- 1 52 53 64 65 66 67 Pre Lab Assignment Table 1 Percent Water In The Hydrates Formula 1 Mgso4 7h20 2 Mgso4 5h20 3 Mgso4 1 (30.85 KiB) Viewed 77 times
- 1 52 53 64 65 66 67 Pre Lab Assignment Table 1 Percent Water In The Hydrates Formula 1 Mgso4 7h20 2 Mgso4 5h20 3 Mgso4 2 (50.61 KiB) Viewed 77 times
- 1 52 53 64 65 66 67 Pre Lab Assignment Table 1 Percent Water In The Hydrates Formula 1 Mgso4 7h20 2 Mgso4 5h20 3 Mgso4 3 (48.23 KiB) Viewed 77 times
- 1 52 53 64 65 66 67 Pre Lab Assignment Table 1 Percent Water In The Hydrates Formula 1 Mgso4 7h20 2 Mgso4 5h20 3 Mgso4 4 (41.55 KiB) Viewed 77 times
Hydrates As we learned in the "Chemical Change and Nomenclature" lab, a hydrate is an ionic compound that has a definite number of water molecules chemically, mostly by the coordinate covalent bonds and hydrogen bonds connected with each cation and therefore enclosed within its crystal lattice. Those water molecules are called water of hydration. The general formula of a hydrated salt is Cat.An, nH₂O (Cat=cation, An-anion, n-number of moles of water per formula unit of salt). MgSO4 7H₂O → MgSO4 x H₂O + (7-x) H₂O (Eq.1) .. H waters of hydration- MgSO, 6H₂O, (hexahydrate) MgSO SH₂O, (pentahydrate) MgSO 4H₂O, (tetrahydrate) MgSO, 3H₂O, (trihydrate) Cu²+ When the metal hydrate salt is heated, the supplied heat breaks the bonds between water molecules and cations, and water evaporates from the crystal lattice. When heated, the salt hydrate usually converts to a salt hydrate with fewer moles of water (Eq 1.), than the initial amount, and then to its anhydrous form (Eq.2), which means no water. Most salts are thermally unstable, and dehydration can be accompanied by decomposition at higher temperatures (Eq.3). These changes are illustrated below using magnesium sulfate as an example. At 130-160°C magnesium sulfate heptahydrate decomposes and the following hydrates form: MgSO 2.5 H₂O, (no prefix) MgSO, 2H₂O, (dihydrate) MgSO, 1.25H,O, (no prefix) MgSO, H₂O, (monohydrate) H H₂O Picture 1. The simplified structure of the formula unit of copper (II) sulfate hydrate. Only there out of five molecules of water bonded to the copper (II) ion are shown.
Formula of Hydrate At higher temperatures, a MgSO4 0.5H₂O (hemihydrate)forms. Anhydrous amorphous salt forms around 300-350°C according to the final equation: MgSO4.7H₂O(s)→ . MgSO4(s) + 7H₂O(g) (Eq.2) Finally, at temperatures above 950 °C the decomposition reaction takes place: MgSO4(s)→ MgO(s) + SO2(g) +O2(g) (Eq.2) Many of the salts contain a transition metal such as copper or nickel and therefore are colorful. The color of these compounds is determined by the number of water molecules attracted to cation, so the anhydrous compound will have not only a different structure but also a different color than the hydrate: CuSO4.5H₂O(s) (blue) CuSO4(s) (white) + 5H₂O(g) This process can be reversed by adding liquid water to the anhydrous CuSO,: CuSO4(s) (white) + 5H₂O(I) CuSO4.5H₂O(s) (blue) The percent of water in a hydrate can be easily determined experimentally by measuring the mass of hydrated salt, dehydrating it by heating it to the constant mass, and measuring the mass of anhydrous salt. The mass of the water is calculated by subtracting the mass of the anhydrous compound from the mass of the hydrate. The formula of the hydrate is calculated by finding the mole ratio of anhydrous salt to the water. Empirical and Molecular Formulas of the Compounds Empirical formula gives the relative number of atoms of each element in a compound. Molecular formula gives the actual number of atoms of each element in a molecule of the compound. The numbers in a molecular formula will be whole number multiples of the numbers in an empirical formula. To determine the molecular formula of a compound, we need to know both the empirical formula and the molar mass of the compound.
Using Spectrophotometry (Colorimetry) to Determine the Formula of Hydrate EXPERIMENT In the first part of the experiment, you will be given a magnesium sulfate hydrate. You will be asked to determine the identity of the unknown using two approaches: comparing the % H2O in your compound to those on a list of possible compounds below and calculating the empirical formula of your hydrate. The unknown sample that you will be given is one of the six following hydrates: 1. 3. MgSO4.7H₂O MgSO4.5H₂O 5. 6. MgSO4 • 2H₂O MgSO4 H₂O 2. 4. You will measure the mass of hydrated salt, then remove the water from the crystals and measure the mass of the anhydrous salt. Complete dehydration cannot be detected by visual inspection. Thus, you will reheat the sample until the values of successive weighing of the residue will indicate that all the water of hydration has been driven off and only the anhydrous compound remains in the crucible. This is called heating to a constant weight. The data gathered will allow you to determine the percent water in the hydrated salt and the empirical formula (the simplest whole-number ratio of atoms in a compound) for the hydrated salt. The mass of the water is calculated by subtracting the mass of the anhydrous compound from the mass of the hydrate. The percent of water is calculated by dividing the mass of the water by the mass of the hydrate, then multiplying by 100%. MgSO4 • 4H₂O MgSO, 3H₂O ●