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Is the percentage of iodine the same in all CuI mineral supplements?

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    Our earlier discussion about the essential mineral iodine indicated that mass is a very important characteristic of atoms—it does not change as chemical reactions occur. It is a much more reliable measure of how many nutrient atoms are present than Volume, for example, because atoms or molecules pack together more tightly in liquids and solids or become more widely separated in gases when a reaction takes place. From the time Dalton’s theory was first proposed, chemists realized the importance of the masses of atoms, and they spent much time and effort on experiments to determine how much heavier one kind of atom is than another.

    Earlier when we were looking at sources of iodine, we noted the use of CuI (copper (I) iodide, or cuprous iodide) as a salt additive to supply "iodine". To know if CuI provides the RDA of iodine, we must be assured that its composition is constant. That would depend only on the weights of the Cu and I atoms. It was asserted that an atom of iodine is 2.00 times as heavy as an atom of copper in the salt additive CuI. Where did that value come from, and how is it related to the composition of CuI?

    Early chemists determined that a 20 g sample of a compound of copper and iodine contains 6.67 g of copper and 13.33 g of iodine. When there were no simple ways to determine the microscopic nature of a compound or its formula, they made the simplest possible assumption, that the formula contained the minimum number of atoms: one of copper and one of iodine. We now know that the formula CuI is correct (luckily, CuI2 doesn't exist). However, erroneous assumptions about the formulas for other compounds led to half a century of confusion about atomic weights.

    Since the formula is CuI, chemists argued that the ratio of the mass of iodine to the mass of copper in the compound must be the same as the ratio of the mass of 1 iodine atom to the mass of 1 copper atom: 

    \[\dfrac{\text{Mass of 1 I atom}}{\text{Mass of 1 Cu atom}} = \dfrac{\text{Mass of I in CuI}}{\text{Mass of Cu in CuI}} = \dfrac{13.33g}{6.67g} = 2.00\]

    In other words the mass of an iodine atom is about two times as great as the mass of a carbon atom. Notice that this method involves a ratio of masses and that the units (grams) cancel, yielding a pure number. That number (2.00) is the relative mass of an iodine atom compared with a carbon atom. It tells nothing about the actual mass of an iodine atom or of a carbon atom–only that carbon is three-quarters as heavy as oxygen. The relative masses of the atoms are usually referred to as atomic weights. Their values were are in a Table of Atomic Weights, along with the names and symbols for the elements. The atomic-weight scale was originally based on a relative mass of 1.00 for the lightest atom, hydrogen. As more accurate methods for determining atomic weight were devised, it proved convenient to shift to oxygen and then carbon, but the scale was adjusted so that hydrogen’s relative mass remained close to 1. Thus nitrogen’s atomic weight of 14.0067 tells us that a nitrogen atom has about 14 times the mass of a hydrogen atom. The fact that atomic weights are ratios of masses and have no units does not detract at all from their usefulness. It is very easy to determine how much heavier one kind of atom is than another.

    Example \(\PageIndex{1}\): Iodine vs Copper

    Use the Table of Atomic Weights to show that the mass of an iodine atom is about 2.00 times the mass of a copper atom. (This tells us--as we saw above--that CuI is 66.6% iodine by mass.

    Solution The actual masses of the atoms will be in the same proportion as their relative masses. The atomic weight of iodine is 126.90 and that of copper is 63.546. Therefore

    \[\dfrac{\text{Mass of 1 I atom}}{\text{Mass of 1 Cu atom}} = \dfrac{\text{relative mass of an I atom}}{\text{relative mass of a Cu atom}} = \dfrac{126.90g}{63.546g} = 1.997\nonumber\]

    or Mass of an I atom = 1.997 × mass of a Cu atom

    The atomic-weight table also permits us to obtain the relative masses of molecules. These are called molecular weights and are calculated by summing the atomic weights of all atoms in the molecule.

    Example \(\PageIndex{2}\): Cupric Iodide

    How heavy would a cupric iodide molecule (CuI2, copper (II) iodide) be in comparison to a single iodine atom? (As we saw before, although it looks like a good source of iodine, cupric iodide is unstable, and decomposes to CuI. Neither compound has individual molecules, but extended -Cu-I-Cu- linkages).

    Solution First, obtain the relative mass of an CuI2 molecule (the molecular weight):

    1 Cu atom: relative mass = 1 × 63.546 = 63.546
    2 I atoms: relative mass = 2 × 126.9 = 253.9
    1 Hg2Br2 molecule: relative mass = 317.4


    \[\dfrac{\text{Mass of a CuI}_2\text{molecule}}{\text{Mass of an I atom}} = \dfrac{317.4}{126.9} = \text{2.501}\nonumber\]

    The CuI2 molecule is about 2.5 times heavier than a iodine atom.

    From ChemPRIME: 2.5: Atomic Weights

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