Many chemical reactions reach a state of equilibrium if conditions are right. In an equilibrium system, forward and reverse reactions occur at equal rates so that no net change is produced. When equilibrium is reached by a reaction in a test tube, it appears that changes have stopped in the tube. Once equilbrium has been reached, is it possible to produce further observable changes in the tube? If so, can you control the kinds of changes? If not, why are further observable changes impossible? You will observe several chemical systems in this laboratory activity. A careful study of your observations will enable you to answer these questions.
Obtain a test tube rack, six small (13 x 100 mm) test tubes that are clean but don't have to be dry, and a test tube clamp. The test tubes should be placed open end up in the test tube rack. Prepare a hot water bath: Half-fill a 250 mL beaker with tap water. Start to heat the water (as your teacher directs) so that the water will be near boiling when you are ready to use it. Prepare an ice water bath: Fill a 250 mL beaker with crushed ice. Add enough tap water to make "slush". Set up a data table with column headings as indicated below (The last column will be completed after data have been collected.) System Disturbance Observed Change Direction of Shift 1 2
etc. As you set up equilibrium systems and add disturbances to them in the procedure, enter appropriate information in each of the first three columns of your data table. Mix chemicals in test tubes by holding the top of the tube with one hand while you flick the bottom of the tube with your other hand until the tube contents.
System 1: Iron(III) and thiocyanate
Setting Up the Equilibrium
Half-fill the first tube in your rack with distilled water. Add two drops of 0.1 M Fe(NO3)3 and two drops of 0.1 M KSCN to this tube. Mix the contents thoroughly. If the contents of the tube are not red-orange, repeat Step 2 until the solution is red-orange. Divide the red-orange solution in the first tube among six tubes so each tube contains the same volume.
Chemical Equation for the Equilibrium System Fe3+(aq) + SCN-(aq) FeSCN2+(aq) + heat Colorless Colorless Red-orange from Fe(NO3)3 from KSCN
Disturbing the Equilibrium
Leave Tube 1 undisturbed; use it as a control. Use a clean, dry spatula to add a small crystal or two of solid iron(III) nitrate, Fe(NO3)3, to Tube 2. Mix. Under Disturbance on your data table, record what you did or added to the system to cause the change you observed. In this and all other observations, pay particular attention to color and color change. Always compare with the control tube or you may miss slight color changes. Phrase your Observed Change so the kind of change you observe is indicated, e.g., "lighter red" or "from grey to pink." Use a clean, dry spatula to add one or two small crystals of solid potassium thiocyanate, KSCN, to Tube 3. Mix. Record observations. Add 5 drops of 0.1 M sodium hydroxide, NaOH, to Tube 4. Mix, observe, and record.
Use a test tube clamp to place Tube 5 in a hot water bath. When the contents of the tube are hot, observe and record. Use a test tube clamp to place Tube 6 in an ice water bath. When the contents of the tube are cold, observe and record. (Data check: Obtain your teacher's initials.) Discard all test tube contents in the waste container provided by your teacher. Do not pour anything in the sink. Rinse the tubes with tap water; remove as much water as possible by shaking before standing the tubes upright in the test tube rack. Follow these same disposal and rinsing procedures after you complete each system below.
System 2: Bromothymol blue
Setting Up the Equilibrium
Half-fill three test tubes with distilled water. Add three drops of bromothymol blue indicator to each tube. Mix thoroughly.
Chemical Equation for the Equilibrium
Bromothymol blue is a weak organic acid with a complex formula. For our purpose, its formula can be abbreviated to HBb. HBb(aq) H+(aq) + Bb-(aq) Yellow Colorless Blue
(Green can be observed if approximately equal amounts of yellow and blue forms are present.)
Disturbing the Equilibrium
To Tube 2 add two drops of 0.1 M hydrochloric acid, HCl, and mix. Observe and record. To Tube 3 add two drops of 0.1 M sodium hydroxide, NaOH, and mix. Observe and record. Explore what happens when you now add NaOH to Tube 2 or HCl to Tube 3. See whether your observations are in agreement with observations you have already recorded.
System 3: Complex Ions of Copper(II) (Cu2+)
Setting Up the Equilibrium Half fill a test tube with 1.5 M copper(II) chloride, CuCl2, solution. Divide so five tubes contain approximately equal volumes. Equilibrium has already been established in the solution.
Chemical Equation for the Equilibrium CuCl42-(aq) + 4 H2O(l) Cu(H2O)42+(aq) + 4 Cl-(aq) + heat Green soln Colorless Light blue soln Colorless
Disturbing the Equilibrium
To Tube 2 add a small quantity (the size of a rice grain) of solid calcium chloride, CaCl2. Mix to dissolve the solid. Repeat the addition and dissolving of solid CaCl2 until no more solid will dissolve. Observe and record. To Tube 3 add enough ethyl alcohol, C2H5OH, to triple the volume of the solution. Mix, observe, and record. Place Tube 4 in a hot-water bath. When the solution is hot, observe and record. Place Tube 5 in an ice-water bath. When the solution is cold, observe and record.
System 4: Dinitrogen tetroxide (N2O4)
Setting Up the Equilibrium
Dinitrogen tetroxide, N2O4, can decompose into nitrogen dioxide, NO2, a reddish brown poisonous gas. So that you may work with these substances safely, your teacher will provide two sealed tubes each containing a mixture of these subtances. Equilibrium between N2O4 and NO2 has already been established in the tubes.
Chemical Equation for the Equilibrium N2O4(g) + heat 2 NO2(g) Colorless Reddish brown
Disturbing the Equilibrium (Caution: N2O4 and NO2 in the sealed glass tubes are poisonous. Handle the tubes carefully to avoid breaking the tubes and releasing the gases.) Place one sealed tube containing the equilibrium system in a hot water bath. When hot, compare to the unheated tube and record. After removing the tube from the hot water bath, cool it under running cold tap water. Then place the tube in an ice-water bath. When cold, compare to the unchilled tube and record.
System 5: Complex Ions of Cobalt(II) (Co2+)
Setting Up the Equilibrium Half-fill a test tube with 1.5 M cobalt(II) chloride, CoCl2. Divide the solution so five tubes contain approximately equal volumes. Equilibrium has already been established in the solution.
Chemical Equation for the Equilibrium heat + Co(H2O)62+(aq) + 4 Cl-(aq) CoCl42-(aq) + 6 H2O(l) Red Colorless Blue Colorless
Disturbing the Equilibrium
To Tube 2 add a small quantity (the size of a rice grain) of solid calcium chloride, CaCl2. Mix to dissolve the solid. Repeat the addition and dissolving of solid CaCl2 until no more solid will dissolve. Observe and record. To Tube 3 add enough acetone, CH3COCH3, to double the volume of the solution. Mix, observe, and record. Place Tube 4 in a hot water bath. When the solution is hot, observe and record. Place Tube 5 in an ice water bath. When the solution is cold, observe and record. Wash hands thoroughly before leaving the laboratory.
Data Analysis and Concept Development
To complete the fourth column on the right side of your data table (headed Direction of Shift), decide whether each disturbance caused the equilibrium system to shift left or right. Record the direction of shift in this column. How do you decide direction of shift? Consider the equilibrium system A B Yellow Green
If a disturbance causes the system to become more yellow, chemists would say that the equilibrium position has shifted to the left because the system must have moved to produce more of the yellow molecules shown on the left side of the chemical equation. If the system shifted to the right you would observe more green in the system. The direction of shift is "right". Use these ideas to decide and record the direction of shift caused by each disturbance. Use your data table to find all cases where a disturbance was caused by heating. After you have found all of these cases, answer the following: How does the direction of shift relate to the side of the chemical equation on which the heat term is written? Write a rule which would allow you to predict how other equilibrium systems would shift when disturbed in this way. Use your data table to find all cases where equilibrium systems were disturbed by cooling.
How does the direction of shift relate to the side of the chemical equation on which the heat term is written? Write a rule which would allow you to predict how other equilibrium systems would shift when disturbed in this way. Use your data table to examine all cases where a disturbance was caused by increasing the concentration of a substance already present in the equilibrium system. Hint: Adding solid Fe(NO3)3 to System 2 increases the concentration of Fe3+(aq) and NO3-(aq) when the solid dissolves. Adding HCl solution to System 3 increases the concentration of both H+(aq) and Cl-(aq) in the system. Write a rule which would explian how the direction of shift relates to the side of the chemical reaction on which the substance with increased concentration is written. In some cases the equilibrium system was disturbed by decreasing the concentration of a substance in the system. Usually this is done by adding another substance not involved in the equilibrium which reacts with a substance in the system, changing it to a different substance. For example, in System 1 you added 0.1 M NaOH (containing aqueous Na+ and OH- ions). OH- reacts with Fe3+ to form the precipitate Fe(OH)3(s). This decreases the concentration of Fe3+(aq) remaining in the solution. Concentration can also be decreased by adding another solvent (acetone or alcohol) to dilute the water in the system. Identify substances whose concentration is decreased in as many cases as you can. For each, explain what causes the concentration of a particular substance to decrease. Write chemical equations where possible.
The equation for the example above is: Fe3+(aq) + 3 OH-(aq) Fe(OH)3(s) For each case involving a decrease in concentration, identify the substance that is decreased in concentration, on which side of the equation this substance is found, and which way the equilibrium is observed to shift. Consider cases where equilibrium was disturbed by decreasing the concentration of a substance in the equilibrium system. How does the direction of shift relate to the side of the chemical equation on which the substance with altered concentration is written? Write a rule which would allow you to predict how other equilibrium systems would shift when disturbed in this way. Write a general rule that would cover all of the types of disturbances you have observed. Write your rule so it can be used to predict the effect of any temperature or concentration disturbance on an equilibrium system
Ronald George Wreyford Norrish was born in Cambridge on November 9th, 1897. His father, a native of Crediton, Devonshire, came to Cambridge as a young Pharmacist to open one of the early shops of Boots, the Chemists, and remained there for the rest of his long life.
After spending his early years at the local Board school, Norrish obtained a scholarship to the Perse Grammar School in 1910. He remembers with deep gratitude his early teachers, in particular Rouse, Turnbull and Hersch, who gave dedicated and individual help to promising young scholars. In 1915 he obtained an entrance scholarship to Emmanuel College, Cambridge in Natural Sciences, but left in 1916 with a commission in the Royal Field Artillery for service in France. He was made prisoner of war in March 1918 and spent the rest of the war in Germany, first at Rastatt and later at Graudenz in Poland. Repatriated in 1919, he returned to Emmanuel College where he has remained ever since, first as a student and after 1925 as a Fellow. Norrish's early research was inspired by Eric Redeal (now Sir Eric Redeal) under whose lively supervision he first took up the study of Photochemistry.
In 1925 he was made Demonstrator and in 1930, Humphrey Owen Jones Lecturer in Physical Chemistry in the Department of Chemistry at Cambridge and upon the death of the first Professor of Physical Chemistry, Dr. T.M. Lowry, was appointed to the Professorship in 1937. He occupied the chair until 1965 when he retired as Emeritus Professor of Physical Chemistry in the University.
Norrish has had the good fortune to work with many gifted students and with them has carried out a wide range of research in the fields of Photochemistry and Reaction Kinetics, including Combustion and Polymerisation. As the study of Chemical Kinetics developed, there was a fortunate integration in the various aspects of the study in which his school of work was engaged, as a result of which the importance of Photochemistry and Spectroscopy to Chemical Kinetics in general emerged. All this was sadly brought to a temporary halt in 1940. During the second world war, while still continuing to direct the Department of Physical Chemistry and to teach, Norrish was concerned with a good deal of research work in connection with various ministries and was able to collaborate with his colleagues on various government committees. It was after the war in 1945 when research was recommenced that work was started with the object of observing short lived transients in chemical reactions. In collaboration with his student, now Professor George Porter, this led to the development of Flash Photolysis and Kinetic Spectroscopy which has had considerable influence on the subsequent development of Photochemistry and Reaction Kinetics, and in the hands of workers in many parts of the world is continuing to develop as a powerful technique for the study of all aspects of chemical reaction.
In 1926 Ronald Norrish married Annie Smith who was Lecturer in the Faculty of Education in the University of Wales in Cardiff. They have two daughters and four grandchildren. Much of their time has been spent in travel.
Norrish has served on the Councils of the Chemical Society, the Faraday Society of which he became President in 19531-1955 and on the Council of the Royal Institute of Chemistry of which he was Vice President from 1957 to 1959. He delivered the Liversidge Lecture to the Chemical Society in 1958, the Faraday Memorial Lecture to the Chemical Society in 1965, and the Bakerian Lecture to the Royal Society in 1966. He was President of the British Association Section B (Chemistry) in 1961, and in the same year was made Liveryman of the Worshipful Company of Gunmakers. In 1958 he received the honorary degree of D. de l'U. at the Sorbonne in Paris and also honorary degrees D. Sc. in Leeds and Sheffield in 1965, Liverpool and Lancaster (1968) and British Columbia (1969). He is an honorary member of the Polish Chemical Society and Membre d'honneur of the Soci?t? de Chimie Physique in Paris. He is a foreign member of the Polish and the Bulgarian Academies of Sciences, a corresponding member of the Academy of Sciences in G?ttingen and of the Royal Society of Sciences in Liege. He is a honorary member of the Royal Society of Sciences in Uppsala and the New York Academy of Sciences. He has received the Meldola medal of the Royal Institute of Chemistry (1926), the Davy medal of the Royal Society (1958), the Lewis medal of the Combustion Institute (1964), the Faraday medal of the Chemical Society (1965) and their Longstaff medal (1969). He was elected Fellow of the Royal Society in 1936 and is still endeavouring to continue to prosecute his scientific activities in Cambridge.
To mark his retirement in 1965, many of his old friends and younger colleagues now occupying distinguished positions in academic and industrial work in Great Britain and abroad collaborated to publish a work entitled "Photochcmistry and Reaction Kinetics". To them and to all others with whom he has worked for over 50 years he is deeply grateful.
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