate solution to the 20 ml of water.
  5. Connect the 9V battery to the battery connector. Then connect the alligator clip to one piece of the pencil lead to create your graphite electrodes.
  6. Without allowing the electrodes to touch one another, place the graphite electrodes in the petri dish of water. Electrolysis should occur immediately. Record observations including the color of the overall solution and the changes occurring at each of the electrodes.

Questions

  1. Describe the reaction that you are observing. Why are we observing this color change? (Hint: How is an indicator used to detect changes in pH?)
  2. Write the oxidation and reduction half-reactions and indicate which is happening at the anode and which is happening at the cathode.
  3. Now write the overall balanced net ionic reaction for the electrolysis of water. (Hint: Balance the half-reactions by adding OH and H+.)
  4. Based on the structure of water (H2O) which electrode would you expect to produce the most gas? Did you observe this result? Why?

Part 2 A – Determine the E0 for a Cu and a Zn Cell

  1. In a small beaker, obtain roughly 30 mL of 0.1 M CuSO solution. In another small beaker, obtain roughly 30 mL of 0.10 M Zn(NO3)2.
  2. You will need to make two different galvanic cells with each of these solutions. To prepare your first, use a piece of the metal in the solution, connect the leads of the digital multimeters to the metals so that you get a positive voltage. Place a salt bridge (KNO3 soaked paper) between the two beakers.
  3. Remove the metal electrodes, rub them with steel wool and re-immerse them in the solution, record another reading. Do this one more time so you have three trials.

B- Prepare and Test Two Concentration Cells

  1. Set up a concentration cell using one of your three metals. Obtain about 25 mL of 1M solution. Make a 1:20 dilution of this metal (using a graduated cylinder). Using the metal electrodes, digital multimeter, and salt bridge to set up a concentration cell.
  2. Remove the metal electrodes, rub them with steel wool and re-immerse them in the solution, record another reading. Do this one more time so you have three trials.

– Ksp of an Insoluble Salt

  1. Set up a concentration cell using copper. One cell should have a solution with a concentration of 1.0 M CuSO4 (about 25 mL). The other should have a concentration of 0.05 M of CuSO4 (about 20 mL). Prepare this through dilution if necessary or use one of the previous set ups.
  2. Mix 10 mL of 0.05 M NaOH with the beaker that contains the 0.05 M CuSO4. This will form an insoluble Cu(OH)2 salt.
  3. Insert copper metal and a salt bridge to measure the potential between the two. Remove the metal, rub it with steel wool, and re-immerse it. Do this to get three trials.
  4. Disassemble your cell, and pour all solutions into the waste containers.

Calculations

  1. The average Ecell of a Cu/Zn cell.
  2. The average Ecell for the copper concentration cell with sodium hydroxide.
  3. Determine the concentration of copper in the concentration cell.
  4. Calculate the Ksp of Copper (II) hydroxide. Results and Discussion
  5. Discuss how your experimentally determined E0 values for each galvanic cell compare to the expected values. Explain any difference between the two.
  6. Discuss how your experimentally determined E0 values for each concentration cell compare to the expected values. Explain any difference between the two.
  7. Explain any difference between your Ksp result and the expected result.

Portions of this document have been adapted from the following sources.

  • M. Davis, Chemistry 203: Lab 6a – Electrochemistry http://faculty.ccc.edu/mdavis/Documents_common_2010_Fa/Labs/Lab_6a_el ectrochem_variety.pdf, Accessed April 15, 2014.
  • M. Davis, Introduction to Electrochemistry A Two-Week Chemistry Experiment : Week One – Concentration Cells,

http://faculty.ccc.edu/mdavis/Documents_common_2010_Fa/Labs/lab_6_ech em.doc, Accessed February 10, 2014.

  • D. A. Skoog, Principles of instrumental analysis, Saunders College Pub.; Harcourt Brace College Publishers, Philadelphia; Orlando, Fla., 1998.

 

 

 

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