Notice!

I’ve found that this book project has been showing up on more and more search engines lately and is also being directly linked to for the information it contains(1).  I therefore find it necessary to warn all persons viewing this document that it is a work in progress, and as such it contains errors of all kinds, be them in experimental procedures that may cause harm, or in faulty reasoning that would get you slapped by nearly any chemistry instructor.  Please for now take the information here with a grain of salt.

Most Importantly!

By reading further you agree not to hold the authors of this document responsible for any injuries/fatalities that may occur from attempting to make any of the products or following any of the procedures that are outlined within.  Chemistry inherently possesses a degree of danger and you must understand this, wear gloves and more if the situation calls for it, your safety is in your own hands, not mine!

Also note that this project is open for contribution by any party on the internet.  Simply submit a section to Rob.Vincent@gmail.com and it will be added into the text pending editing and such within a few weeks.  Any person contributing will have their name mentioned in the credits.  Thank you for reading this, and enjoy!

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Although this document may be directly linked to, it will not work in that manner as I have hotlink protection for PDF documents, however directly linking to the html document is possible, still though I would prefer links be to the main book project page.

 

3.0 Lab Techniques (Basic) 

       This section will cover a number of techniques that can be very useful to most chemists.  Refluxing allows for a more complete reaction of reagents by allowing them to react at a temperature very close to or at their boiling points.  Distillation is infinitely useful for the separation of volatile reagents from their non-volatile counterparts, or for driving a reaction where the product can easily boil off.  Filtering is an essential method of removing solid particulate, this comes in handy when dealing with removing solid catalysts from a reaction mixture, insoluble products, or simply some floating pieces of gunk that come in your reagents bought from hardware stores or your products from nearly any reaction.

       Electrolysis is a difficult concept to master but it opens a whole world of possibilities for the manufacture of a number of reagents.  Powerful reducing and oxidizing agents are both possible to make though electrolysis, as the process of electrolysis is itself both a powerful reducing and oxidizing agent.  Covered in three sections including the rarely covered molten salt electrolysis it is the hope of this text just to give the basic overview to the process, nothing more, as such details would take considerable space. 

      Shifting gears again this section goes to titrations which are quite useful.  Usually used for determining the amount of acid or base present in an unknown they can also be used with some modifications to determine the oxidizing or reducing potential of a solution or even its metal content.  Titrations should be preformed with a somewhat expensive piece of equipment called a buret however careful use of a graduated cylinder can give close enough figures for the home chemist.  This section continues with tips on heating and cooling and maintaining temperatures during reactions followed by two methods of purification, removing water from gasses, solids, and liquids along with the ever useful recrystalization.

       This section rounds out with methods of measuring mass and volume and how useful precise measurements are.  All in all these techniques should find use with most chemists who wish to pursue their hobby at home and should hopefully offer a number of alternatives to the professional ways of carrying out these procedures, enough so to allow utilization of these methods in a home environment.

3.1 Refluxing

** Before you ever heat up a reaction, you should try to assure yourself that this will not cause the reaction to get out of hand.  Always start your reactions cold or at room temperature.  If you believe that heat is necessary, heat the reaction after you mix the reagents. **

 

       Refluxing is a common technique, typically used to accelerate reaction rates.  The most common method of refluxing a reaction mixture in a flask is to submerge the flask in oil, water, or a sand bath that is at a constant temperature slightly higher than the boiling point of the solvent.  Attached to the flask should be a reflux condenser (see section 2.1) to cool the gaseous solvent when it rises out of the flask.  The condenser must be long enough and cool enough (a constant flow of cold tap water is usually sufficient) that the gaseous solvent condenses well below the top of the condenser this is usually apparent, a visual vapor line indicating the height of vapor in the column.  This is especially important if long reaction times are necessary, as it will minimize the amount of solvent loss.  Always be aware of the temperature of the bath and try to keep it as constant as possible.  Monitor your reaction closely to be sure that the gaseous solvent never reaches the top of the condenser!

 

       As the reaction flask heats up some of the liquid(s) will volatize and climb into the reflux condenser.  There they intermingle and fall back into the reaction flask.  Then once again they boil off and climb into the column and mingle then fall back in, it’s somewhat similar to simmering in a culinary process to get the flavors to intermingle.  By refluxing for long lengths of time you can react systems with more then one phase or react insoluble solids effectively with liquids or simply speed up a reaction.  The bubbling helping the mixing of the phases and the heat speeding up the reaction. Refluxing is a very helpful and simple technique that is easy to master and should not be looked upon with apprehension.

 

          Note that if you are trying to accelerate the rate of a reaction you should not go from room temperature to refluxing in one step.  The best method is to heat the reaction in increments of 10 to 20 degrees Celsius and see if the reaction occurs at a lower temperature than reflux.  This is much, much safer than immediately refluxing.

 

 

 

 

3.2 Distillation

        Distillation is one of the most essential procedures in all of at home chemistry in my honest opinion.  When it all comes down to it, distillation is about removing a volatile product from a solution, usually for purification and by the same logic, extraction.  The most simple of distillation apparatuses is to the right.

        It consists of four essential parts.  The tube on the left is the distillation flask (this is also known as the reaction vessel), or in this case the test tube.  This leads to a condenser, in this case a glass tube.  The condenser serves the essential role of taking the gasses produced from heating the distillation flask and cooling them down, causing you to end up with a liquid.  In this case the condenser is the most basic design, air cooled, not very efficient, but made more efficient by a fan blowing on it.  Other better condensers utilize running water to cool them or mixtures that can obtain even cooler temperatures.  

        The condenser tube leads into a receiving flask.  This can be cooled as shown in the picture by ice, this is especially necessary if your product that you are distilling over is significantly volatile.  The final part of the setup, the one that really makes it work is your heat source.  A Corning hot plate is about top of the line for an at home lab but basically anything that gets hot and hopefully has a heating control will work. As usual for heating applications borosilicate is preferred.  For a more precise heat control the use of a 'bath' is advised, this is just a beaker full of a fluid (oil in the picture above) that is heated directly, thereby heating your distillation flask more evenly and consistently and if things get out of control, it can act as a heat sink.  The bath temperature is usually about 5-10°C higher then the reaction flask temperature as a rule of thumb.

       Now, heating a liquid mixture to boiling and condensing the vapor may sound easy, but there are several small but important factors that need to be kept in mind.

 

       First of all, the intent of the distillation is to end up with the most volatile compound of your mixture in the receiving flask and the less volatile compound(s) in the distillation flask. No problem you say, because I am condensing the most volatile component. This is, unfortunately, not entirely true. The other components in your distilling flask also have a vapor pressure, which rises with the temperature. As a result the vapor you are condensing mainly consists of the most volatile compound, but it will also contain a fraction of the less volatile compound(s). This fraction will be large if the boiling points of the compounds don't differ much (less than 40°C) and small when the boiling points are quite different.

 

       For example, distilling wine (usually 12,5 vol% of ethanol in water) will yield a solution which is greatly enriched in ethanol content, but it will still contain a considerable (about 50-40%) amount of water.

 

       Secondly, as your distillation proceeds, the concentration of the most volatile compound decreases and the boiling point of your mixture rises (sic). As a result, your vapor will contain more of the lower boiling compound(s) as the distillation proceeds towards completion.

 

       There are several ways to solve this problem, the most important being fractional and azeotropical distillation.

 

       Fractional distillations is actually quite simple, the idea is to redistill your distillate until it's pure. Now, if we take our previous example of wine, about three distillations would be needed to attain reasonably pure ethanol. It goes without saying that this is both time and energy consuming. Being the inventive chaps they are, some chemists came up with a clever solution to this problem: The fractionating column. This simple but nifty device allows you to separate close boiling compounds in one run or in considerable less runs than a simple distillation would require.

 

       The fractionating column is placed between the heated flask and your still head with the thermometer. Vapors which pass through it cool down as they rise and eventually condense, the compound(s) with the highest boiling points condense first against the walls and whatever else is filling your column. As a result there are countless condensation cycles taking place in your column. The condensation of the highest boiling compound(s) delivers condensation heat which evaporates the most volatile compound which is running down the wall. As a result the vapor coming out of the top will almost exclusively contain the most volatile component, unless you are distilling an azeotrope. One could say that inside the column several successive distillations are taking place simultaneously.

 

       There are several types of fractionating columns and they all have their specific uses. The most commonly used column is the Vigreux column (picture?).  It has a relatively small surface area but a high flow rate. The standard 30cm Vigreux column is ideal for separating compounds that have a difference in boiling point of 20°C or more and it can be used for vacuum distillation. Vigreux columns can be made longer or stacked to improve separation, but above a length of one meter one should consider filled columns. Filled columns come in all different sorts and sizes, but they al work on the principle of maximal surface area and therefore they are usually filled with irregularly shaped objects. They retain a lot of liquid and they are also not very well suited for vacuum use. If you are planning to distill compounds with high boiling points or you are using large or long columns you should consider insulating the column to minimize heat loss.

 

       A very important point when performing a fractional distillation is monitoring the temperature, certainly when separating more than two compounds. You’ll notice that the temperature suddenly skyrockets when the most volatile compound is depleted from your mixture because of the good separation of your column. Good measuring of temperature can only be achieved by correct positioning of the thermometer, which should be just below the bend towards the cooler, so that it’s being immersed in the vapor. (PICTURE!!!!!) Slight deviations of a few mm can cause temperature-reading failures of several degrees centigrade.

 

       Azeotropical distillation could be explained as cheating. The trick here is that a third substance is added to your mixture. This then forms an azeotrope, preferably with the compound you do not need to isolate.  An azeotrope is a mixture of two or more substances, which can’t be concentrated anymore by distillation because both components have the same vapor pressure at the azeotropic point.

 

       In this case, the azeotrope needs to have a boiling point that differs substantially from the other component in your mixture. Then you can remove the otherwise hard to separate component azeotropically with the third substance. This process is often used during esterification. Toluene is added to the mixture, which forms an azeotrope with water and boils off. After cooling down the water and toluene separate back into two layers, but this is not always the case.

 

       Azeotropes can also complicate a good separation attempt. A good example is nitric acid. Nitric acid forms an azeotrope with water, which contains 69.2% nitric acid. To concentrate further simple distilling won’t work and neither does fractional distillation. So what now? The most commonly used method is to break the azeotrope by adding another substance. This is completely the opposite of azeotropical distillation where a substance is added to form an azeotrope, so beware of confusion.

 

       In the case of nitric acid, the most applied method is to add sulfuric acid, as this has such a great affinity towards water it will easily “steal” the water from nitric acid. As a result the nitric acid behaves as if it were all by itself in your flask and thus obtains its’ normal boiling point, which is far lower than the hydrated sulfuric acid. One can also snoop the water off by adding a very hygroscopic salt like magnesium nitrate, the only condition being that it does not react with the acid.

 

       The regulation distillation setups used in chemistry labs use ground glassware.  The joints are all tapered glass and fit together snugly, with or without the use of a sealant, which can be anything from silicone gel to concentrated sulfuric acid.  There are two extreme sizes for the professional distillation setup, the 14/20 sets which would be considered the smaller scale, and the 24/40 sets, considerably larger scale.  For example, the largest flask I have seen for a 14/20 set is 250 ml, the largest for a 24/40 is 10 L.  The smaller setup has the advantage of taking up less space and using less of the distillate to wet the vessel therefore resulting in a greater yield, it is great for distilling small amounts but the purification of 3.8 L of over the counter paint thinner may be a pain.  In contrast a 24/40 setup is perfect for this larger project.  But is not good for small amounts, when very small amounts are used the apparatus first needs to be ‘wetted’ with vapor and as such a majority of your liquid might end up lost on the walls of the apparatus and also the larger setup of course takes up a larger space.  It is really up to the chemist and what scale they will be working on to decide.  However the 24/40 joints are more readily available online and elsewhere then 14/20.  But all of this varies country by country, in Belgium for example the most common joint size is 29/32.  Finally in addition to ground glass a time tested method is to use rubber stoppers and glass tubes to connect parts of a distillation apparatus, the drawback being that rubber is attacked by a number of reagents such as oxidizing agents and organic solvents.

 

Picture 1 A 14/20 distilling apparatus the reaction flask being heated in an oil bath.

 

The Metal Distillation Apparatus:

Although the materials section goes to great lengths to show that glass is a great all around material in which to perform reactions there is no reason a metal distillation apparatus should not be considered for some undertakings.  The most notable advantage of a metal apparatus being the high working temperature and the availability of the materials with which to construct one with.  Shown above is a copper condenser in production.  Such an apparatus could be used to distil any of a number of solvents for purification purposes.  It could also be used for inorganic applications such as distilling sodium, something that would be impossible in normal glassware due to the high working conditions. 

Such an apparatus can be simply put together with exoxy and then connected to glassware through rubber stoppers and tubing.  Most chemists will limit themselves to the normal glassware distillation apparatus, but metals are more cost effective and they are more durable.  Definitely a good investment.  Copper is a great choice with which to work, the metal is bendable with the proper tools and it is fairly inert, not to mention most of the connections are readily available as is shown buy the photo above of the distillation apparatus prior to assembly.  However distillation apparatus can be made from other metals as well.

This crude apparatus is a condenser fabricated from steel threaded pipes.  The connection on the right leads from the distillation vessel and connects to the side of it.  The left side actually connects to a jar with the lid attached there to distill into it.  And the copper piping around the vessel cools it, as well as additional piping added to simply hold it in place.  Crude but effective there are many other ways to prepare distillation apparatuses from metals and over the counter items if one is willing to take the time to make one.

 

Here is a quick checklist to follow before performing a distillation:

  1. Check if your compounds form an azeotrope. If not, or if your starting material is below or above the azeotropic concentration, proceed with a normal distillation first. If your compound does form an azeotrope, proceed to step five.
  2. Check the difference in boiling points.
  3. Check the vapor pressure of the higher boiling component at the boiling point of the lower boiling compound. If this is significant and purity is required, use a column.
  4. If you can’t achieve sufficient separation with a column or you don’t have a sufficient column at your disposal, check weather you can use an azeotrope to remove ONE of the components.
  5. If your compounds form an azeotrope, which you want to separate, check if there’s something that’ll form a stronger azeotrope with it or which has a greater affinity for the compound.
  6. Last but not least, check if your components won’t decompose at the necessary temperatures. If yes, carefully read section 8.6 under the advanced techniques on how to work with and distill under vacuum .

       

 

        Distilling HNO3 (Nitric Acid) from a mixture of a nitrate salt and sulfuric acid is a time tested way to isolate this useful oxidizing acid.  So an adventurous chemist combined an unweighed quantity of NaNO3 (sodium nitrate) into a flask with a large quantity of 94% H2SO4 (sulfuric acid) and attached a condenser though which water was run and in turn this ran to a receiving flask.  Some boiling stones were also added (pieces of obsidian) to help ease the boiling process and make it smoother.  However complications where run into shortly after heat was applied, the black boiling stones started to color the mixture black, and the gas running into the condenser was a dark red, not what would be expected for a clear to off yellow acid.  Regardless distillation was continued and in the end the chemist ended up with 20 ml of a dark red volatile distillate.  Upon addition to water it decolorized and left an acid solution.  Although not obviously apparent to the chemist at the time they had distilled NO2, a highly toxic gas that is one of the decomposition products of nitric acid.  Their ambitiousness and inexperience resulted in them heating the reaction mixture too high and the concentration of their acid only would have allowed 95%+ HNO3 to make it over, which would have called for a vacuum because high concentration HNO3 decomposes somewhat more readily then the more common 70% grade.  The chemist eventually realized their mistake by simply observing the physical properties of the product ran in contrast to those of the desired product they thought they had.  Luckily the chemist made it out unscathed.

 

Important Safety Concepts:

Heating:

       When distilling one should use a safe method of heating, which prevents your glass from cracking and improves its life. It also protects you from a bursting apparatus, which showers you in dangerous chemicals. If heating with a flame try to keep away from that section of your apparatus.

 

Flames:

       Open flames are NOT a safe way of heating. Flames cause localized overheating and are especially a hazard when distilling flammable substances.  Flames that are non-diffuse such as flames from a torch can cause extreme stress and failure of glassware, either use flames to heat a bath that your reaction vessel is in, or use a very diffuse flame.

 

Electric Heating:

       Electric heating mantles or hotplates can be used with certain exceptions. Under normal pressure one can safely use a hotplate (preferably with magnetic stirring) to heat an Erlenmeyer. However, one should be careful with flammable substances. If the temperature of the hotplate surface is higher than the auto ignition point of your substance, it cannot be used unless the whole apparatus is sealed to the entry of atmospheric oxygen and the output gasses are property taken care of. Same goes for electric heating mantles. Heating mantles should be used for flasks because a hotplate causes localized overheating with its flat surface! Note that heating mantles usually impair magnetic stirring, unless you buy a magnetic stirrer which fits the heating mantle, but these are usually within the >500$ range.

 

Baths:

       Baths are ideal for heating flammable substances, certainly water baths. Water has the advantage of being not flammable and it’s high heat capacity can be beneficial when distilling. However, water evaporates rapidly when used above 80C. Oil baths don’t evaporate as quick as water, but they have other disadvantages. Most commonly available oils smell and are flammable. They also pose a severe explosion hazard when distilling oxidizing material. Ideal are silicone baths or other non flammable synthetic oils, but these are usually hard to get and/or expensive.  For lower temperature applications or if you just don’t care about your hot plate a sand bath can be used by simply filling a pan with sand and putting on your plate, but at higher temperatures the insulating effect of the sand can burn out the heating element in a hot plate so be warned.

 

SealingJoints:
       Sealing your joints is very important. Improperly sealed joints will cause losses but they also pose a safety hazard, because they allow air to enter. This is mainly dangerous when distilling flammable substances or when distilling under vacuum. Joints are commonly sealed with grease. There are several types of grease and they all have their disadvantages. Vaseline is cheap to get and easy to use, but it’s chemical resistance is limited and it will contaminate organic solvents. Other options include silicon oil, but this is expensive and does not provide 100% chemical resistance either, and specialty high vacuum greases composed of higher fluorinated hydrocarbons but they can cost a bundle. Therefore the author recommends cheap, white, teflon tape available from any hardware store to seal pipe joints. This easily fits between the joints and has excellent chemical resistance. Wrap one layer around the ground glass male side and press into the opening opposite, giving a slight twist, do not force or twist too much though or you might crack or otherwise damage your glassware.

 

        This is a common procedure for the inorganic chemist, less so for the organic chemist but none the less important.  Evaporating to dryness is a feasible method of recovery of a pure product providing:

1.      Your intended product will not decompose at the temperatures necessary to volatize the solvent. [Note: Vacuum can decrease the necessary temperature to remove the solvent and prevent decomposition of your product]

2.      Any other compounds in your solution besides your intended product are also volatile.

3.      Your intended product will not volatize to any major degree at the temperatures used.

4.      Your intended product will not explode at the temperatures used/is not pyrophoric at these temperatures.

5.      An extremely pure product is not required/additional purification will take place.

        Examples of simple procedures would be:

·         Making AgNO3 by dissolving silver in HNO3 and boiling off the HNO3/Water.

·         Neutralizing BaCO3 with HCl then boiling off the water and HCl.

Examples of procedures that will not work are:

·         Boiling off the water from commercial bleach (NaOCl decomposes, NaCl/NaOH impurities)

·         Dissolving Na in MeOH then boiling off the alcohol (NaOMe decomposes)

·         Dissolving Al in HCl then boiling off the HCl solution (AlCl3*xH2O decomposes to oxychlorides)

        Things to watch out for consist of azeotrope distillation when involving liquids and carry over of less volatile insolubles.  Additionally when a solution has nearly evaporated there may be a heavy precipitate on the bottom, this can cause 'bumping' in the flask which can bounce a flask off a hot plate or even crack it due to the pressure of the vapors rising though the precipitate.  To avoid the worst of this you can cool the solution when a precipitate starts to form then filter it then resume heating, or resort to magnetic stirring, or heat at a lower temperature then the boiling point of the liquid.  When creating a salt by reacting an acid with a metal/carbonate/hydroxide etc. be sure to use the stoichiometric quantity of acid if at all possible, excess acid will only have to boiled off resulting in noxious clouds that kill grass, other plants, your eyes, and lungs (this is assuming the acid is volatile, e.g. H2SO4 will not be easily volatile).

·        Sublimation

Sublimation is usually considered the process by which a solid can go to the gas phase without passing through the normal intermediate liquid stage.  A number of substances are known to sublime, common substances include naphthalene, paradichlorobenzene and the most famous of all, iodine.  However many substances can and do sublime even if a liquid phase does exist.  The main thing to consider is a vapor pressure.  All substances theoretically have a vapor pressure, even high melting solids.  Normally a substance has a higher vapor pressure the closer it is to the melting point and from there the closer it is to the boiling point.  But many substances, especially organics have vapor pressures that are appreciable at room temperature.  If the vapor pressure of a substance can be increased, by applying vacuum or by heating, or both then there is a better chance of subliming the product.

Sublimation can be considered a form of short path distillation and can be done providing two criteria are met.  1)  The substance to be sublimed must have a high vapor pressure.  2)  The substance from which the product is to be sublimed must have a relatively low vapor pressure.  Sublimation at room temperature is usually slow, but as stated above, applying vacuum and heating increase the vapor pressure of substances and allow more of the molecules to escape in the gas phase and therefore allow more of them to recondense on the conveniently placed cooler areas of a vessel.  The picture shown here is just one way to setup a sublimation apparatus using an Erlenmeyer flask with a side arm to which vacuum is applied.  The crystals cooling on the test tube inserted into the rubber stopper and full of ice.  When the sublimation is complete the stopper is removed and the crystals are scraped off.  Other simpler apparatuses, can easily be improvised just knowing the normal parts of sublimation, basically you need a partially enclosed environment, a cooler part of a vessel, and a warmer part (a simple example being a jar with a bowl sitting in the opening full of ice, with mild heat applied to the bottom).  But even at room temperature sublimation will occur as stated, in which case crystals will usually sublime to the top of a test tube. 

       Sublimation is great for substances with high freezing points which could otherwise allow the substance to solidify in distillation apparatus and possibly clog it during operation.  It is also great for substances that might decompose at higher temperature or have prohibitively high boiling points (some substances with high boiling points can sublime at considerably lower temperatures).  Usually not a procedure for the large scale, sublimation does offer yet another tool with which a chemist can retrieve a product from a reaction mixture or a reagent from an over the counter product. 

3.3 Filtering 

       So, you have a mixture of a liquid and a solid that you want to separate.  Filtration is the answer.  It’s easy, fast, and effective.  The only things you really need are a funnel and a piece of filter paper (a coffee filter will work too, or even a wad of cotton). To determine which type of filtration is best, you need to know whether you want to keep the solid or the liquid.  There are three common types of funnels that are useful to a chemist: liquids, powder, Buckner.  A liquid funnel has a long, narrow spout and is usually best for simple filtrations, using a cotton wad or a properly folded filter paper or coffee filter.  However because the spout and subsequent area that drains into the spout is so small it is only good for filtering a tiny amount of precipitate otherwise it may become too packed to filter, even with vacuum filtration.  Powder funnels have a much wider spout and are useful for filtering things that might clog a smaller spout funnel.  Buckner funnels are the best for isolating solids, but they work best with a vacuum (more on that below).

       You may be wondering, what good is filtration anyways?  I can always decant my mixture can’t I?  Well, yes you can, but there are some inherent problems associated with decanting.  1)  If the solid material left at the bottom of your flask was your goal, what you have left will contain appreciable reaction mixture, and some of your solid will probably be lost during the actual decanting process.  Or 2) If the liquid is your goal appreciable liquid will be left in your solid that could otherwise be obtained by filtration, and also small particulate may come over during the actual decanting process ruining the purity of your new reagent or the solid may not totally settle out.  Filtration is almost always preferable to decanting.  Also before selecting a filtering method one should consider a few things:

1.      What quantity of liquid will I have to filter?

       Number one is important because gravity filtration is not a good method to filter large quantities of liquid, both due to the time involved and the possibility of clogging of filter paper, in this case vacuum filtration is more appealing.

2.      Am I trying to recover my liquid component?

       If you are trying to recover your liquid component and not your solid component, sand or diatomaceous earth could be put onto the filter paper to increase the efficiency of the filtration and speed it up.

3.      Am I trying to recover my solid component?

       If the solid component is your goal then special care should be taken to make sure your solid is recoverable in high quantity from the filtration method of your choice.

4.      Do I think the crystals, both in volume and by their size might clog my filter paper?

       The possibility of crystal masses clogging filter paper is exceptionally high in the case of really fine precipitates, which can cause a complete stop to filtration, in addition large quantities of crystals can do the same or overflow my filtration method. 

5.      Do I have fine particulates that might pass though filter paper?

       If there are ultra fine particulates in the filtrate there are very fine filters that could be used to remove them.  Letting the solution rest for a few days can ‘age’ the precipitate resulting in a more manageable solution or filtering though sand or diatomaceous earth or even a fine sintered glass filter can remove many fine particles.

6.      Do I have to filter it while hot?

       If you have to filter a solution when hot the first addition of the solution to the filter paper can result in crystallization on the filter paper resulting in almost no filtration ability, which can be catastrophic in some situations where a very hot solution is teetering around.  Pre-washing the filter paper with hot solvent or heating the glass/porcelain parts of the apparatus in an oven can help here.

7.      Does my substance to be filtered contain components that may not take nicely to filter paper?

       Finally, my last point to consider, how will the actual substance passing though the filter paper affect it?  Most of the time it won’t affect it greatly, but some copper complexes can attack cellulose, as can strong sodium hydroxide solutions or strong oxidizing solutions like concentrated nitric acid, passing these though a normal filter paper could spell doom for the reaction and will at the least call for you to filter again, for oxidizing agents try filtering with glass wool(1).

       So how do I filter my substance?  Well that depends on what type of filtration you want to take advantage of (see below).  But the normal procedure is to slowly decant your liquid into the funnel without adding the solid at the bottom.  In this way the filtration can proceed smoothly, because the solid particles will slow it greatly.  The liquid mix is added in small increments to your filtration method, if you are taking advantage of the insolubility of a salt yields may be increased by cooling the mixture until it is near the freezing point to depress the solubility.  After each addition you should wait until the addition is nearly though the filter paper.  Then add more, keeping up the process till all of your solution is used.  Once only a small amount of liquid is left in your beaker swirl around the liquid and dump it in all at once to get your solid onto the filter paper. 

       When all of your liquid has been added the, solid left in the filter paper is washed with cold solvent (whatever your solution is) is added to the flask you were filtering from and swirled around to remove any solid left and this is added to the filter cake to wash it and add this extra little bit of solid to the batch.  The cake can then be washed with additional aliquots of solvent.  After your solid is thoroughly washed it can be removed from the filtration apparatus and spread out on filter paper or in watch glasses to dry.  The basic components of a filtration system are a membrane to separate the filtrate from the non-filtered solution, and a place for each to go.  There is room for emergency improvisation here, the author of this text has for instance seen a beaker with a coffee filter over it held in pace with a rubber band inverted over a collection vessel and heated externally with a torch, the increasing vapor pressure in the flask forcing the liquid though the filter.  Although this is not a recommended method it just goes to show filtration is a mechanical process and can easily be modified to be accessible in your situation.

(1)

In the case of filtering oxidizing solutions you can give glass wool a try, or alternatively fiberglass insulation, this should first be cleaned by stirring with hydrochloric acid and drying, if your product however looks bad you may want to find a different source for your glass wool substitute.

3.3a Selection of filter paper 

       First and foremost: size.  The filter paper you use must fit your funnel or bad things will happen.  The filter paper should fit entirely within the sides of the funnel.  For a Buchner funnel, the filter paper should cover all of the holes on the flat portion, but the edges of the filter paper should not touch the vertical sides of the funnel (it won’t seal under vacuum otherwise).  If you don’t want to buy all kinds of different sizes of filter paper, remember that a big piece can become a small piece, but not vice versa.

The type of filter paper you use really depends on the size of the solid material you are filtering away.  For most applications coffee filters are sufficient, but coffee filters are relatively thin and can’t take much abuse.  If you can get real filter paper, do it (The stuff from Whatman is the best, and not all that expensive).  For larger crystals a wad of cotton packed into a liquid funnel works very well.  If you have a Buchner funnel, you need a flat piece of filter paper so coffee filters are no good unless you cut them up to fit your Buchner.  For filtration of very fine particles, like charcoal dust, you can find very high efficiency filter paper, but the cost is rather high.  If you really need it, the best stuff is made from Teflon (a slightly cheaper alternative is a nylon filter paper but this is still quite expensive and a special commodity) and can be found through chemical suppliers or chromatography suppliers.  Using Teflon filter paper requires vacuum filtration.  You can sometimes sidestep the need for Teflon filters by using a “filter aid” such as diatomaceous earth, sold by the trade name “Celite.”  Celite is essentially very, very fine dirt, but it really works well!  (See the vacuum filtration section for details on its use).  As a final note, all filter paper, no matter how expensive, is really only good for one use.  It’s not worth ruining an experiment to save a few cents by reusing a piece of filter paper.

3.3b Gravity Filtration 

       Gravity filtration is slower than vacuum filtration, but requires less equipment and is generally more effective.  Gravity filtration is better for isolating the liquid phase than the solid phase, but if you don’t have a vacuum, it will still work reasonably well for isolating the solids.  To perform a gravity filtration, you need a funnel (liquid funnel is best, but a powder funnel is okay too so long as you use filter paper and not cotton), filter paper, and a flask to collect the liquid.  There are many ways to fold a flat piece of filter paper in to a cone shape.  The best way is called “fluting” and consists of folds in opposite directions along the diameter of the filter paper (FIGURE).  A simpler, although somewhat less effective method, is to fold the circle in half and then fold the semi-circle in half again.  This wedge can be opened to form a cone, such that half of the cone has a one-ply layer of paper and the other half has a 3-ply layer.  Fluting the filter paper will give a significantly faster flow rate!  Fold the filter paper whichever way you choose and then place it inside the funnel.  Clamp your receiving flask in place (important!) and then set the funnel on top.  Slowly pour the mixture into the filter paper cone.  The liquid level should never be higher than the top edge of the filter paper or the mixture will spill over and go through unfiltered. 

       Gravity filtration is relatively slow (remember, fluted filter paper has a faster flow rate, it’s worth the extra folds!) so just add enough of your mixture to get close to the top and then let all of the liquid flow through before adding more.  This may seem unnecessarily slow, but if you slip and add too much of the mixture and some gets through unfiltered, you’re just going to have to start over again.  Once you’ve poured in all of your mixture, rinse your reaction vessel with the solvent and then rinse the solids thoroughly.  Let the solvent drain completely before adding more to rinse.  If the solids are slightly soluble in the solvent, be sure to use ice cold solvent and try to use as little as possible.  Gravity filtration cannot be used reliably for solids that precipitate as ‘gels’ such as iron or aluminum hydroxide, these are very difficult to remove water from and easily clog filter papers if not assisted by vacuum.  Also note that filtering organic liquids of low density or just in general through filter paper can prove difficult without vacuum as their low density gives little incentive to pull them through the filter paper and in addition filter paper contains water as papers are hydrophilic and as such they can ‘intimidate’ organic liquids from readily passing through them.

3.3c Vacuum Filtration 

Filtration using a hand vacuum pump and a Buckner Funnel in a Erlenmeyer Flask with Side arm.

 

       Vacuum filtration is the fastest way to filter a mixture, but it’s not always the most effective.  If you want to isolate the solid material, this is the best way to go.  You will need a Buckner funnel, a piece of filter paper, a rubber sleeve for the funnel, an Erlenmeyer flask with a vacuum sidearm, and a vacuum pump.  First, put the rubber sleeve around the spout of the Buchner funnel.  The sleeve should fit snuggly and be larger than the opening at the top of the Erlenmeyer (we want it air-tight, remember).  Clamp your Erlenmeyer in place (important!) and then set the Funnel on top of the Erlenmeyer and drop in the filter paper.  Wet the filter paper with a little bit of the solvent from your reaction (this will help the vacuum seal).  Connect the hose from your vacuum to the tubing adapter on the Erlenmeyer.  Turn on the vacuum and then slowly pour the mixture onto the filter paper.  The solvent should be sucked through quickly.  Try to keep the filter paper moist since it won’t seal when it’s dry and the solid can be sucked through around the edges of the paper.  Once all your mixture has been poured onto the filter rinse your reaction vessel with solvent a couple of times and then rinse the solid a few times with a little solvent (if the solid is slightly soluble in the solvent, be sure to use ice cold solvent and try to use as little as possible).  Pull air through the solid for a minute to dry the solid.  Then turn off the vacuum and you’re all set.

 

For filtration with Celite, set up the filtration apparatus as described above, then add an even layer of Celite 2-3 cm thick over the top of the filter paper.  Turn on the vacuum for a few seconds; the layer of Celite should compress a little.  Wet the Celite with solvent and then turn the vacuum on again briefly.  Repeat the process of adding a little solvent to keep the Celite wet and then turning on the vacuum to pack it down.  This process is essential to separating very fine dust from your mixture.  When you are satisfied with the packing of the Celite, turn on the vacuum and slowly pour your mixture onto the filter.  Try not to disturb the Celite too much, you want it to stay as even as possible.  Next rinse the reaction vessel and the Celite.  I repeat, rinse the Celite.  A lot.  There is a lot more to rinse than in a regular filtration.  My favorite method for this is to turn off the vacuum, add a volume of solvent that is roughly the same as the volume of Celite (the solvent shouldn’t go through the Celite filter with the vacuum off) and then suck it through by turning on the vacuum.  I usually repeat this process 2 or 3 times.  Finally, if you are using Celite as a filter aid, it is all but impossible to recover your solid material, so be aware!

3.4 Electrolysis 

              An amateur chemist can easily define electrolysis as any reaction that calls for the direct application of an electric current to a chemical, either on its own, or in solution, for purposes other then heating.  Our point here is to understand how we can transform chemicals using electricity. The key word is “ions”. What is an ion? A quick and dirty definition would be: “It’s an electrically charged atom or molecule”. The positively charged are called “cations” and the negatively charged are called “anions”. Ions have different chemical and physical properties than the original atom or molecule. All anions have their own name, so Br- is baptized Bromide and NO3- is Nitrate. Cations with more than one possible charge also have names, so Cu+ has one electrons missing and it’s baptized Cuprous, Cu++ has lost two of them and becomes Cupric. [See table in section 1.3 for a list of common cations and anions]

        Many chemicals, including all salts, are made of opposite charged ions “glued” together by electrical forces. And in these ‘ionic compounds’ electrolysis can help unglue them.  There are basically three schools of electrolysis that you should familiarize yourself with.

Pure Compound / Molten Salt Electrolysis

This is the simplest form of electrolysis.  Not in its practice, but in concept.  A pure compound. ionic in nature, i.e., consisting of a cation and anion, is heated until it becomes liquid.  Ionic liquids are good conductors of electricity and therefore are able to be electrolyzed directly.  Once molten, electricity is applied and the compound breaks down into its constituent parts.  For example, molten sodium chloride when subjected to DC electrolysis will break down into liquid sodium metal and chlorine gas.

Aqueous Electrolysis

The most common form of electrolysis for the at home chemist.  An ionic salt is dissolved in water and a current applied.  Depending on the reduction and oxidation potentials of your cations and anions you get different products.  For example, electrolysis of water, with a little salt added to aid in conductivity will yield hydrogen and oxygen gas, excess salt and a higher current will yield hydrogen and chlorine gas.

Non-Aqueous Electrolysis

A compound is dissolved in any liquid other then water and current applied.  Different products are possible under different conditions.  Products are possible with non-aqueous electrolysis that are impossible in aqueous electrolysis or would require high temperatures for molten salt electrolysis.  For example, it is possible to obtain lithium metal as a deposit in the electrolysis of lithium chloride in pyridine, whereas lithium metal would react instantly in water and molten salt electrolysis would require temperatures of several hundred degrees Celsius.

       There are two different types of electricity available.  There is alternating current (AC) and direct current (DC).  Alternating current is the type of electricity that comes out of the wall, this is not good for electrolysis, alternating current changes which side of your cell is your cathode and which is your anode about 60 times a second, this means that almost nothing can be accomplished with it(1).  If you have a cell full of water and something to make it conductive and put two electrodes into it, plugging it into the wall the only thing you will do is heat your solution to boiling with resistance heating, make a random explosive gas mixture above your water, and deform your electrodes.  Therefore your only real choice for productive electrolysis is DC.

        So where does DC come from? I will not detail the electronics here. There is plenty of information elsewhere on the internet.  But you usually have two sources, converting the output of your wall adaptor to DC, or using batteries.

 

·        Batteries: Not very useful, except for simple demonstrations. Those little square 9V batteries make me laugh! If you insist on using batteries, be a man, use 4 “C” size (big) batteries or a lantern battery or even a car battery, but if you go with the car battery then you’ll need a charger, and you could just use that directly anyways.

·        Adapters: These are very common these days and most households have one or two spare ones from old equipments that broke. 6 to 9 volts is fine for most experiments and they usually deliver above 0.5A. Another common source of power that falls into this category packing a little more punch is the car battery charger.  A good one can supply from 0 – 12 V and from 0 – 55 A, and can be procured cheaply from second-hand stores.  Old ATX power supplies from computers are marvellous for electrolysis because they yield high amperages (up to 20A) at low voltages (5 and 12V).

·        Build your own simple power supply: If you have some skills on electronics, you can built your own power supply just using a transformer and a single 1N4001 diode.

 

Build your own variable power supply: An extra potentiometer (100Kohms) and a power transistor like  TIP 31 (3A, 40W),  TIP 41 (6A, 60W)or TIP3055 (15A, 90W; a.k.a. 2N3055), can give you control over the voltage supplied by the batteries, the adapters or your home built power supply. Don’t expect precision or stability though:

 

A power source capable of delivering at least 0.5 ampere would be nice. The product yield per hour depends on the current, so your current should be proportional to your hurry and your electrode surface area (too much current per square centimeter may cause unwanted results).

The electrolysis process is sensitive to the voltage applied, but if you want a “rule of thumb” number, 9V will perform most tricks. 

An exercise in calculating yields involving electrolysis.

Let’s say that you want to make bromine, and just to simplify things let’s say you have some lead (II) bromide laying around to perform molten salt electrolysis on.

First off you divide your reactions into half reactions, the half reactions, when added together cancel out but separately they give the number of electrons necessary and help in visualization.

Pb2+ + 2e-  Þ  Pb(l)

2Br- Þ Br2(g) + 2e-

Notice how the number of electrons on one side of the equation match the number of electrons on the other side of the other equation.  That is because while one thing is reduced (gains electrons) another thing must be oxidized (loses electrons) a good way to remember this is the mnemonic OILRIG [Oxidation Involves Loss, Oxidation Involves Gain (of electrons)].  Couple this with the mnemonic RedCat [Reduction occurs at the cathode] and you can figure out where your products will be produced at.  When wires are color coated in their normal manner the cathode is the black wire, the negative (-) wire.  By default then the red wire is the positive (+) wire, is the anode.

Now that you know your reaction, what kind of power source are you using?  Maybe you’re using a simple wall adaptor, possibly from a phone charger and the phone broke.  Looking on the plastic you might even find that it is 6.2 V and .5 A.  That’s all the information you need off that.  The next step is to calculate the number of coulombs ( C ) of electricity that pass though the molten mass.  A coulomb is a unit of measurement specific to things involving electricity, it is equal to one amp times one second.  So it actually measure the quantity of charge moving though the cell.  So let’s say that we are going to run this for 30 minutes:

.5A . .5h . (3600s/1h) . (1C/1A.s) = 900 C

To explain the above equation you can see the that .5A came from the power of the power supply, the .5 hours came from the time the cell was running, the other numbers are conversions to the number of seconds in an our and the number of amp-seconds in a coulomb.  This is the generic representation and you can just plug in your numbers in the above equation to get your own unique answers.  Now we see from the two equations way overhead that for every one mol of Br2 generated  two mols of electrons are simultaneously brought into existence.  Now we need a new unit of measurement for electrolysis, a Faraday (F).  A Faraday is the number of electrons necessary to reduce one mol of a single charge unit.  A Faraday is equal to 96,500 C.  From here we use a new equation:

g. Br2 = 900 C . (1 F / 96,500 C) . (1 Mol Br2 / 2F) . (159.8 g Br2 / 1 Mol Br2) = 1.5 g Br2

In the above equation we took the number of coulombs that we got from the first equation and multiplied it by the conversion for Faradays to coulombs and that by the mols of Br2 and the number of Faradays involved, e.g., in this case two mols of electrons are involved and therefore there are 2 F.  This is all multiplied by the grams of Br2 per mol to give 1.5 g of Br2 produced overall.  To get lead you could convert the grams of Br2 produced and make it into mols, then you could multiply that by the grams per mol of lead or just substitute that information into the last part of the above equation.

Well, most of you are thinking, “Only 1.5 g…. what the heck, I wanted a liter!” Well, this is not the setup you would use for massive Br2 production.  But electrolysis is good for producing small quantities of hard to obtain chemicals, and you could increase your yields by either increasing the amps of your power supply or by running the setup longer.

 

 

       Aside from a power supply the second most important consideration are the electrodes.  Many of the metals are similarly conductive for all intents and purposes and therefore it is better to consider them in terms of their chemical resistance relating directly to whatever environment you are planning to perform electrolysis in and the ease of procurement of the electrode material.

 

·        Copper can be easily obtained from common wires. Copper has a wide resistance to many environments.

·        Zinc can be found as the outer metallic shell of common batteries (the cheap ones, called carbon-zinc). Not good for acidic environments or basic environments.

·        Carbon or graphite: These are very useful electrodes, since they do not oxidize as anodes. Well, they don’t last forever; they are attacked by oxygen, originating CO2, or maybe just disaggregate in the solution, but are much, much, cheaper than platinum electrodes, so they are widely used. The most common are pencil’s graphite. These have very different compositions, and may or may not last long. In fact, they may even be very bad conductors. Another source of carbon electrodes is the carbon rod that every carbon-zinc battery has inside. They are better than pencil’s graphite. The best carbon electrode I found is a rod of graphite covered with a copper layer found in solder’s supply shop. When you strip the copper with ferric chloride (or electrolysis), a resistant graphite electrode is left. Another option is graphite from electric motors sliding contacts. They are small, but easy to find and quite resistant. I have a couple from large polisher that are 23x16x6 mm.

·        Lead is used often as an inert electrode. You will probably find it in a sporting goods shop: fishing weights, gunshots in general, airgun bullets etc(3). It melts easy and you can cast your electrodes melting it with a blowtorch and a scoop.

·        Nickel: Coins from some countries contain an appreciable amounts of nickel(2) and can be used as electrodes, additionally nickel can be procured from scrap yards or specifically for electrolysis.  Nickel is an excellent material for electrolysis of highly basic solutions.

·        Iron:  Iron is attacked readily under acidic conditions and somewhat slower under basic, but it is commonly available and may find some use in a pinch.

·        Platinum: If you have the money, it’s the most resistant anode I know. Fine wire is ok, but for larger surfaces, other metals plated with platinum can be used.

·        Silver/Gold:  A cheaper alternative to platinum somewhat more reactive.  Available from coin suppliers in the form of collectable currency either can be melted and cast into the appreciable shapes desired.  Silver wire can be purchased from jewelers as can platinum.

·        Misc. Electrodes:  Mercury(4), lead dioxide(5), rare earth oxides plated on titanium, tantalum, there are hundreds of possible electrodes that one may come across that are not covered here.

 

(1)

AC electrolysis is really a good way to fry electrodes.  Nothing will survive AC electrolysis of hydrochloric acid, even platinum will succumb to this, which is a good way to prepare soluble platinum compounds ironically.  Even carbon will be destroyed.

(2)

Older American nickels contain decent amounts of nickel but older Canadian coins contain an even higher amount.  The amount of nickel in coins in your country can usually be found easily by searching online.

(3)

Lately there has been a push in some areas to replace these lead items with the more environmentally friendly bismuth, at this time though products that have been replaced with bismuth usually proudly proclaim it as being environmentally friendly so determining what you have is still relatively easy, lead in these items is also usually alloyed to a small percent with other elements such as antimony.

(4)

Mercury electrodes are famous in the mercury cathode cell for the preparation of sodium hydroxide.  In this cell the mercury electrode forces the reduction of sodium cations to sodium metal rather then the production of hydrogen, the result is a mercury-sodium amalgam, containing a few percent sodium metal, this reacts with water slow enough to allow recovery of it, to look further at this phenomenon search ‘over voltage’ online.

(5)

Lead dioxide electrodes must be formed carefully, they have their claim to fame in the preparation of perchlorates by electrolysis of chlorates by the armature chemist.

3.4a Molten Salt Electrolysis

       Performing electrolysis on a molten binary salt usually yields predictable products providing you have a simple anion coupled to your metal ion.  Another advantage is that electrochemistry can produce many elements that would not be possible to produce under aqueous conditions.  The most common examples including producing the alkali and alkali earths electrochemically.  Of course the most prohibitive feature of molten salt electrolysis is the high temperatures usually employed.  The corrosiveness of molten salts and some of the products produced by the electrolysis, especially at the temperatures employed these factors usually result in molten salt electrolysis being beyond the reach of the beginning chemist.  However, many chemists would like to do molten salt electrolysis, and they can, the procedures outlined within this section will cover the basic points for this method of electrochemical production.

Working with Molten Salts

          Great care must be exercised when working with molten salts.  For the most part this is due to their high melting points.  Think about what happens when drop of water hits hot grease, the spattering and violence of the sudden evaporation of the water, now consider molten salts can be several hundred degrees Celsius hotter.  Water hits these salts like a bullet, sending liquid everywhere.  Due to the temperatures it can greatly dehydrate anything organic and it can also start fires on wood and paper and such.  It should also be noted that molten ionic salts are good solvents for…. Whatever, they attack a multitude of things and some they flat out destroy.  Specific examples like molten sodium hydroxide (above) destroy wood and such on their own, let alone at high temperatures, getting a quantity of this on your skin would be disastrous.  When working with molten salts your apparatus should be firmly placed on a flat surface, and once the salt is molten it should not be moved and you should limit how much you ‘mess with it’ as many mistakes can occur.  Wear welding gloves that can easily be removed if you get your salt on them and avoid inhaling the fumes coming off the melt too much, they can cause sickness and throat pain and loss of vision (as I once experienced).  Finally, when you are done with a reaction, allow the salt to cool to room teperature on its own, in case of emergency the salt can usually be dumped in sand.

       Nearly everything in your electrolysis cell will have to be metal due to the temperatures employed, although less used, ceramics may also be employed for some applications.  Some lower temperature cells may safely employ Teflon (not for use with molten alkali’s) or even some epoxies or cements.  The type of cell construction depends entirely upon the types of products produced.  The physical state of the products and their inherent reactivity after genesis is of the utmost importance and of equal related importance is weather the products thus formed will react with each other or even with the melt.  So here is your quick checklist:

1.      What is the best melt, and eutectic composition or a straight composition and at what temperature does it run?

2.      What products will be formed through electrolysis, if an eutectic mixture, will only one product be formed at the anode and one at the cathode?  Are you sure of both your anode and cathode products (check their potentials against charts(1)).

3.      At the temperature employed, what will be the physical state of the products formed?

4.      If the products are solid then they will not mingle, however if they are either liquid or gas, will they react back with each other and thus make it necessary to divide the cell or otherwise separate the products from one another?

5.      Will the products formed react with the melt(2), and if so, at what speed will the reaction occur?

6.      What corrosive properties might your melt have, testing though melting small portions of it on a small scale is usually necessary before making a full size vessel, if the metal can hold up to the melt for over an hour without pitting or significant loss of weight it is usually good, but products formed from electrolysis may still attack the vessel.

7.      How is the cell going to be heated to the melting point?  Flames involve additional precautions when working with cells that produce flammable gasses but are not out of the question.  Resistance heating, that is the heat generated though the electrolysis of the melt can supply a large quantity of heat and may maintain your cells fluidity although the melt must still be initially melted, optimally an embedded heating coil with a variable power setting will keep the cell temperature exactly where you want it at.

       Assuming you have two products that are going to be produced that are going to react back with each other (which is usually the case) there are several cell designs from which to choose depending on the physical form of the products and the density of liquid products compared to the melt.  The ultimate variable however in any electrolysis procedure is the electrodes.  The surface area of an electrode exposed to the melt, their distance from one another, and the amount of current flowing though them more often then not makes or breaks an electrolysis reaction.  Unless one finds highly accurate detailed sources for performing electrolysis on a molten salt they will not have the electrode distances or current densities.  And even if you do happen to find them, they usually employ apertures that are beyond what many normal people are capable of producing.  The only advice the author can give you on this subject is to try different things and experiment, that is what this is all about anyways, keep a journal so you can notice trends, cells such as the first one detailed pose the extra difficulty of the electrodes being stationary, the misplacement of an electrode in this circumstance is a tragedy, requiring shutting down the cell and possible part replacement, experiment with variable electrode cells before switching to the fixed electrodes for this reason.

Sample Design #1 (In Theory)

 A : This is the body of the cell, it is a reducing adaptor for plumbing, this one goes from 2 inches at the top to 1.5 inches at the bottom, both the top and the bottom are threaded.  B : This is a plumbing piece called a brushing, it reduces from 1.5 to .5 inches, the bottom hole being threaded.  C : This a plug, .5 inches in diameter that fits the hole in the brushing, these three pieces comprise the containment body of the cell.  D : This is the anode, it is inserted into a hole drilled into the plug ( C ) and does not touch the plug, thus the hole is somewhat large, if it touched the plug it would short out the cell and perform electrolysis elsewhere.  E : This is the cathode, it is inserted in the same way as the anode ( D ) the reason for it being shorter then the anode (see picture) is that in being shorter it has a higher current density, the cathode is where reduction happens, it’s where metals are formed, and just as a general rule of thumb higher current density gives better yields of metals.  F : This a material that the hollow part of the plug was filled with, to prevent the molten contents of the cell from leaking out, to hold the electrodes in place, and to prevent them from touching the base of the plug.  It can be any material inert to the cell, high heat resistant epoxy, fireplace mortar, concrete.  Just be sure to position the electrodes so they don’t touch the metal of the cell before setting them permanently in this design.  G : This a piece of metal that sets loosely inside the cell, it is rectangular with the sides of the metal touching the sides of the cell, but it does not extend to pass between the electrodes, if it did this it would short out the cell.  By being here it separates the products formed.

(In Practice)

                    

Removable Plug shown Removed          Overview of cell with Divider in Place w/ plug       

Additional additions to cell design : What this cell lacks is a heating mechanism. Flame heating is somewhat out of the question unless you have multiple heat sources, because you can’t heat from beneath and hot spots in an electrolysis melt are usually undesirable.  Wrapping the cell in an electric heating coil from a stove top coil range then embedding it within some refectory is reasonable for supplying most temperatures up to 500 °C.  Initially however it may be necessary to heat a quantity of your electrolyte separately over a flame to liquefy it, then add it to a preheated cell with the current ready to run, immediately upon the liquid entering the cell it will want to solidify unless the temperature is sufficient, however by the molten salt coming into contact with instantaneous electrolysis it should hopefully keep it molten(3), from there the cell can be fed periodically with solid electrolyte until the desired level has been reached.

What is this cell good for? : For one the products produced by electrolysis must be lighter then the melt (which is usually the case) and they must be liquids or gasses.  Plus they must not react substantially with the melt.  If the electrodes in the bottom are sufficiently distant from one another then products that would normally react with each other can possibly be obtained.  Upon their formation they rise though the melt and before they have a chance to move around much the metal divider prevents them from crossing over to the other side.  In this way efficient separation at the top of the cell is possible.  For example, if sodium chloride and calcium chloride are electrolyzed in this cell sodium metal will rise to the top above the cathode and chlorine gas would rise to the top above the anode.  The gas would require special handling, such as copper metal piping and such, and have to be taken care of with special precautions.  On the other side of the divider, liquid sodium metal would pool and be able to be ladled off.

Sample Design #2

Stripped down version of the Downs cell for sodium production from sodium chloride.

A :  Cell body inert or resistant to high temperatures and molten electrolyte.  B :  Electrolyte level in cell, rises above inverted funnel ( D ).  C :  Molten electrolyte.  D :  Inverted funnel located below the electrolyte level ( B ) and directly above the anode ( F ) to collect gaseous products ( H ) and move them away from the cell to deal with them separately and to prevent them from reacting with the cathode products.  Or it may be the intended product such as in a fluorine cell.  E :  Pipe leading away gasses from anode.  F :  Anode, inserted though bottom of cell and carefully positioned to avoid touching the metal lining of the cell itself, held in place with a resistant non-conductive material such as cement, some epoxies, silicates, mortar, etc. (In the Downs cell(4) this would be where the chlorine gas was produced.)  G :  Cathode (In the Downs cell this is where the sodium would be produced, there would be an inverted funnel above this too where the sodium could be drawn off during the electrolysis) the electrode is held in place same as the anode.

Sample Design #3

A :  Cell body, resistant to heat and electrolyte.  B :  Anode or Cathode held in place with a resistant material and set so as to prevent it from coming into contact with the metallic cell body.  C :  Direct heat such as a hot plate or torch can be more easily applied to this cell design.  D :  Anode or Cathode held in place with a resistant material and set so as to prevent it from coming into contact with the metallic cell body.  E :  Wire gauze cell divider.  F :  Electrolyte level. G :  Electrolyte

Sample Design #4

The most basic electrolysis cell.

A :  Cell body, heat resistant and resistant to the molten electrolyte.  B :  Direct heat can be applied from below in this cell.  C : Anode or Cathode, suspended above the cell by some mechanism and dipping into the melt a variable amount.  D :  Anode or Cathode, suspended above the cell by some mechanism and dipping into the melt a variable amount.  E :  Electrolyte level.  F :  Molten Electrolyte.

Classic Cell Example:

The Famous Castner Cell (Patent No. 452,030   May 12, 1891)

This cell runs with straight sodium hydroxide, originally run in an iron pot ( S ), and heated by a ring of flames ( G ), set in brickwork ( R ).  The sodium hydroxide ( A ), is melted and kept about 20 C above the m.p. (318C) of sodium hydroxide. The cathode made of nickel [or iron or copper] ( H ), comes up though the bottom of the pot.  Holding it in place is a cake of solid sodium hydroxide ( K ) cast into the cell when it was cool.  This cell has multiple anodes ( F ) that drop down around the cathode from above.  A cylindrical vessel ( N ), floats in the fused alkali above the cathode, and the sodium ( D ), and hydrogen liberated at the cathode collect here, the sodium is protected under the hydrogen atmosphere ( C ) which is allowed to escape and is burned off at the exit.  This are can be opened at periodically to remove sodium that gathers here.  Metal gauze ( M ), separates  the anode from the cathode.  Oxygen that is simultaneously produced escapes from vent ( P ).

 

Potassium hydroxide electrolyzed with nickel electrodes in a steel U-Tube

       Despite all these ideas a simple design such as above, where the anode and cathode areas are simply in a U-Tube (named such because it is in the shape of a U) affords enough separation to work for moderate periods of time, notice the distinct color difference between the anode (right) and cathode (left).  Such a cell could be easily constructed from materials at a hardware store and heated at any point where the melt solidifies directly. 

 

       Famous examples of cells using molten salts are abundant, but the most famous examples include; (1) The aforementioned Castner cell which uses a melt of sodium hydroxide to yield hydrogen gas, oxygen gas, and sodium metal.  (2) The Downs Cell(4) which uses an eutectic of sodium and calcium chloride to give sodium metal and chlorine gas. (3) The Hall-Hẻroult process whereby aluminum is produced from molten Al2O3 with cryolite to depress its melting point to a manageable 950 °C (manageable in industry at least).  There are other examples but these stick out among the rest.

 

(1)

Note, the electrolytic potentials in a molten salt are totally skewed with respect to the aqueous potentials found everywhere, these can only be used as an extremely rough guide.

(2)

An eutectic mixture of KOH and NaOH might be tempting to use for the preparation of the alkali metals due to its low melting point (~150 °C) however neither Na or K will be the product, instead you end up with a mixture of the two, this mixture, the sodium-potassium eutectic can spontaneously catch on fire at room temperature and is a liquid at room temperature, you don’t want to think about how it will behave at 150 °C with oxygen and hydrogen being produced all around it!

(3)

A note on this, the initial addition of molten salt to the electrodes can be a dangerous affair, without the electrolyte over the electrodes there is a period where the electrolyte added may only touch between the electrodes in a small space, when this is the case the electrolyte may pop and short out the cell, sending molten salt out everywhere, during an attempt at sodium metal this happened to the author from a small quantity of sodium being generated causing a small disaster, be warned when working with molten salts.

(4)

The Downs Cell uses a mixture of calcium chloride with sodium chloride to reduce the melting point of the sodium chloride from 801 °C to 580 °C, it contains roughly 33% NaCl and the remainder is CaCl2, more NaCl is added as the reaction progresses to replace the NaCl being electrolyzed.  The resulting Na metal contains a very small calcium metal impurity.  The actual eutectic between NaCl and CaCl2 lies at a different point giving an even lower melting point but the melt here is the melt used in the industry.  The Downs Cell is still the main method to produce sodium metal, replacing the Castner Cell quite some time ago due to NaCl being cheaper then NaOH, however the lower melting point of the Castner Cell makes it more attractive to the amateur chemist then the Downs Cell, coupled with the fact that no chlorine is generated in a Castner Cell.

 

3.4b Aqueous Electrolysis

By Tacho.

1- Ions:

         Water and other solvents called “polar solvents” do something interesting: their molecule has two poles, a positive charge on one side and a negative charge on the other. When you put a salt in water, the salt’s negative ions are surrounded by water molecules pointing their positive side to it. Of course, the positive ions get surrounded by water molecules pointing their negative side to it. The result is that the salt dissolves in a liquid “soup” of ions that we will call “solution” or “electrolyte” from now on.

          There must be enough anions to neutralize all cations. The whole thing must be neutral. Yes, it can built up charges, and momentarily have an imbalance, but mother nature will find a way to put thing back the way she likes it.

         In the solution, you don’t have the individual salts anymore. Just ions. So, if you put two salts in water, say... sodium chloride and potassium bromide, you only have sodium cations, potassium cations, chloride anions and bromide anions. You can extract from this solution the original salts or sodium bromide and potassium chloride! If you could substitute both of those anions for the hydroxide anion, than you could extract sodium hydroxide and potassium hydroxide from the same solution.

         That’s what this work is all about: the substitution of ions in aqueous solutions using electrolysis. We do that by pushing electrons in and out of atoms or molecules using electricity.  

2- How?  

        Suppose I want to get a copper salt, dissolve one of my electrodes. I mount the following setup, put distilled water in the flask and turn on the power:

        What do I get? Nothing! Why? Pure water practically does not conduct electricity! So we must add something that will make water conductive but won’t be part of the reaction. Let’s say we want copper sulfate, so let’s choose something with a sulfate ion attached to an “inert” ion. MgSO4 should work fine. Magnesium sulfate is Epson salt that every pharmacy should have. For now, take my word that Mg won’t be part of the reaction. Lets dissolve 3g of MgSO4 (the hydrate is okay.) in 30ml of distilled water and try again.

        Immediately you’ll see bubbles on the cathode and some blue tint by the anode. Cool! Copper ions should be blue! What is in those bubbles? The only thing in this soup that becomes a gas when it gains electrons (that’s what cathodes do, they give electrons) is hydrogen. Water molecules get broken and hydrogen bubbles away.

        But wait... after a couple of minutes we notice that something is very wrong! The blue tint is solid and it’s precipitating!  Copper sulfate doesn’t do that! What is happening?

        Here goes the explanation: When the hydrogen of water receives an electron, it joins a friend to become a gas molecule and bubbles away. It leaves behind the other half of the water molecule: OH- ion. This ion meets Mg+ ions or Cu++ ions and forms an insoluble hydroxide that precipitates (actually copper ions form more complex insoluble compounds, but let’s pretend it’s plain hydroxide). If you carefully put a piece of indicator paper close to the cathode while its bubbling, you will see that it’s alkaline.

        So how do we make our copper sulfate?

        The little snotty armchair chemist now smiles and says with his squeaky voice: “I know! I know! All you have to do is to use two half-cells connected by a salt bridge! It’s so simple! Like this:”

        In the real world, however, this setup is not efficient. If you use the 10g MgSO4 in 100ml water solution, under about 10 volts, you just can’t see any bubbles! That’s because ions are not agile movers, the salt bridge is a long way for them. Like any electric circuit, shorter paths increase current. So, a better design for the amateur would be:

        I tested this for the copper sulfate production and it works. For some reason, it works best using the lower part as the anode cell, where the sulfate forms. The solution by the anode is called anolyte and now is a mixture of Mg2+, Cu2+ and SO42- ions. How to separate pure CuSO4 from it? If you look up the solubility information for the two salts you will find that 100 ml of cold water will dissolve 31 g of hydrated copper sulfate whereas it will dissolve 71 grams of magnesium sulfate therefore if you ran your electrolysis long enough then heated to reduce the volume of the solution, then cooled, the first thing to come out of solution should be the copper sulfate, this is just one possibility, one of the difficult parts of electrolysis involves this problem applied to systems involving many more ions. 

In the above setup the cotton acts as a membrane to prevent the passage of certain ions.  Another type of system accessible to the at home chemist consists of a flowerpot, unglazed and previously soaked in a strong acid placed into a larger container.  The flowerpot and the larger container are filled so they contain the same level of electrolyte and an electrode is inserted in the flowerpot and another in your main solution outside. The electrons will migrate though the flowerpot and electrolysis will commence but your anode liquid will be separated from your cathode liquid, the strong acid soak necessary to facilitate the permeability of the flowerpot to ions, remember that strong acids are incredibly conductive.  This is yet another way to separate your anolyte to prevent extraneous reactions.

        I substituted the two copper wires of the first setup (the one with the connector) by two pencil leads (graphite) and put them in a pure copper sulfate solution procured from a gardening shop. When I turn on the power, quickly something starts building up in the negative electrode. That’s copper. It’s probably powdery and maybe too dark to look metallic, but it’s copper. That’s the principle behind electroplating. Don’t expect shiny metallic deposits though(1). What’s happening is that the Cu2+ cations in the solution gain a couple of electrons and become Cu, the metal.

3- Giving names to things happening.

        What is happening in the anode is called oxidation. It doesn’t matter if that there is no oxygen involved! Loosing electrons by the anode is called oxidation and don’t you argue! Usually this oxidation results in metals becoming ions, ions gaining oxygen atoms, oxygen evolving in bubbles or halogen ions becoming the element.

        What is happening in the cathode is called reduction. It’s usually a metal ion becoming a metal, ions loosing oxygen atoms, elemental halogens becoming ions or hydrogen evolving in bubbles.

4- What happens when?

        Question: What happens when I have many different metal pieces in the same electrolyte, connected and acting as anodes? Do they all go into solution at once? Do some of them go first? And what happens when I have many different metal cations by the cathode? Do they all get reduced together forming an alloy? 

        Answer: There is a priority list. Theoretically, if you have two different metal pieces, like copper and zinc connected to the positive pole of a battery and immersed in an electrolyte, all zinc will go into solution before the first atom of copper gets oxidized.

        Here is an incomplete list (potential table) that shows the tendency of an atom or molecule or ion to gain or loose electrons. If it shows a reduction potential of –2.71 for something, the oxidation potential for the same something will be +2.71:

Oxidized creature

Reduced creature

 

Reduction potential:

Li+

 

Li

 

 

-3.04

K+

 

K

 

 

-2.92

Ca2+

 

Ca

 

 

-2.76

Na+

 

Na

 

 

-2.71

Mg2+

 

Mg

 

 

-2.38

Al3+

 

Al

 

 

-1.66

2 H2O

 

H2(g) + 2OH-

 

 

-0.83

Zn2+

 

Zn

 

 

-0.76

Cr3+

 

Cr

 

 

-0.74

Fe2+

 

Fe

 

 

-0.41

Cd2+

 

Cd

 

 

-0.40

Ni2+

 

Ni

 

 

-0.23

Sn2+

 

Sn

 

 

-0.14

Pb2+

 

Pb

 

 

-0.13

Fe3+

 

Fe

 

 

-0.04

2H+

 

H2(g)

 

 

0.00

Sn4+

 

Sn2+

 

 

0.15

Cu2+

 

Cu+

 

 

0.16

ClO4- + H2O

 

ClO3- + 2OH-

 

 

0.17

AgCl(s)

 

Ag + Cl-

 

 

0.22

Cu2+

 

Cu

 

 

0.34

ClO3- + H2O(l)

 

ClO2- + 2OH-

 

 

0.35

IO- + H2O(l)

 

 I- + 2OH-

 

 

0.49

Cu+

 

Cu

 

 

0.52

I2(s)

 

2 I-

 

 

0.54

ClO2- + H2O

 

ClO- + 2OH-

 

 

0.59

Fe3+

 

Fe2+

 

 

0.77

Hg22+

 

2Hg(l)

 

 

0.80

Ag+

 

Ag

 

 

0.80

Hg2+

 

Hg(l)

 

 

0.85

ClO- + H2O

 

Cl- + 2OH-

 

 

0.90

2Hg2+

 

Hg22+(aq)

 

 

0.90

NO3- + 4H+

 

NO(g) + 2H2O

 

 

0.96

Br2(l)

 

2Br-

 

 

1.07

O2(g) + 4H+

 

2H2O

 

 

1.23

Cr2O72- + 14H+

 

2Cr3+ + 7H2O

 

 

1.33

Cl2(g)

 

2Cl-

 

 

1.36

MnO4- + 8H+

 

Mn2+ + 4H2O

 

 

1.49

H2O2(aq) + 2H+

 

 2H2O(l)

 

 

1.78

Co3+

 

Co2+

 

 

1.82

S2O82-

 

2SO42-

 

 

2.01

O3(g) + 2H+

 

O2(g) + H2O

 

 

2.07

F2(g)

 

2F-(aq)

 

 

2.87

 

        One should be able to get all the information one needs from this table, but lesser creatures like me always get confused trying to draw practical conclusions from it! So I organized the following two specific tables for things that can happen at your anode and things that can happen at your cathode from practical observations:

 

Things that can happen at your anode (where oxidation takes place, positive pole in the electrolytic cell) in order of priority:  

Priority

a.k.a. Oxidation Potential

in

Volts:

Situation:

What I think should happen:

What happened when I tried it:

higher than +0.76

If your anode has alkali metals, aluminum or magnesium ...

...Alkali metals react with water, and can’t be electrodes in aqueous solutions. Aluminum behaves strangely as an anode(2) magnesium will likewise react with the solution and little relevant electrolysis can be preformed with it.

+0.76

If your anode has zinc...

...Zn2+ ions go into solution. It may form insoluble salts (precipitate or form chunks in the anode) if it finds certain anions in the electrolyte like carbonate, hydroxide or hypochlorite. The Zn in the anode “dissolves” in the electrolyte.

It works. Nice and easy. Using a two half- cell setup with MnSO4, the zinc anode dissolves into Zn2+ ions, a colorless clear solution.  Insoluble impurities may show themselves at this point.

+0.74

If your anode has chromium...

...Cr3+ (chromic) ions go into solution. It may form insoluble salts (precipitate or form chunks in the anode) if it finds certain ions in the electrolyte like borate or cyanide. The Cr in the anode “dissolves” in the electrolyte.

+0.41

or

+0.04

If your anode is made of iron...

...Fe2+ (ferrous) ions go into solution at “priority” of  +0,41V, or Fe+3 at “priority” of +0.04V. It may form insoluble salts (precipitate or form chunks in the anode) if it finds certain exotic ions in the electrolyte. The Fe in the anode “dissolves” in the electrolyte. See -0.77 priority.

Funny things happen with an iron electrode. In my MgSO4 cell, at voltages lower than about 5 volts, a yellow/brown tint and a greenish precipitate appear due to iron ions(?), but if the voltage is above that, oxygen evolves from the anode and almost no color shows in the solution.

If a piece of zinc from a battery is strapped to the anode, almost no brown color evolves .Makes sense.

+0.40

If your anode is cadmium...

...Cd2+ ions go into solution. It may form insoluble salts (precipitate or form chunks in the anode) if it finds certain ions in the electrolyte like carbonate or borate. The Cd in the anode “dissolves” in the electrolyte, take note of cadmiums toxicity.

+0.23

If your anode has nickel...

...Ni2+ go into solution. It may form insoluble salts (precipitate or form chunks in the anode) if it finds certain ions in the electrolyte, like carbonate or borate. The Ni in the anode “dissolves” in the electrolyte.

+0.14

If your anode has tin...

...Sn2+ (stannous) ions go into solution. It may form insoluble salts (precipitate or form chunks in the anode) if it finds certain ions in the electrolyte. The Sn in the anode “dissolves” in the electrolyte. See -0.15 priority.

 

+0.13

If your anode has lead...

...Pb2+ (plumbous) ions go into solution. It may form insoluble salts (precipitate or form chunks in the anode) if it finds many common  ions in the electrolyte, like sulfate, chloride or carbonate. The Pb in the anode “dissolves” in the electrolyte. Another heavy metal you should be careful about.

0.00

If you make the incredible “hydrogen bubbling on platinum” electrode...

 

...Literature is unanimous that hydrogen would become an H+ ion, this is the standard by which this system of potentials was based.

-0.15

If your anolyte (the solution of ions around your anode)  has Sn2+ (stannous) ions...

...The Sn2+ is oxidized and becomes Sn4+ (stannic) ions.

-0.16

If your anolyte has Cu+ (cuprous) ions...

...The Cu+ become Cu2+ (cupric) ions.

-0.17

If your anolyte has ClO3- (chlorate) and OH- (hydroxide = alkaline medium)...

...The ClO3-  ion becomes the ClO4- (perchlorate) ion and some extra water is formed.

Perchlorate is the sweet darling of pyrotechnics. I have not done it, but the whole electrolysis sequence Cl- -> ClO- -> ClO2- -> ClO3- -> ClO4- is well described  on the internet(3).

-0.22

If your anode has silver and your anolyte has Cl- ions...

...Insoluble AgCl (Silver Chloride)  is formed.

The AgCl forms an insulating white layer on the silver piece that reduces the current sharply. It resembles white (well, maybe cream) paint that darkens if exposed to sunlight for long.

-0.34

or

-0.52

If your anode has copper...

...Cu2+ will go into solution. Or Cu+ at –0.52. It may form insoluble salts (precipitate or form chunks in the anode) if it finds certain ions in the electrolyte, like carbonate or hydroxide. The Cu in the anode “dissolves” in the electrolyte.

My results with a MgSO4 cell are described in the text.

-0.35

If your anolyte has ClO2- (chlorite) ion and OH- (hydroxide = alkaline medium)...

...The ClO2-  ion becomes the ClO3- (chlorate) ion and some extra water is formed.

See note for chlorate (oxidation potential –0.17).

-0.49

If your anolyte has I- (iodide) and OH- (hydroxide = alkaline medium)...

...The I- ion becomes IO- (hypoiodite) ion and some extra water is formed.

-0.54

If your anolyte has I- ion.

I2 is formed.

It does. In a potassium iodide solution, using pencil’s graphite electrodes, brown iodine color rapidly develops around the anode while hydrogen bubbles form at  the cathode.

Interesting note: If you apply 60Hz AC directly from the transformer (6V) into the electrodes, a yellow color slowly develops, showing this reaction is not fully reversible. At least mine wasn’t.

Also interesting: according to this list of priorities, a copper anode should not generate iodine, instead, the copper should dissolve into copper ions. That did not happen in my test. Brown color developed near the copper anode with no blue tint. Go figure!

-0.59

If your anolyte has ClO- (hypochlorite) and OH- (hydroxide = alkaline medium)...

...The ClO-  ion becomes the ClO2- (chlorite) ion and some extra water is formed.

See note for chlorate (oxidation potential –0.17).

-0.77

If your anolyte has Fe2+ (ferrous) ion...

...The Fe2+  becomes Fe3+ (ferric).

-0.80

If your anode has Hg Mercury...

...Hg2+ go into solution. It may form insoluble salts (precipitate or form chunks in the anode) if it finds the certain ions in the electrolyte. The Hg in the anode “dissolves” in the electrolyte. Dangerous heavy metal, by the way.

Mercury, besides being toxic, has some difficult to understand chemical behaviors such as its ability to introduce overpotentials and such.

-0.80

If your anode has silver...

...Ag+ go into solution. It may form insoluble salts (precipitate or form chunks in the anode) if it finds certain anions in the electrolyte, like hydroxide, carbonate, iodide, bromide or chloride. Silver has many insoluble salts. The Ag in the anode “dissolves” in the electrolyte.

-0.90

If your anolyte has Cl- (chloride) and OH- (hydroxide)...

...ClO- (hypochlorite) ion is formed. And some extra water.

See note for chlorate (oxidation potential –0.17).

-1.07

If your anolyte has Br- (bromide) ion...

...Bromine (liquid) is formed(4).

It does. In a sodium bromide solution, using pencil graphite electrodes, brown bromine color rapidly develops around the anode while hydrogen bubbles form at  the cathode.

This one is fully reversible: If you apply 60Hz AC directly from the transformer (6V) into the electrodes nothing seems to happen, although bromine and hydrogen are being oxidized and reduced 60 times per second!

Also: If you use a copper anode, no brown color develops, only a bluish precipitate (bromide? hydroxide?). This is expected, since copper is above the bromide ion in the priority list.

-1.23

whatever aqueous...

...At this point, in aqueous solutions, water molecules nearby the anode are split in O2 (gas), that evolves from the anode, and 4 H+ ions, that remain in the solution. Theoretically, no oxidation should take place beyond this potential, because water should get oxidized first. Ha ha! So much for theory! Something called overpotential makes, among others, the next two oxidations possible:

-1.36

If your anolyte has Cl- (chloride)...

...Chlorine gas evolves.

This one sure works. The electrolysis of a concentrated NaCl solution gives off hydrogen and chlorine. In theory, if you keep the voltage low enough, you can get a small amount of oxygen and no chlorine. But that may depend on the anode composition.

-1.49

If your anolyte has Mn2+ (manganous) ion...

...Mn2+ becomes  MnO4- (permanganate) ion.

I have not tested this one personally, but I read in a book a detailed description of a laboratory procedure to obtain potassium permanganate by electrolysis. The book seems trustworthy.

Things that can happen at your cathode (where reduction takes place, negative pole in the electrolytic cell) in order of priority:

Priority

a.k.a.

Reduction Potential

in

Volts:

Situation:

What I think should happen:

What was actually observed:

1.49

If your catholyte (the solution of ions around your cathode) has MnO4- (permanganate ion) and the solution is acidic...

...MnO4- turns to Mn2+ (manganous) and some water is formed.

1.07

If your catholyte has Br2 (liquid bromine)...

...Br- (bromide) ion is formed.

As I said before, if you apply 60Hz AC directly from the transformer (6V) into graphite electrodes in a sodium bromide solution, nothing seems to happen, although bromine and hydrogen are being oxidized and reduced 60 times per second! So, it works!

0.96

If your catholyte has NO3- nitrate ion and your solution is acidic…

...NO (Nitrogen Monoxide) gas and some extra water are formed.

 

0.90

If your catholyte has ClO- (hypochlorite) ion...

...Cl- (chloride) and OH- (hydroxide) are formed.

0.85

If your catholyte has Hg2+ ions...

...Hg (metallic mercury) is formed.

0.80

If your catholyte has Ag+ ions...

... Ag (metallic silver) is formed.

0.77

If your catholyte has Fe3+ (ferric) ions...

...Fe3+ becomes Fe2+ (ferrous) ion

0.59

If your catholyte has ClO2- (chlorite) ion...

...ClO2- becomes ClO- (hypochlorite) ion

0.54

If your catholyte has I2 (iodine) dissolved...

...Iodine  becomes I- (iodide) ion.

 

As I said before, if you apply 60Hz AC directly from the transformer (6V) into graphite electrodes in a potassium iodide solution, a yellow color slowly develops, showing this reaction is not fully reversible. But the fact that is a slow development shows that most iodine was oxidized and reduced 60 times per second. So it works!

However, a practical method of turning elemental iodine into I- is beyond my reach . I tested a solution of ethanol, iodine and MgSO4 and another with ethanol, iodine and NaCl. None seem to have any change under low voltage, At higher voltages, hydrogen evolved but no change in color. The anode was graphite.

0.52

or

0.34

If your catholyte has Cu+ (cuprous) or Cu2+ (cupric) ions...

...Both get reduced to metallic copper. Cu+ at 0.52 priority and Cu2+ at 0.34 priority.

This one is easy. But to get a shiny smooth deposit is an art(1).

As expected, when I did electrolysis using graphite electrodes in a solution of copper sulfate and zinc chloride, only copper deposited in the cathode (I tested it regularly dipping the electrodes in 20% HCl where copper does not react, but zinc bubbles). Only when almost (!) all blue tint from copper ions have gone from the solution, zinc starts depositing. Chlorine evolved from the anode, by the way, so by default, zinc sulfate was left in solution.

0.49

IO- (hypoiodite)...

...Becomes I- (iodide) and OH-.

 

0.35

ClO3- (chlorate) ion...

... becomes ClO2- (chlorite) ion

0.17

ClO4- (perchlorate) ion...

...becomes ClO3- (chlorate) ion.

0.15

Sn4+ (stannic)...

...becomes Sn2+ stannous ion.

0.00

If your catholyte is acidic and therefore has H+ ...

H2(g) is generated. Therefore, none of the following metal ions can be reduced in acidic electrolyte. Which makes sense, since these metals are attacked by acids. In fact, they are attacked by acids exactly because they have a lower reducing potential then H+, but that’s another story.

-0.04

Fe3+ (ferric) ...

metallic Fe deposits at the cathode.

No it doesn’t. At least not in my ferric chloride solution with graphite electrodes. I could get no deposit.  This could be due to the acidity of ferric chloride solutions though.

-0.13

Pb2+ (plumbous)...

metallic Pb deposits at the cathode.

-0.14

Sn2+ (stannous)...

metallic Sn deposits at the cathode.

-0.23

Ni2+...

metallic Ni deposits at the cathode.

-0.40

Cd2+...

metallic Cd deposits at the cathode.

-0.41

Fe2+ (ferrous)...

metallic Fe deposits at the cathode.

-0.74

Cr3+ (chromic)...

metallic Cr deposits at the cathode.

-0.76

Zn2+...

metallic Zn deposits at the cathode.

It works. The deposit from a zinc chloride solution is powdery and dark, but it’s zinc allright.

-0.83

Whatever aquoeous...

Water by the cathode gest split in H2(g) that evolves from the cathode and  + 2OH-, that remain in the solution.  End of the line. No exceptions and no reductions from here!

 

Aluminum ions, Magnesium ions and alkali metals ions...

Forget it! Not in aqueous electrolytes! Water gets reduced first(5)!

 

5- Practical notes for amateur experiments:

 

 

C) What can be made using aqueous electrolysis?

 

       Electrolysis can be inefficient and slow. Consumes lots of electricity and takes a long time to produce low yields that, in most cases, must be submitted to other procedures to isolate the pure products.  If you are hoping to make a bottle of bromine or iodine in one sunny Thursday afternoon, you will be very disappointed. On the other hand, it’s simple and within the reach of any amateur. It’ a good option when you need just a little bit of a chemical; instead of buying half a kilogram of an expensive, polluting and carcinogenic salt, make a little bit as needed. Also, it’s a way of obtaining chemicals that you just can’t buy.

 

·        Salts in general can be obtained like sulfates, chlorides, nitrates, chlorates and perchlorates;

·        Hydroxides and oxides;

·        Metal powders and metal deposits on surfaces or purification of a metal(6);

·        Gases like oxygen, hydrogen and chlorine;

·        Bromine and Iodine;

·        Dilute acids;

·        Organics like ethane, chloroform, etc.

 

 

A Small Mistake and a Learning Experience

 

 

       In an attempt to prepare chloroform a quantity of acetone was mixed with water and some sodium chloride.  The mixture was subjected to electrolysis using a carbon electrode and a nickel electrode.  The reasoning behind the experiment was simple at the time; (1) Electrolysis of the sodium chloride would give chlorine gas at the anode and hydroxide at the cathode.  (2) The mixture was being subjected to magnetic stirring so most of the formed chlorine would be dissolved and reacted with the hydroxide being produced to give a solution of sodium hypochlorite (the active ingredient in bleach).  (3) The formed sodium hypochlorite would react with the acetone present in a haloform reaction (5.3c) to give the desired chloroform.  This however was not what happened, instead the nickel electrode went into solution as the electrolysis progressed and ‘salted out’ the acetone.  This term is used to describe the lack of solubility of the nickel chloride in acetone causing the water and acetone to become insoluble in one another.  As a result the reaction that I hoped for had ceased.  (Notice the green color indicative of nickel salts in the bottom water layer which is denser then the acetone layer above.)  It should also be noted that chlorine gas can react with acetone directly giving chloroacetone, a potent lachrimator, which could have definatley caused me some breathing problems if formed in decent quantities.  So although the reaction sounded good at first, there were definitely some bugs to work out before I would be able to run it again.

 

    

(1)

As some have concisely stated, electroplating can be an art.  Developing a uniform coating of metal of significant thickness on nearly any substrate requires technique, and a decent investment in time and effort of calibrating a standard technique to work with  your equipment, providing you don’t just rush out and buy a setup specifically for electroplating. The more advanced form of this is a process known as electroforming, whereby an object is coated with a thick mechanically sound layer of metal, non-metal objects can be formed from wax or other materials then rendered conductive with some kind of paint then electrolyzed and coated with enough metal to make them useable for machining purposes.  However electroforming in the home lab is considerably more difficult then electroplating.

(2)

When aluminum is used as an anode certain characteristics are expressed that are beyond the scope of this text, information relevant to this can be found using key word combinations such as “Anodized Aluminum” or “Anodization”.

(3)

There are many places on the web covering the whole series from chloride to chlorate to perchlorate through electrolysis at home, searching is quite easy and rewarding, it can be done and people have succeeded in making large amounts of perchlorates in this manner, http://www.geocities.com/CapeCanaveral/Campus/5361/chlorate/chlorate.html is a nice site to start out at.

(4)

If the anode and cathode are separated even by a sheet of paper the bromine can be collected due to its density compared to water.  It will sink to the bottom of the apparatus and a small beaker can be placed there under water, the bromine will actually collect in the beaker.  For an example of the experiment look at http://www.crscientific.com/article-bromine.html

(5)

As stated previous, using a mercury cathode a mercury-sodium amalgam can be made containing a small amount of sodium metal, this however is more a curiosity as there is only a small amount of sodium, and winning it from the amalgam involves some dangerous manipulations such as distillation of the amalgam, that much mercury at those kind of temperatures can contaminate everything within a large radius, around an entire house!

(6)

A standard procedure industrially for copper but can be applied to other metals.  The metal to be refined is placed at the anode where it will dissolve in the solution as electrolysis prceeds, if the metal has a fairly low reduction potential then it will also re-deposit at the cathode before other impurities in the anode will carry over, details can be found far and wide on the internet through simple searching: http://doccopper.tripod.com/copper/er.html

 

3.4c Non-Aqueous Electrolysis

 

       Sometimes water as a solvent just won’t cut it.  Rarely this is due to the limited solubility of a compound in water, more often then not electrolysis of a substance in a non-aqueous environment is necessary due to the reactivity of some compound that is being solvated or a compound being produced with the water in which it is dissolved, therefore excluding it as a solvent.  Usually in these cases it is some metal that is being formed, the alkali metals being the prime example, that would react with the water, but it need not always be a cation that is being reduced, shown here is a picture of a cell for the preparation of fluorine, not only would the preparation of fluorine be inhibited by its high potential in aqueous solutions, but it would react readily with water, therefore the electrolysis of a salt, usually potassium fluoride is done in anhydrous hydrogen fluoride.

 

Moissan's original cell for electrochemical fluorine production.

The actual electrolysis takes place in the cell marked B.  Contained within was a U-tube made of platinum metal.  The stoppers on the top were composed of shaped calcium fluoride and the anode and cathode were both platinum.  The cell ran on a mixture of potassium fluoride and hydrogen fluoride at low temperatures.  The temperature of the cell being maintained by the methyl chloride that entered the area surrounding the U-tube.  Jar C also contains methyl chloride which cooled down the resulting gasses to condense out any hydrogen fluoride that may have made its way through.  The bulbous parts ( D ) and ( E ) contain sodium fluoride, which forms the acid salt NaF*HF when in contact with hydrogen fluoride to remove the last of the hydrogen fluoride.  The whole tubing system leading from the fluorine production area was composed of platinum in this original setup.

 

       When it comes to solvents available for use, the requirements are a little more stringent then just any liquid available that contains water.  As a matter of fact finding a good non-aqueous solvent is near impossible, at least one that has the wide range of possibilities as water, therefore selection of a solvent is highly dependent on what you want to accomplish, and each solvent that you have the option of using has its own reactivates and conditions that it best works under that must be taken into account [Note that the reductive or oxidative potential of anions and cations is influenced by their solvent and therefore the table in the previous section on aqueous electrolysis cannot be followed strictly for electrolysis in non-aqueous mediums].  The solvents that do work for non-aqueous electrolysis well are hard to come by, and few of them are over the counter, still, here is a brief listing (Note that when it comes to availability, the reference is to the availability of the anhydrous substance, and of course the availability of substances depends heavily on your country):

 

[Solvents highlighted in blue have some merit of being used in electrolytic procedures beyond the curiosity stage.]

 

Liquid

Availability

Properties

CH3CN Acetonitrile

Somewhat Difficult to Acquire

Not as toxic as one would infer from the cyanide group, it is widely studied for use as a non-aqueous medium, it dissolves a nice variety of ionic compounds to some extent.

H2SO4 (Sulfuric Acid)

Somewhat Available

The usefulness of sulfuric acid for non-aqueous electrolysis is limited, more often then not, it is the reagent in a technique to produce peroxysulfuric acid or ozone.

(CH3CH2O)2CO Ethyl Carbonate

Varies Greatly

Coming into its own as a solvent in lithium ion batteries, somewhat inert and works well for increasing the conductivity of a solution.

HF Hydrogen Fluoride

Very Difficult to Acquire

Hydrogen fluoride is incredibly dangerous and preparing it at home in the anhydrous state would be ridiculously dangerous, it finds very limited use except as a solvent for the preparation of high oxidation state compounds (See Moissan’s cell at the opening of this section).

SO2 Sulfur Dioxide

Somewhat difficult to Acquire (Could be made)

Sulfur dioxide is a widely investigated non-aqueous solvent.  It is however a gas at room temperature and is very irritating to work with. 

CH3SOCH3 Dimethyl Sulfoxide (DMSO)

Somewhat Available

Displays a wide solvating ability and conducts fairly well, although somewhat easily attacked by oxidizing agents, still a good non-aqueous solvent.

NH3 Ammonia

Difficult to Acquire (Could be made)

Ammonia is a gas at room temperature so cooling is necessary to work with it.  Ammonia solvates many things including the alkali metals, it can also produce them through electrolysis of their salts, along with other interesting things, such as hydrazine.

HCOOH Formic Acid

Difficult to Acquire (Could be made)

Has a significant polar character and will dissolve many things that water will, finds limited use in the non-aqueous electrolysis field.

CH3COOH Glacial Acetic Acid

Difficult to Acquire

 

HCON(CH3)2 Dimethyl Formamide (DMF)

 

 

C5H9NO

N-Methylpyrrolidone

 

 

C6H5NO2 Nitrobenzene

Difficult to Acquire (Could be made)

Nitrobenzene is slightly more polar then most aromatics and is therefore more useful then many in electrolysis.  It has the famous reputation of allowing for the preparation of the alkali metals by electrolysis of an alkali halide fused with aluminum halide.  The downsides of this solvent being the acquisition of it coupled with its toxicity.

C5H5N Pyridine

Difficult to Aquire

Somewhat able to dissolve ionic compounds, the classical example of its abilities is the preparation of lithium metal by electrolysis of a solution of lithium chloride in pyridine(2).

(-CH2CH2CH2CH2O-) Tetrahydrofuran THF

Can be acquired and purified from OTC sources.

THF is used in some OTC applications involving PVC pipe and can be purified from there.  It does well in organometallic applications because it has a somewhat pronounced polar character and thus works well for fun with electrolysis.

(CH3CH2)2O Diethyl Ether

Difficult to Acquire (Could be made)

Shows only a small solvating ability, not a good choice for electrolysis as many dissolved salts do not conduct.

CH3NO2 Nitromethane

Difficult to Acquire

 

CH3COCH3 Acetone

Widely Available

 

CH3CH2NH2 Ethyl Amine

 

 

Aromatics (Toluene, Xylene, Benzene)

Widely Available

Aromatics are had to work with in that they hardly dissolve anything considered ionic, let alone make conductive solutions, but things can be done(1).

 

 

Electrolysis of lithium chloride in DMSO with Nickel Electrodes

(1)

Although if something is solvated in an aromatic hydrocarbon it really doesn’t contain any ions, electrolysis of aluminum bromide in toluene with an amine to take up the liberated bromine will give deposits of aluminum metal, see US patent 3,997,410 for an example.  This is due to complexes generated between the aluminum and bromine giving compounds of the sort AlBr4+ which can carry the current through the solution leading to productive electrolysis.

(2)

In practice this electrolysis from lithium chloride to lithium metal is very sensitive and does not give good results, the solutions are only slightly conductive as well, facilitating the need for very large surface areas on the electrodes.

 

3.5 Titration 

        A titration is perhaps the cheapest and most useful quantitative test that can be done in home chemistry and can serve in a variety of circumstances for analyzing a plethora of products. Also remember there are many different types of titrations, ranging from acid-base, to redox, to iodometric, and many more. It is important to select the right method to get the fastest and most accurate results. It is most useful when you want to know the concentration and purity of a given reagent, for instance, hydrochloric acid. However, before you could do this, you still need to know the basics and have the necessary materials.

Materials:
To do an accurate titration, you will need:

A small glass funnel to facilitate filling the buret
A glass buret (50mL in 0.1 w/ PTFE stopcock is recommended)
A standardized titrant
A stirrer plate with PTFE magnetic stir bar
An indicator suitable for the pH range
An accurate balance for weighing dry materials
Buret Clamp
Ring stand


 The Basics


       Now that you have all of the materials, it is time to understand the some of the jargon associated with doing a titration and some basic procedural tidbits. Concentration is usually expressed with molarity or normality. Molarity (M) is moles per liter of solution, while normality (N) is a function of equivalents ( I.e. molarity * n, where n is an equivalent/factor determined from a reaction). Molarity is by far the more commonly used.

Equivalence Point-  The point at which stoichiometrically equivalent amounts of reactants have reacted.

End Point-  The point at which an indicator changes its color and the completion of the titration, this is not necessarily the same as the above. Now is when you take your final number.

Starting Point-  The point at which the titration begins. Now is when you take your initial number.

Midpoint (Half Titration) -  The point at which one of the reactants is half-reacted. In some cases (with a weak acid or base) this pH correlates to the pKa or pKb of that weak acid/base.

Titrant-  The reagent used which is usually of a known concentration and is in the buret.

Standardization-  The process of making a standard solution for further use. This is normally done to verify the concentration of a known solution. (ex. .0216g of an organic acid is titrated with an unknown M concentration of NaOH).

Useful things to remember that can lead to more accurate results:

Keep your buret clean—Always rinse twice with distilled water and then three times with your titrant.  Remember, accurate results depend upon the purity of your reagents and lack of contamination!

Use an electric stirrer—Many people have trouble swirling the flask with one hand and manipulating the stopcock with the other.  And electric stirrer remedies this and makes it easier to concentrate.

Use a white piece of paper underneath the flask—Titrations are able to give results of great accuracy.  To obtain such accuracy, one must look for the first indication that the reaction is finished, and since one is using an indicator, this usually means a color change.  For instance, say phenolphthalein is being used, it is imperative to stop when the slightest pink color is observed.  The contrast of the solution with the white paper underneath aids this.

Fill the tip—Before you take your first initial number, fill the tip of the buret with the solution by opening the stopcock.

Once, twice, three times, and four!—The first titration gives you an indication of what to expect, but irregardless, proceed with caution!  The second, third, and fourth titrations are usually within +/- 0.1 ml.  More titrations mean more accuracy when determining your concentration.  It is suggested to exclude the first titration, and take the mean of the others as your number; finding the standard deviation is also helpful in a quantitative setting.

Pick the right reaction—It is important to select the right type of titration for the job.  Determining the concentration of a KMnO4 solution requires a redox reaction and is different from determining the concentration of HBr, an acid-base reaction.  Even determining the concentration of halide ions (I.e. an iodometric titration with a starch indicator.) can be done with the right reaction.

Other considerations


       For the analysis and interpretation of your results, graphs are invaluable. Much of the calculations are basic solution stoichiometry, however, remember, that is only as good as your reaction coefficients (balance well!). As far as standardized solutions go, time was when the home chemist could order a supply house for standard NaOH soln. and whatnot. Now, that's not the case; your best bet is to standardize your own solution. To do that will require knowledge of buffers and acid-base reactions.       When someone says titration the first thing that comes to the average persons mind is determination of acid concentration.  The regular acid/base titration to determine acid concentration in a solution using a base or acid of known concentration is a useful and widely known titration method.  However titrations can be used to determine other things, such as metal ion concentrations and oxidizing power of a solution.  Titrations at home are possible and with just one moderate investment you will be able to perform a variety of useful titrations for some time to come.

       In case you are completely new to the concept of a titration here is a complete overview of the process.  Let’s say that you have bought some battery acid from your local auto supply and you have no clue of what the concentration is.  Reading the label you realize the ingredients are sulfuric acid and water, no added products.  So, how much sulfuric acid do you have?  In an acid base titration you would take your acid of unknown concentration and react it with a base of known concentration until the acid has been completely reacted (or vice verse with a base of unknown concentration).  There are two key points to consider, your standard, and your indicator.

 

The standard solution

          Your base or acid that you are reacting against your unknown should be of a known concentration, this solution of a known strength that you are basing your calculations off of is known as your standard solution.  A standard solution must either be purchased from a chemical supplier of prepared.  Many liquids over the counter can vary in their acid of basic strength for one reason or another so you cannot take them at label value and can therefore not use them as standard solutions.  For example, the amount of ammonia and OTC ammonium hydroxide can vary by several percent, it evaporates, the plant that made it had a slight variation with your batch, these things are meant for cleaning, not for acid-base titrations after all.  So the best thing to actually start with is usually a solid product that when dissolved in water will yield an acid or basic solution whose concentration can easily be calculated.  The solid should be something non-hydroscopic or at least have a known water content (i.e. known water of hydration) better yet it should be dried shortly before preparing the solution if you are unsure how much water it may contain, not at exhaustive temperatures but maybe 150 °C for a few hours, assuming the compound will not decompose by then. 

       Most solids that you will come across are only weakly basic or acidic, for example, sodium carbonate of sodium hydrogen sulfate.  However for an acid-base reaction to be complete and to give good results you should always titrate with at least one of the titratrants being strong, either a strong base or strong acid.  Therefore the first thing you should titrate against should be a strong acid or base, assuming for example you decide to go with sodium bicarbonate as your primary, it would be wise to titrate it against some over the counter hydrochloric acid, in this way you will get better results and you will additionally have standardized a strong acid for use in standardizing other bases. 

       Now, to prepare your standard you are going to need a scale and a graduated cylinder or another container that can measure liquid fairly accurately.  If your scale doesn’t measure that accurately, i.e. no decimal placed, you might want to make a large amount of solution in order to minimize the relative errors.  If your are desperate you can find the solubility of a substance at room temperature and saturate a solution with it at 25°C then decant off the saturated solution, this will contain a known amount of your solid, however it may deposit solid on cooling and therefore many induce error so the solution should be used immediately to standardize a different solution.  With your standard solution prepared you are part way to the desired goal of being able to determine the concentration of an unknown.

Sample Standard Preparation of a Sodium Carbonate Solution

Taking your grams of sodium carbonate you would put it into a flask that is able to measure volume exactly, a volumetric flask or if all you have is a graduated cylinder, that (Note, you cannot use a beaker for this, beakers can be significantly different then the volume expected).  So, lets say that you weigh out 25 grams of anhydrous sodium carbonate and you dissolve this in water and bring the volume up to total 250 ml.  25 g Na2CO3 / 106 g/mol = .235 mol then you have 250 ml of solution which is .25 L, dividing the mols Na2CO3 by the liters you get mol/L which is M and therefore you have your solution of known concentration, in this case it being .94 M Na2CO3.

Indicators

       How do you know the acid has been completely reacted.  With the Na2CO3 you can get a very inaccurate idea of when the reaction is done by watching to see when it stops bubbling.  However the best method would be to add a chemical indicator.  There are many acid-base indicators, but for going from an acidic solution to being able to tell when it’s basic the quintessential indicator is phenolphthalein which would have to be purchased, it is used in a dilute solution.  In the presence of acid or in neutral conditions solutions containing this organic molecule are clear, but around a pH of 8 the solution sharply turns pink.  (Note that only a few drops of indicator in dilute solution are needed). 

       There are naturally occurring indicators, a number of natural extracts will change color depending on the pH of the solution and could therefore potentially be useful for home preparations of acid-base indicators.  The most famous of all of these natural indicators is the Red Cabbage Indicator(1), that specific phrase pulling up many many hits on google some specifically on the preparation of the solution.  Because a color change can be relative though it would normally be practice, when using such an indicator that undergoes such a wide color range to use a standard solution of known acidity of bascisity and subject the indicator to it, that way you can compare that color with the color of your solution that you are titrating until they are equal.  However for a simple determination these small details can be ignored.

Performing the Titration

       Your equipment in this operation is a piece of glassware called a buret.  It is most easily described as a tall thin buret with a stopcock on the bottom where the liquid comes out.  It is filled to a specific level with the liquid you are titrating with and a flask containing a known amount of your unknown solution with a small amount of indicator is put below it.  The stopcock is turned and the flask is slowly rotated to stir while the titrating liquid enters your unknown.  Once the color starts to become apparent in areas the flow of the titrating liquid should be lessened and it should be added as drops until the liquid turns to its indicator color and stays that way for 30 seconds.  Congratulations, knowing the initial volume of liquid in your buret and the final you can determine the amount of liquid of a known concentration that it took to react with your unknown.  For example, let’s use the .94 M Na2CO3 solution I mentioned above and let’s keep with our sulfuric acid example.

       The initial volume of the .94 M Na2CO3 in your buret was 0.00  (They start at 0 and increase as they go down) and the initial volume of sulfuric acid in a beaker beneath it is 20 ml.  On your first attempt 43.10 ml of Na2CO3 was needed to turn the indicator pink.  On your second attempt 41.2 ml was needed to do the same, and on your third attempt 41.27 ml was needed (Multiple titrations followed by averaging of results usually gives better results, the less accurate the lab equipment the more titrations necessary to off set it).  First, what is the reaction?

Na2CO3(aq) + H2SO4(aq) Þ Na2SO4(aq) + H2O(l) + CO2(g)

       Although the process is complicated by H2SO4 being diprotic so there are two protons that must be neutralized and because of this complete conversion to Na2SO4 cannot be totally confirmed, it is somewhat safe to assume in this case.  So the stoichiometry is 1 : 1 so the number of mols Na2CO3 used equals the number of mols H2SO4 in solution.  If this had been hydrochloric acid instead the ratio would be 1 mol sodium carbonate to 2 mols HCl so the number of mols HCl would be double the number of mols Na2CO3 in the amount of solution used.  But back to the example.   Here are the calculated molarities.

(Molarity of titrating solution) x (Volume of titrating solution in ml) x 1/1000 (converts volume in ml to volume in liters) = mols titrating solution

[(mols titrating solution) x (conversion ratio [1/1 in this case])] / (volume of solution being titrated in ml x [1/1000])

So we get:

[.94 M Na2CO3 x 43.1 ml x 0.001 x 1 (conversion ratio)]/(20 x .001) = 2.02 M H2SO4

[.94 M Na2CO3 x 41.2 ml x 0.001 x 1 (conversion ratio)]/(20 x .001) = 1.94 M H2SO4

[.94 M Na2CO3 x 41.27 ml x 0.001 x 1 (conversion ratio)]/(20 x .001) = 1.94 M H2SO4

       Adding them all together and dividing my three we get an average molarity of 1.96 M so bottle it and save it for later now that you know.

       Some of you might be thinking now, why sodium carbonate?  Well, this is a good choice for the at home chemist for a few reasons 1) It is usually a somewhat pure product, especially if heated first to drive off excess moisture and decompose bicarbonate 2)  It is readily available  3)  It is measurable without difficulty   But there are some aspects that it lacks that would make it ideal:

1)      Sodium bicarbonate is not a strong base and therefore cannot react completely with an acid :

This is a bad thing, only a strong base can completely titrate a weak acid, however sodium bicarbonate decomposes to carbon dioxide which drives the reaction foreword, to ensure that every last bit of CO2 has been driven out of solution though and ensure the most accurate endpoint possible immediately after a color change is noted that stays for more then a few seconds, it helps to heat the solution being tested to drive off dissolved carbon dioxide which causes a slightly acidic solution.

(1)

The Red Cabbage Indicator has a wide range or colors that it changes between from a pH of 0 – 14, making it useful for quite a few operations, but it is sensitive to oxidizing conditions as well as keeping it at high or low pH, also solutions of it tend to break down over time but the reagents for its preparation are cheap and available, one site detailing the preparation of the indicator along with it color ranges can be found at http://chemistry.about.com/library/weekly/aa012803a.htm as can more information on acid-base chemistry.

3.6 Temperature control/Measuring

        Temperature control varies in its importance from reaction to reaction.  Two extreme examples would be a nitration reaction where you are controlling the temperature of the reaction to a range from 10C to 30C and at the total opposite might be trying to make phosphorus where your temperature might be 1200C and you are trying for as high a temperature as possible.  Both of these reactions pose their own difficulties for both measuring the temperature and controlling it.

        Your basic reactions are readily controlled with water in some form or another.  Water has a high specific heat and it can absorb a lot of heat before rising in temperature, or conversely, it can hold enough heat to warm another mass significantly before loosing its full heat.  So a cold water bath might be good for cooling, a warm/hot water bath for heating, and ice is always good to have around just in case.  However the other time tested method of cooling mixtures even cooler are eccentric mixtures.  KOH with ice can achieve extremely low temperatures, HNO3 and ice can too.  As can the dry ice/acetone slurry that is occasionally used.  Even cooler mixtures can be obtained with hydrocarbon baths with liquid nitrogen added periodically.  Liquid nitrogen could even be used directly, much lower then this is hard to accomplish in a home lab, but the lowest temperature coolant you will run across would be liquid helium, but I doubt you'll find this in any local super market.

The Cooling Effect of Winter:

Many of the people who read this text will experience some form of winter cooling and people further south and north of the equator will experience even further cooling and snow.  Winter is a good time to perform those exothermic reactions that would usually require external cooling or precooling of the reagents, just sit them outside.  Many chemicals are also less volatile under these conditions, bromine is more friendly the colder you get so if you were trying to make somewhat large quantities of bromine winter would at least make it more manageable.  The preparation of chloroform from acetone and sodium hypochlorite is aided by allowing the reagents to come to temperature outside then mixing and cooling in a snow bank, it’s cheap and it’s effective.

There is one little catch about the snow though that many people often forget, it may be cold but it’s also an insulator.  Where the snow is touching the container to be cooled it is doing it’s job, but in short order if it was actually warm enough to warrant external cooling then the snow around the vessel will melt and make a nice insulating air pocket.  For this reason snow is best used as a slush around the reaction vessel to be cooled with occasional stirring, in this way heat is conducted further away faster and there is no insulating effect.

        On the opposite side of the temperature scale you are shooting for heating.  Most any heating you will be doing will be the work of either electric heating, as in a hot plate, or chemical heating in the form of combustion.  Commonly electric heating concepts can get to 250C or so, lab grade hot plates can get even higher.  But to get really high you will have to use combustion, lacking a suitable furnace that is.  There are different kinds of torches, and different kinds of burners, Meker burners, Bunsen burners, and more each have their own limitations.  Common butane torches can reach 700 - 800C but to get higher the use of MAPP gas can bring you there with an appropriate torch head.  Methane is also a good carrier of potential energy so hooking directly into your home gas line has its advantages.

        Measuring these temperatures though pose their own difficulties.  Common thermometers are fine for common temperatures, -30C - 300C can be found on many thermometers such as candy thermometers, but beyond this on either side of the temperature scale it becomes necessary to deviate from the norm.  Bimetal thermometers that hook into electrician tools can reach 1000C but you will have to invest in a good probe and the thermometer itself is sensitive to chemical attack.  There are three different types of probes, the most common is the K type, these can cost quite a bit, and you need to also purchase the reader to plug it into, and not all K type probes are the same, they have different ranges, always stay within the manufacturers temperature recommendations to prevent destroying your expensive probe.  Another type of thermometer useful for high temperature is the infrared thermometer.  Just point and click and you get an non-invasive temperature of up to 1000C for upper class models, but you will pay high for this too.

        The cheapest, and most easily accessible tool for high temperature measurement available to the amateur chemist is the melting point of other compounds.  Molten metal baths can give an approximation of the temperature being used to keep the bath molten.  Different substances can be found to give a wide variation of temperature baths, just be wary of decomposition problems at these temperatures.  These can be used to at lest get an estimate of the lower end of the temperature present.

3.7  Removing water from gasses/solids/liquids with drying agents

       Referring to the section on dehydrating agents and desiccants, section 4.6 you can find a list of compounds that are good for removing water.  Both dehydrating agents and desiccants (which are also known as drying agents) can be used to remove water from a system, however the action of dehydrating agents is strong enough that it can attack chemical bonds (beyond any normal chemical reactions that may be associated with a compound) and this extra reactivity should be considered.  Drying a chemical is a common procedure for matter in all of its common forms.  The basic premise is to find a compound that really likes water, more so then the chemical you want to dry, and adding that chemical in such an amount to tie up the water present, then removing the desired compound from the desiccant which has since been used to remove the water.  Some chemicals hold water very strongly and their dehydration proves to be very difficult, but generally procedures for removing water are simple and can be generalized as such:

            Removing water from a solid:  The most common method to remove water from a solid is to heat it to drive off the water or use heat in combination with vacuum.  But this is not always be the case, and when a compound is subject to decomposition from heating, drying agents can be a reasonable thing to try.  The same method to keep a solid dry can also be used to dry a solid that already contains water, the use of a desiccator.  A desiccator is just an air tight container, with a desiccant (drying agent) at the bottom and an area suspended above the drying agent where a solid sample can sit.  As the sample sits above the drying agent it looses water to the environment, in which almost no water is present, this water is then absorbed by the desiccant at the bottom.  This process can occasionally take weeks, so patience is a key to a very dry sample.  A less used method is also available to remove the water from organic compounds involving dissolving in the minimum of hot organic solvent in which water is insoluble, it will then, in theory, float to the top where it can be pipeted off and the solvent then removed by evaporation.  Of course using non-aqueous environments and anhydrous reagents can ensure water does not get into the reaction environment at all and therefore not in the product to begin with although this can be next to impossible for some reactions.

            Removing water from a liquid:  Once a majority of water has been removed from a liquid usually through careful distillation or washing a fairly insoluble liquid with a saturated salt solution, drying agents can be good for getting out that last percent.  The usual procedure is to take the liquid and add a bit of drying agent that is insoluble in it to the liquid.  It is then stirred and if the drying agent at the bottom seems clumpy then more is added until the drying agent is free flowing, like dry sand between your hands.  That is when most of the water has been removed.  In some substances that are exceptionally hydroscopic though this process can take quite a while, the drying of ethanol with anhydrous copper sulfate can take days or weeks for example.  The drying agent in this case does not have to be a solid, sulfuric acid can be used to dry some solvents that do not react with it and are not soluble in it such as bromine and acetonitrile.

            Removing water from a gas:  Gas drying tubes are available for plugging on the end of distillation apparatuses.  What they are is a short length of wide glass tube, followed by a bulge that tapers down to thinner glass tube.  The thinner part usually is attached to the exit/entry point for the gasses in the apparatus by a small length of rubber tubing.  A piece of cotton is crammed in the larger part and the bulbous part filled with a desiccant such as calcium sulfate and another piece of cotton pressed on top of that, gasses entering the apparatus are thus treated to remove some of the water.  In gas generation removing water from the gas produced is sometimes refered to as scrubbing (see the section on gasses), in this case the gas is bubbled through concentrated sulfuric acid or through a hygroscopic solid to remove water, sometimes three and four times before the gas is ultimately consumed.

3.8  Recrystalization

        When the purity of a product is in question and you can spare a little bit in the quest for a more absolute product something can usually be worked out with recrystalization.  The basic principle is to pick out a solid that you want purified, say, ammonium nitrate.  The next step is to find a suitable solvent for it.  The solvent should posses the following properties:

  1. Be able to dissolve a large quantity of the desired product when hot and have only a low solubility when cold.
  2. Not affect the product be it by causing it to decompose or react with it.
  3. Be useful under atmospheric conditions, not possess properties that are affected by heating and cooling.

        Finding such a solvent is usually difficult.  Many places do not list the solubility of a substance in anything but water, let alone finding hot and cold solubilities of a substance of different solvents on the same page.  As such some trial and error may be involved, or research can help, finding out what solvent a pioneer in the field used to recrysatalize one of your products may be a good start.

        Once your solvent is picked out, in the case of ammonium nitrate water can be used.  The first step is to heat the water to a high temperature but not quite boiling, then to saturate the solution with as much of the solute as can be dissolved, if there is still solid solute in the solution either it can be spooned out or more solvent can be added.  After the substance is all dissolved and while still hot a quick filtration can be used to remove insoluble materials such as glass/sand particles.  Careful though, a hot saturated solution can crystallize on a cooler solid surface quickly and can plug filter paper, have something to scrape the filter paper with handy.

        After your hot solution has been quickly filtered you let it cool.  Usually once a certain temperature is reached crystals will automatically start to come out of solution.  However on occasion a solution may become super-saturated, i.e. the solution should have crystals forming but there is nothing for them to form on, that is one way to look at it.  In this case one of two things can be done, a crystal of the original compound can be added, this is called a seed crystal, and new crystals will grow off it.  Or you can scrape the inside rim of your container right at the liquid air interface, this agitation can cause the growth of crystals.

        Allow the solution to keep cooling but don't mechanically cool it too low, if for example you cool a water solution to near 0C then most of the impurities may crystallize out with your intended product defeating the purpose of recrystalization.  But after you get a good crop of crystals then filtration is the logical next step.  Vacuum filtration is the best as you may be filtering off a large quantity of solid but gravity filtration may work depending on your circumstances.  Your freshly grown crystals should be washed while still in the filter with a few quantities of cold solvent to remove adhering particulates.

        Now that you have your crystals they may need to be dried, in a desiccator or under high vacuum are the two normal choices, the desiccator being the most readily available to the amateur chemist.  But this step may not be necessary, check in a chemistry book to see if the salt you produce is hydrated, if that is the case heating the salt will usually be necessary to create the anhydrous product, if your product is disquecent or hygroscopic a desiccator may be a good first choice, keeping it there for a few days may prove to be a good move, followed by immediate storage in an air tight container to prevent the re-entry of water.

        In addition to cooling a solution to cause a precipitate of crystals another solvent can be added to the solution.  In this case the solubility of a solid is less in both the solvents then it is in either one alone.

3.9  Measuring Weight and Volume

       Quantitative chemistry requires one to know the amounts of reagents that are to be mixed for each reaction.  Measuring weight is most easily done with a digital scale, those used for weighing mail are well within the price range of most chemists and more expensive ones are available if you want to splurge.  A good scale should at least go down to measure by the gram, but down to the first decimal place is much better.  Such scales usually go up to one kilogram or so and therefore cover a wide range of useful measurements.  However if a digital scale is not available to you, there is always the old standby of a good old-fashioned balance.  Using a balance requires a set of weights or other objects of known mass to balance against.  The substitutes for a weight set are many but need to be a series of objects that are of a known mass and are nearly identical to one another, even something such as sheets of paper can work, or metallic currency, the weight of which is available with some searching online for each country.  Knowing this one can make a rudimentary scale and balance against a known weight of currency and obtain the correct weight of your desired reagent. [Note that water at room temperature is roughly 1.0 g/ml and as such since it is easier to measure volume you can make a counter weight in a balance scale to the exact specification you want by adding the correct amount of water.]

       Measuring volume is considerably easier as most grocery stores carry objects for just this task.  In America however volumes are measured in the system of cups, quarts, pints, gallons, etc.  The conversions for these measurements are readily avalible online and with some searching, even in American supermarkets one can find measuring tools that have some scale in milliliters.  In a pinch one can use a container of known volume, be it a pop can or a milk jug, just for a rough approximation of the volume it contained.