A -cis to -trans Conversion

Experiment 1: "A cis to trans conversion" Saturday, June 25, 2005 4:19 PM Notes on Theory • Alkenes o They are hydrocarbons with a C=C double bond • Double bonds are stronger and more reactive than single bonds o Hydrocarbons with double bonds used to be known as "olefins", because they had an oily appearance • "Oleum" means oil • "Ficare" means make • Isomers o Geometric isomers have the same molecular formula, but a different geometric arrangement • They also have different physical properties o Alkenes exist as two different geometric isomers • This is because the double bond in an alkene cannot twist very much, so there are two distinct geometric configurations • Cis and trans o The two isomers are generally named cis and trans • "Cis" means the same, which is because the cis isomer will have some important molecular group on the same side of the molecule (think about what C4H8 looks like) • "Trans" means opposite, so the important molecular groups will be on opposite sides of the molecule o Heats of formation, denoted as ΔHof, are the change in heat for some system when a molecule is formed • If they are negative, it means that heat is released…more heat is released when a reaction is more exothermic • Exothermic reactions happen naturally and thus their product is stable • Trans alkenes are more stable than cis alkenes by about 1 kcal/mole, which means that their heats of formation are more negative • Steric hindrance o This is when atoms or groups in a molecule are too close to each other, and so there is repulsion between their electron clouds o Consider cis-2-butene, which has 2 CH3 groups on the same side… • If you measure the van der Waals radii (the radius of an imaginary sphere surrounded the molecule) for one CH3 group, it is 2A…so if the two CH3 groups were as close together as possible but not touching, their centers would be 4A apart • However, they are measured to be about 3A apart…meaning that they are too close to each other! Hence the presence of steric hindrance o In this lab, we are investigating the interaction between the cis and trans isomers of 1,2-ethylenedicarboxylic acid • The cis isomer is also known as "maleic acid" • The trans isomers is also known as "fumaric acid" • Hybridization of carbon o sp3 hybridization

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This involves one 2s and three 2p orbitals (do you know what an orbital is?) forming 4 sp3 orbitals § The bonds between hybridized orbitals are called sigma bonds § The bonds between unhybridized orbitals are called pi bonds sp2 hybridization • Here we have one 2s and two 2p orbitals combining to make 3 sp2 hybrid orbitals, leaving one p orbital unchanged § The p orbitals lie perpendicular to the plane of the hybrid orbitals (remember that the p orbitals bond with a pi bond) § Remember that pi bonds are weaker than sigma bonds sp1 hybridization • Only one 2s and one 2p orbital are hybridized here to form two sp hybrid orbitals • The p orbitals are (once again) parallel to each other, and perpendicular to the line formed by the hybridized orbitals •

Notes on Procedure • Put maleic anhydride into boiling water…this will dissolve it into maleic acid o And it will be a supersaturated solution, which means that since the water is so hot (boiling!), more maleic acid is dissolved in it than is normally possible • Then we let the solution cool, and also put it on ice o This lowers the temperature of the water enough that excess maleic acid (which now cannot be dissolved in the water) crystallizes • Take the crystals out of the beaker by pouring the whole thing through a Buchner funnel (like a coffee filter-type thing)…and use water to rinse the crystals out if they're sticking to the beaker o Note that we shouldn't use too much water here, or else the crystals will just go right back into solution since there is enough water to dissolve in! o However, if water gets through the filter and the Buchner funnel, that's cool…because it'll just add to the dissolved maleic acid which is already in water! • So now we have maleic acid in two forms: dissolved in water, and crystallized • Take the stuff that was dissolved in water (the filtrate) and REFLUX it… o So here is what refluxing is:z • You have a reaction going on in some vessel, but at the top of the vessel we have a condenser • So as the reaction happens and vapors are formed, they have nowhere to go but up this condenser (it is like a pipe thing) • But there is cold water running through an outer wall of the condenser, so that as the vapor travels up, it is condensed and drips back down into the vessel at the bottom • This allows the solution to happen without the loss of solvent! o So we add concentrated HCl and two boiling chips to this dissolved maleic acid which we have, and we reflux for 10 minutes • When the reaction is done, crystallize the product once again by putting it in an ice bath o Also, suction filter it and wash with a little bit of cold water o The product is fumaric acid, by the way :P

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Now we compare properties of our two isomers (maleic and fumaric acid) o Use a MelTemp melting point apparatus to compare the melting points for each product o Also, place each of the crystals in water to compare the solubilities of each isomer

Notes on Results • The melting point of fumaric acid is higher than maleic acid • Fumaric acid is insoluble and maleic acid is highly soluble Experiment 2: "The Structures of Crystals" Sunday, June 26, 2005 3:40 PM Notes on Theory • Solids normally exist as building blocks of atoms, molecules, or ions o These are assembled in an orderly manner into what is called a crystal structure • We describe these crystal structures (since there are different types) by their unit cell o A unit cell is something which, when moved a distance equal to its own dimensions in various directions, can generate the entire structure Notes on Procedure • Simple cubic lattice o Notice that for a unit cell, the spheres are in contact along the cube edge - and so the edge length of the unit cell is twice the radius of one sphere • Body-centered cubic lattice o Remember that the spheres are in contact along the cube diagonals, but not the cube edges! • Face-centered cubic lattice o The spheres are in contact along the face diagonals! • Hexagonal close packing o Here we place as many spheres as possible around a central sphere the most we can get is 6 (that is why it is called hexagonal close packing) o If we take a layer of 6 (plus the one in the middle) of these and put a triangle-shaped layer of 3 underneath it, we have 2 options for the one on top: • If our top layer has its three spheres directly on top of the spheres in the bottom layer, we call it "ABAB type of hexagonal structure" § Zinc and magnesium have this structure! • If we rotate the top 60o so that the balls on top are not directly over the balls on the bottom, it is called an ABCABC structure • Sodium chloride lattice • Structures of some salts o A "tetrahedral" site is made when we have three spheres in a trigonal arrangement, and then on spheres on top of the hole…

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Often we would put a small cation in this hole - this is how the structures of zinc sulphide and silver iodide crystals are arranged We have an "octahedral site" if we can fit 6 spheres around the hole (or the cation which occupies the hole) We can figure out what kind of structure a molecule will have (tetrahedral site or octahedral site) by checking the "radius ratio", which is the ratio of the cation's diameter to the anion's diameter •

Experiment 3: "The Iodide-Persulphate Reaction" Monday, June 27, 2005 12:38 AM Notes on Theory • The point of this experiment is to measure the rate of oxidation of iodide ion by persulphate ion o Reaction #1: (NH4)2S2O8 + 2KI -> I2 + (NH4)2SO4 + K2SO4 o Net ionic equation: S2O82- + 2I- -> I2 + 2SO42• We measure the rate of reaction by using the fact that the amount of iodine produced in a given time gives us the SAME information as the amount of time taken to produce a given amount of iodine o Note that the iodine which is produced reacts with sodium thiosulphate as follows: • Reaction #2: I2 + 2NA2S2O3 -> 2NAI + NA2S4O6 • Net ionic equation: I2 + 2S2O32- -> 2I- + S4O62• In Reaction #1 alone, we end up with a violet color solution, because that is what iodine produces • However, if the iodine reacts immediately with the thiosulphate, it just all becomes clear o We only need to make sure that we have 2 moles of thiosulphate for every mole of iodine • So we can let the reaction run and see how long it takes the solution to turn violet o When this happens, we know that all the thiosulphate is used up…and since we know how much thiosulphate we had originally, we also know how much iodine has been produced! So we know how much iodine was produced in a given amount of time! • Note that we also add starch to this experiment because it forms a blue complex with free iodine, so that the color change is even more accurate • • As we know from CHEM 123, the rate law for this reaction is: Note that also as we learned, you might expect k to be constant throughout! However it is not…because it is strongly dependent on ionic strength (a measure of the effective concentration of ions) o So we try to keep ionic strength constant by adding an electrolyte such as (NH4)2SO4 or KNO3 o When there is a high concentration of ions in solution, the repulsion between 2 negative ions is REDUCED because one or both of them probably attracted a positive ion towards it which can shield the negative charge o We calculate ionic strength using this formula:

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Additional notes… o Some complications can arise throughout this process: • What if S2O82- reacts with S2O32-? • I- may react with I2 and the product I3- could react with the persulfate ion • Reaction 2 has to be much faster than Reaction 1!

Notes on Procedure [This is pending, depending on what Sue says regarding whether we need to know procedure or not!] Experiment 4: "Weak Acid-Strong Base and Weak Base-Strong Acid Titrations" Tuesday, June 28, 2005 12:29 AM Notes on Theory • Bronsted-Lowry: o An acid is a substance which can donate a hydrogen ion (H+) o A base is a substance that can accept a hydrogen ion (H+) • Lewis: o A substance that can accept a pair of electrons o A substance that can donate a pair of electrons • The acid strength of a Bronsted-Lowry acid refers to the ability of the acid to relinquish protons o The stronger the acid is, the more easily it loses protons o And of course, vice versa for bases… • The strength of an acid is inversely related to the strength of its conjugate base o Think about the ionization constant for a strong acid and what the equation is like…you should be able to figure out why! • A weak acid in water ionizes PARTIALLY, and there is an equation for this (try to think of an example) o The equilibrium constant (you know about these!) of this equation is called an "acid ionization constant"…denoted by Ka (as opposed to Kc) • Fun facts on pH… o pH = -log[H3O+] o It is measured with indicator dyes and pH meters • Indicator dyes change color over different pH ranges • pH meters measure the electromotive force of a cell, using the solution as an electrolyte pH meters indicate pH by measuring the electromotive force (e.m.f.) developed by a cell o The meter has a calomel electrode and a glass membrane electrode, and both are immersed into solution o Charge is sent from the calomel electrode through the solution, then picked up by the glass membrane electrode • So the stronger the acid is, the more electrolytes there will be, and the stronger the charge picked up by the glass membrane electrode

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Or in other words, the response of the glass electrode is proportional to the hydrogen ion concentration in the solution EMF is a linear function of pH, so we can figure out what the pH is pH meters require calibration with a buffer solution of known pH usually close to the range under test •

Notes on Procedure • Part A: Titration of a weak acid with a strong base o Create an acetic acid solution of known molarity o Titrate it with an NaOH solution of unknown molarity o Use the indicator phenolphthalein o As the endpoint of the reaction comes (which is when the indicator changes color), add titrant more slowly because the pH rises rapidly here • This occurs at a pH of 5.5 • Part B: Titration of a weak base with a strong acid o Create an ammonia solution of known molarity o Titrate it with an HCl solution of unknown molarity o Use the indicator methyl red • Slow down at pH = 8.5

Experiment 5: Buffer Solutions Tuesday, June 28, 2005 1:15 AM Notes on Theory • A buffer solution is a solution that resists changes in pH by maintaining a relatively steady H+ concentration by shifting the equilibrium between a conjugate acid-base pair • We make buffer solutions by dissolving almost equal concentrations of a weak acid and its conjugate base (a salt of the weak acid) in water • The concentration ratio between [A-] (the base) and [HA] (the acid) is called the buffer ratio o The pH of a buffer solution depends directly on the logarithm of the buffer ratio The greater the concentration of conjugate acid and base in a buffer, the better its ability to resist a change in pH - because the difference in the ratio created by using up some acid or base to neutralize some added substance will be relatively smaller The number of moles of strong acid or strong base needed to change the pH of one liter of buffer by one unit is called the buffer capacity Note that the buffer capacity is affected by the buffer ratio If you consider the Henderson-Hasselbalch equation, note that if the [A-]/[HA] ratio is close to one, changes in either [A-] or [HA] will not change the pH as much…but if the ratio is already far from one, the same amount of change will affect the pH much more

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Notes on Procedure • Part 1: Dilution of a buffer

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o We make a buffer out of acetic acid and sodium acetate o Then dilute it - make a 1:5 concentration and a 1:50 concentration Part 2: Addition of acid to buffered and unbuffered solutions o Take the undiluted buffer and add HCl to it slowly… o Take plain water and add HCl to it slowly…

Notes on Results • For the HCl getting added to the undiluted buffer, the pH stays the same for a bit, and then it drops sharply…then it levels out again o It drops sharply because the buffer capacity has been exceeded, and the buffer is no longer able to resist changes in pH • For the HCl getting added to just water, it drops right away because there is no buffer in place to resist the immediate pH change! Appendix C Sunday, June 26, 2005 5:18 PM Quantitative Techniques Laboratory Techniques Quantitative Transfer of Samples Use of a Burette in Titration Use of a Volumetric (Transfer) Pipette Weighing Replicate Samples of a Solid Filtration Safety Information Tuesday, June 28, 2005 5:20 PM General 1. Pay close attention to instructions. 2. Do not perform unauthorized experiments. 3. Be familiar with first aid measures and the locations of eye wash, shower, etc. 4. Wipe up spills and bottle rings immediately. 5. Report all accidents immediately to the instructor. 6. Wear a protective apron. 7. Use fume hoods. 8. Wear safety goggles. 9. Keep aisles free of obstruction. 10. Keep sinks clean. Fire and Explosions • The vapor of nearly all organic solvents is flammable… o So keep electric sparks, open flames, etc. away from solvents o Prevent vapors from escaping containers by keeping the lids on • Rules: o Know where the closest fire exit, fire extinguisher, etc. are located o Leave immediately in case of fire.

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If your clothes catch on fire, roll on ground or wrap yourself in blanket. Do not run. Limit flammable liquids to amount actually needed for immediate use. Avoid inhaling smoke and gases resulting from fire/explosion. Assemble the apparatus so that it is easy to reach the control valves and switches if a fire occurs.

Fire Extinguishers • Carbon dioxide can be used on all fires except those involving Na, K, Mg, Al • Never use water on electrical fires. • When to use which kind of extinguisher? o Dry chemical is good for flammable liquids and electrical fires o Foam is good for flammable liquids and rubbish o Water-type and soda acid are good for wood, papers, textiles, and rubbish o Active metal fires can be smothered with fine dry soda ash, sodium chloride, and/or sand and graphite • After a fire extinguisher has been used, report it to the laboratory instructor Handling Glassware 1. Carry tubing vertically rather than horizontally. 2. When bending glass tubing, place hot glass on wire gauze or asbestos plate. 3. When inserting glass tube into a stopper Handling Chemicals