Two of the basic tenets of atomic theory: Lavoisier’s Law of Conservation of Mass and Proust’s Law of Definite Proportions are applied to chemical reactions involving metallic copper and a copper salt. The data obtained from the reactions can be used to determine the empirical formulas of the compounds and the amount of water in a hydrate.
To determine the empirical formulas ofcompounds
To determine the percent composition of a reactionproduct
To apply the Law of Conservation of Mass to a chemical
Toapply the Law of Definite Proportions to determine the empirical formulas of
To determine the empirical formula of a
To learn how to determine the empirical formula of
1. decanting– gradually pouring off the supernatant, leaving the solid in the original container
2. empirical formula – chemical formula based on the lowest possible integer coefficient of elements
3. formula unit - the ratio of atoms in an ionic compound, the chemical formula; although it does not represent the whole crystal structure of an ionic compound, it is useful for stoichiometric calculations
4. hydrate– inorganic compound containing a determined ratio of water molecules per formula unit
5. Law of conservation of mass – establishes that mass is preserved during the course of a chemical reaction; that is, mass is not created nor destroyed in chemicalreactions
6. Law of definite proportions – establishes that chemical compounds always have the same mass ratio ofelementsmolecular formula – chemical formula based on the actual ratio of itselementsoxidation – loss of electrons in an oxidation-reduction (redox)reaction
7. oxidizing agent – element that causes another element to be oxidized by beingreduced(accepting the other element’s electron or electrons)
reduction – gain of electrons in an oxidation-reduction (redox)reaction
8. reducing agent – element that causes another element to be reduced by being oxidized (giving away its electron or electrons to the otherelement)
9. supernatant – liquid lying over the solid in a chemical reaction orprocess
Chemists write formulas in different ways. The most common is the molecular formula. The molecular formula shows the actual number of atoms of each element present in the formula. A simple form of a chemical formula is an empirical formula. An empirical formula consists of the lowest possible integer ratio of the elements forming the compound. The empirical formula is not how a compound exists, and does not represent the actual structure of the compound, but it is useful since it can be determined easily by experiment. For example, the chemical compound benzene has a molecular formula of C6H6. The lowest possible integer ratio between the carbon and the hydrogen will be 1:1, thus benzene has an empirical formula of CH.
Two of the basic foundations of our atomic theory are Antoine Lavoisier’s law of conservation of mass, and Joseph Proust’s law of definite proportions (also known as law of constant composition or law of definite composition). The law of conservation of mass states that matter is not created nor destroyed, but preserved during the course of a chemical reaction. The law of definite proportions states that a chemical compound will always have the same ratios of its elements. We can observe that benzene will always consist of 92.2 % carbon by mass, and 7.8 % hydrogen by mass.
Each C6H6 molecule will contain six carbon atoms, and six hydrogen atoms. Using the molar masses from the periodic table, we obtain:
No matter how large or small the sample is, benzene will always have this ratio or percent. The law of definite proportions and the law of conservation of mass enable us to determine the formula of a compound if we are able to determine the masses, or the percentages of the different elements in the formula. Once we obtain the masses for the elements, we can use the elements’ molar masses to determine the number of moles of each element and determine the ratio between them.
For example, a chemist found that a given compound is composed of 0.410 g Al and 1.590 g Cl. What is the empirical formula?
- We should expect that due to experimental errors, the ratios will not be exact integer numbers, but within the uncertainty, they can be rounded to the closest integers. If the ratio of any of the elements is close to 0.5 (1.5, for example), the ratios of all the elements in the compound need to be multiplied by a factor of 2 to turn them into integer ratios.
Let’s consider another example. In this experiment, in addition to the law of definite proportion, we will also need to apply the law of conservation of mass to find the chemical formula of a hydrate. Hydrates are salts with a strong affinity towards water. They have a fixed ratio of water molecules.
1. Familiarize yourself with the use of the Bunsen burner. Following instructions of Technique 9 Using the Bunsen Burner, turn on and adjust height of flame and air so all three cones visible.
2. Obtaina clean Check the crucible for cracks. A cracked or chipped crucible will break upon heating or cooling. Support the crucible on a clay triangle and heat with an intense flame
for 5 minutes (see Figure 1). This “firing” process will remove any contaminants from the
crucible that would affect mass change during the experiment.
Allow them to cool to room
1. Measure the mass of the fired, cool crucible. Use only clean, dry crucible tongs to handle the crucible and lid for the remainder of the
Add between 1 and 1.5 g of
2. Heat over the flame for complete reaction (turn into the black copperoxide).
3. Allow the crucible to cool. Weigh the burned product and
4. Determine the empirical formula of the copper
Part B Determination of the Amount of Water in a CuSO4×nH2O Compound
Allow time for glassware and equipment tocool
Technique 2: Using abalance
copper (II) sulfatecrystals
1. Weigh a clean and fired crucible. Add approximately 0.2 g of CuSO4nH2O to the crucible andrecord the mass to the full precision of the balance.
2. Heat for at least 10 minutes until the color changed to Let it cool for five
3. After five minutes, weigh the crucible with the anhydrous
4. Determine the mass of the dehydrated
5. Determine the amount of water in the
1. Dispose of all materials on the corresponding waste containers as indicated by your instructor.
2. Wash all glassware with soap and water, and then rinse with deionized water (Technique 1 Cleaning Glassware). Dry the outsides of the Return all glassware to its place.
3. When using a crucible and lid, why must the crucible and lid be inspected for cracks before using?
4. What is meant by “firing the crucible”? Why must we “fire” the crucible before usingit?
5. A30 g sample of titanium chemically combines with chlorine gas to form 5.16 g of titanium chloride. (a) What is the empirical formula of titanium chloride? (b) What is the percent by mass of titanium and the percent by mass of chloride in the sample?
6. A 0.500 g sample of tin foil reacted with oxygen to give 0.635 g of product.
(a) What is the empirical formula of the tin oxide?
(b) What is the percent by mass of tin and the percent by mass of oxygen in thesample?
7. Epsom salt is commonly purchased in the pharmacy for a variety of uses, anti-inflammatory, laxative,and Epsom salt is a hydrated salt of magnesium sulfate. If 2.000 g of Epsom salt are heated to remove the waters of hydration, 0.977 g of the anhydrous (without water) magnesium sulfate, MgSO4, were obtained. Calculate the number of waters of hydration, and write the chemical formula of the Epsom salt.
Include the data collected in the lab. Determine the corresponding formulas. Show the calculations for Trial 1 (Your instructor may assign more than onetrial).
Answer the following Post-Lab
1. In part A, the crucible was not “fired” before burning the copper. When the copper was burned,volatile impurities in the crucible were burned Will this error increase, decrease, or not affect the ratio of copper to oxygen determined and the empirical formula found?
2. In part A, a student did not completely burn all the copper and some un-burned copper remained. How will this affect the empirical formulafound?
3. In part B, a student burned the hydrate over a very intense heat, and some of the copper(II) sulfatewas turned into a copper(II) Will the number of waters of hydration determined be too high, too low, or remain the same?