Introduction:

The Mole

     As with any area of technical expertise, chemists have a unit of quantity, known as the mole. It is no different from, say, a unit of length; e.g. a meter. Having distance in meters allows one to convert it to any other unit of choice, by simply looking up the relevant conversion factor. So too with the mole, once it is known for any substance in a given chemical reaction, then the quantity of all other substances can be derived by using the appropriate conversion factor. In our case, these conversions are dependent on the reaction in question and can only be arrived at using a balanced chemical reaction. Coefficients before each substance will be used to make our conversion factors. 

     Why use the mole(?), since the fundamental unit of an element or a compound is an atom or molecule, respectively. Why not use the mass of the individual atom or the sum of the atoms that make up the molecule? Size is the issue here, the mass of an atom is so small (~10^-23 g) that, while it is possible to determine, it is impractical to do so in most laboratory experiments. Somewhat analogous to the fact that when you refer to the distance from Amherst to Northampton you speak in terms of ~10miles rather than ~1,600,000 cms! A more practical mass range from the laboratory perspective is in the gram range, 1-300g. In order to define the mole in terms the fundamental unit (atom/molecule), a standard had to be chosen, and this turned out to be the number of atoms in twelve grams of pure carbon twelve (an isotope of carbon with mass number 12). The number of atoms turns out to be 6.023x10²³ atoms, which has come to be known as Avogadro's Number. Apart from its large size there is nothing special about this number, and if a different standard were chosen it would have a different value. A mole can thus be redefined as the mass of any substance that contains 6.023x10²³ units of that substance. Its use is pretty much confined to atoms and molecules and the unit is hardly likely to replace the dozen when it comes to large objects such as eggs! The mass of any element that corresponds to 6.023x10²³ atoms of that element is equal to its molar mass, which for an element can be found in the periodic table. The masses listed on the table are in fact average values, derived from the various isotopes of the elements and their natural abundance's. For molecules, the molar mass is the sum of the individual atoms that make up the molecule. 

     In summary: 1mole is equal to 6.023x10²³ atoms (or molecules) which in turn is equal to the atomic mass of the element or the molar mass of the molecule in grams per mole. 

     In the laboratory, from a stoichiometry point of view, you can classify the substances that you will be dealing with as either, solids, solutions, or gases. For this experiment we will be dealing exclusively with solids where: 

Determining the Formula of a Compound

    One of the most exciting aspects of chemistry is the discovery of some new substance when one element reacts with another. In order to know one has made a new substance, the formula of the compound formed must be determined quantitatively. This must involve determining the mole ratio and not the gram ratio. The gram ratio gives no meaningful information since the individual masses are not related to equivalent number of atoms. 

    While the experiment that you will be performing in this laboratory is not a new discovery, we are going to treat it as such! You will burn a measured quantity of magnesium in the presence of air (your source of oxygen) and then use mass measurements to determine a possible formula for the resultant product: 

Note: 
When no integer appears before a atom or molecule in the balanced chemical equation, it is inferred that only one molecule is produced, thus the product of the above reaction is written as Mg_xO_y and not 1 Mg_xO_y.

    Remember, that an atom is the smallest particle of an element that still retains the chemical properties of that element and thus the mole ratio should turn out to be a whole number or by simple multiplication can be transformed into a whole number. However the resultant formula is not necessarily the actual formula since all that we have determined is the ratio of one atom to the other, e.g. 

Mg/O = 1:1, could be MgO, Mg₂O₂, Mg₃O₃……etc, i.e. any combination whose atom ratio is 1:1

It is now safe to round 4.98 to 5 and obtain 5:3.  This latter method may look cumbersome, but it works all the time and it does not depend on you recognizing fractions!

Mg/O = 5:3, could be Mg₅O₃, Mg₁₀O₆ ….etc, again any combination whose atom ratio is 5:3.

Recording Data and Significant Figures

Significant Figures:

    Since this is your first experiment, and many of the experiments that you will be performing are quantitative in nature (meaning the result is a determined number), lets delve briefly into significant figures. In most experiments many different kinds of measurements are made using a variety of different equipment, some of which measure to a higher degree of precision than others. It then makes sense that a calculated result based on these measurements can be no more precise than the least precise measurement that went into that calculation. This is the essence of significant figures. 

1. In adding or subtracting measurements, the number of decimal places in the answer should equal the measurement with the fewest decimal places.

 
2. In determining the number of significant figures in a measurement, read from left to right and count all digits, starting with the first digit that is not zero: e.g. 
    

1. 0.049g:      2 Significant figures 

2. 2.0g:          2 Significant figures

Constants are considered to have infinite number of significant figures. Put in another way, these quantities should not be used in determining the measurement with the least number of significant figures.
 

3. In multiplication and division, the number of significant figures in the calculated quantity should be the same as the quantity used with the fewest number of significant figures.

 
Note:
When using a calculator, do the calculation using all the digits allowed by the calculator and round off the final answer.

Rounding: round up if the digit after the last significant figure equals or is greater than 5.

With regards to the laboratory only:

• you may treat as a constant any reagent concentrations given in an experiment.
• if in doubt use one more rather than one less significant figure. 
• we are primarily concerned with significant figures only in the final answers. With all other data we are concerned with recording the data to the accuracy of the instrument used.

          This is one area where we place a lot of emphasis. Always record the data so that it reflects the accuracy of the instrument used. With digital instruments that simply means recording all the digits displayed, even zeros! 

           With other equipment you record to 1/10th or 1/5th of the smallest division. In the case of a buret, the smallest division is 0.1 therefore you should record to two decimal places in increments of 0.01 or 0.02 depending on you ability to guesstimate! With a 400mL beaker, the smallest division is 25mL, therefore you should record this in increments of either 2.5mL or 5mL . 

Experimental Procedure:

Using an Electronic Balance

Please be aware that these balances are somewhat delicate and very sensitive. Do not remove the plastic guard around each. While it may not always be practical to do so, it is good practice to stick with the same balance during any particular experiment. Pictured  are the three models of analytical balances that we use in the General Chemistry program here. During the first lab, your TA will demonstrate the various idiosyncrasies of each! When doing any weighing, please observe the following: 

1. Do not remove the plastic guard around each balance.
2. Do not move the balance, this throws off the calibration and necessitates recalibration of the balance.  

1. Weigh chemicals in containers that are dry on the outside. Never weigh directly on the pan.
2. Weigh everything at room temperature.
3. Zero the balance and keep the draft doors closed during weighing. Avoid leaning on the benches as this in fact interferes with the weighing!

Determining the Formula of a Compound
 

1. Wash a crucible with soap, rinse with tap water and finally distilled water. You will not be able to get the inside immaculately clean but anything that remains after washing thoroughly and heating will not affect the reaction. 

 
2. Support the crucible on a tripod using a clay triangle, as depicted in the diagram 

 
3. Heat the crucible for two to three minutes using a hot (blue) flame. This will burn off any moisture and any impurities in the crucible. (A yellow flame will deposit carbon on the crucible and destroy any hope of getting meaningful results.) 

 
4. When cool, handling the crucible with a paper towel, place it on top of a 50mL beaker and weigh it on the analytical balance (Do not place anything hot directly on the balances). Only weigh the crucible and beaker. 

 
5. Cut a piece of magnesium equivalent to ~0.2-0.25g.  magnesium develops a coating of magnesium oxide over time.  To remove this use the sand paper provided.  Please do not do this on the bench top as sand paper leaves unsightly marks!!.  Once you have sand papered the magnesium, transfer it to the crucible.  Reweigh and record. 

		Heat the entire crucible and its contents with a moderately vigorous blue flame. The magnesium will start to glow, but if it ignites like that depicted on the right, quickly using tongs, place the cover on the crucible (as depicted on the left). Continue to heat for 8-10 minutes, removing and covering the crucible any  time it ignites too vigorously. Allow it to cool to room temperature.
	As we are using air as our source of oxygen (remember air is predominantly nitrogen) we may form some magnesium nitride and we need to convert any of this formed to the oxide. Hence step 7.

7. Carefully add 10 drops of distilled water to the crucible using a dropper. Make sure that all the white powder stays inside the crucible. 

 
8. Heat using a blue flame for five minutes. Allow it to cool. Weigh and record. 

 
9. Continue this heating, cooling, re-weighing, until you get two readings that do not differ by more than 0.001g. 

 
10. Repeat the whole procedure with a second sample. 

 
11. Determine the moles of magnesium and oxygen atoms that reacted and then the mole ratio.