4 Exp. 4: Formation of an Oxide

Determination of the Empirical Formula of Two Compounds:

Pre-Lab, may be included on BB pre-quiz:

  1. Define empirical formula

 

  1. Define the Law of Definite Composition

 

 

 

  1. Write the empirical formula for each of the following compounds:
    1. 0.200 mol Al and 0.600 mol Cl

 

 

  1. 10.98 g Ba, 2.57 g S, and 5.12 g O

 

  1. When 2.50 g of copper (Cu) react with oxygen, the copper oxide product has a mass of 2.81 g. Copper oxide has two possible formulas; CuO or Cu2O. Which formula is supported by the given data? Be sure to show a calculation.

 

  1. Calculate the theoretical percent composition, by mass, for each element in the compound MgO.

 

  1. A compound contains only C, H and O. If the combustion of a 0.255 g sample of the compound produces 0.561 g of CO2 and 0.306 g of H2O, what is the empirical formula of the compound?

 

 

  1. Explain the safety concerns in this lab and how to avoid getting injured.

 

 

 

 

 

 

 

 

Introduction

Joseph Proust did much of his pioneering work using inorganic binary compounds including sulfates, sulfides, and metallic oxides. Today you will be working with both sulfates and metallic oxides, and much like Proust, you will use mass relationships to come up with the chemical compound formulas associated with the masses of the elements present. Proust published the law of Constant Composition in 1794 in a paper detailing his work on iron oxides; “A chemical compound always contains the same elements combined together in the same proportion by mass.” In order to frame this law, lets first review chemical and physical change.

When a substance undergoes a physical change, bonds between atoms are not re-arranged.  When iron is melted, it undergoes a physical change from solid to liquid, but the identity of the metal remains the same.  In a chemical change, however, the substance gains a different composition and different properties.  As iron oxidizes, for example, it changes from a dull gray to iron (III) oxide to a ‘rust’ colored substance with different properties and a different composition.

You will be observing a chemical reaction today that falls under the category of a combustion, or combination with oxygen. We are most familiar with the combustion process that occurs with a hydrocarbon such as wood, oil, wax, or gasoline. Such combustions form the most stable oxygen compounds possible, CO2(g) and H2O. The combustion of a metal, such as magnesium, will also form the most stable oxide of the metal with the formula MgxOy. With careful mass measurement and stoichiometric analysis, you will be able to determine the formula of the metal oxide formed with magnesium, or the ratio of oxygen to metal in the product, much as Proust did centuries ago. It is important to note that the ratio is an atom to atom or mole to mole ratio, so the mass values you obtain will all be converted into moles before analysis.

Table 1: Comparison of Physical and Chemical Changes

Physical Chemical
Change in State Formation of a gas (bubbles)
Change in Size Formation of a solid (precipitate)
Tearing Disappearance of a solid (dissolving)
Breaking Change in color*
Grinding Heat given off or absorbed

*Change in color can also be physical.  A hot stove is one example.

Most chemical reactions can be classified into five types, shown in table 2.

 

Table 2: Types of Chemical Reactions

Type of Reaction Description Example Equation
Combination/Synthesis Elements or compounds form a more complex product. P + O2  à P2O5
Decomposition A reacting substance is split into simpler products. KClO3 à  KCl + O2
Single replacement One element replaces another in a compound NaCl + F2 à NaF + Cl2
Double replacement Elements in two compounds switch places AgI + Fe2(CO3)2 à FeI3 + Ag2CO3
Combustion Reactant and oxygen form an oxide product. Fe + O2 à Fe2O3

 

Objective

The objective of this lab is to discover the empirical formula of two compounds, the first by combustion, the second by dehydration of a hydrate. The law of definite proportions will be used.

Safety/Housekeeping

  • When using flame, tie back long hair, remove loose clothing, especially synthetics, wear your lab coat and goggles but not gloves while burning
  • Don’t look directly at the burning magnesium.
  • Remember that hot objects take a while to cool
  • Discard your solid compounds in the inorganic solid waste and aqueous compounds in the labeled waste container in hood 1.

 

Procedures

Finding the Simplest Formula of an oxide

Materials: Crucible, crucible lid, crucible tongs, clay triangle, iron ring, ring stand, Bunsen burner, magnesium ribbon, steel wool, ceramic square.

 

*Use a procedure that makes sense to you. The crucible lid, for example, can fall and break. If it makes sense to weigh your compound without the lid, do so and modify the data table.

  1. Obtain a clean and dry crucible and lid. Secure stains in the crucible will not affect the mass calculations. Place the crucible and cover in the clay triangle and heat for one (1) minute.  Cool until close to room temperature.  Using crucible tongs, weigh the crucible and cover (together) and record.  Do not round any digits from masses that you collect.
  2. Obtain a piece of magnesium ribbon that has a mass of 0.15 to 0.3 g. Remove any tarnish found on the ribbon by polishing with steel wool or sandpaper and weigh once more. Wind the ribbon into a loose coil and place the coil into the crucible. Weigh the crucible and cover with the magnesium.
  3. Begin to heat the crucible with the lid. Oxygen must, however, be replaced by lifting the lid periodically. Watch closely for smoke or fumes, which is the first indication of the magnesium and oxygen reacting. As the magnesium reacts, leave the lid off for a bit to make sure the magnesium finishes reacting. Avoid looking directly at the burning magnesium as it emits enough light to damage your eyes.

Observe the reaction of the magnesium and oxygen about every minute.  When the magnesium no longer produces smoke or a flame, heat the crucible strongly for another five (5) minutes.  Turn off the burner and let the crucible and its contents cool to room temperature.  It may cool more quickly on the ceramic square than on the warm triangle.

*Do not place your crucible on a cold surface such as your ring stand. The thermal shock will crack your crucible. Do not add water to your warm crucible. It will crack.

  1. Reweigh the crucible, cover, and oxide contents. Record their mass and describe the appearance of the oxide product.
  2. Perform the combustion twice. The ratio of Mg to O should be very close. If it is not, perform a third combustion.

You may find the following example data table useful as you work through your calculations; you will need to reproduce it or one like it for your informal report:

Data Table 1: Magnesium Oxide Formula
  Trial 1 Trial 2
Mass of Crucible and Lid (g)
Mass Magnesium (g)
Mass Magnesium Oxide Product (g)
Mass Oxygen (g)
Moles Magnesium (mol)
Moles Oxygen (mol)*
Average Moles Mg (mol)
Average Moles Oxygen (mol)

*Note: Use O rather than O2 to perform your calculations as your calculated oxygen is combined with Mg.

Calculations:

Atoms in a chemical formula combine in whole number ratios. If this is true, this must remain true for moles of atoms in a formula. Divide the numbers of moles magnesium and oxygen by the smaller of the two values. This dictates that the quotient of the smaller value is one (1). The other quotient must be an integer, 1, 2, 3, etc., or easily lend itself to a whole number ratio. For example, 1:1.5, when doubled, becomes a whole number ratio of 2:3. 1:1.33, when tripled, becomes a whole number ratio of 3:4 etc. Using the mole values for magnesium and oxygen, use this method to find the simplest (empirical) formula for the oxide of magnesium. With careful measuring and recording, your data will lend itself to a whole number ratio. You must report your data as a whole number and may not use decimals or fractions. Points will be awarded for obtaining the correct formula. Many compounds have small whole number ratios. If you find a large ratio, 8 to 9 for example, you may consider checking your measurements and simplifying to a smaller ratio.

Data:

Report the ratio of Mg to O as a chemical formula with whole numbers:

_________________________________________________________

 

Part 2: Formation of a Hydrate

Discussion:

A hydrate is a salt that contains a specific number of water molecules called the water of hydration.  The number of water molecules is fixed for each kind of hydrate but differs from one salt to another.  In the formula of a hydrate, the number of water molecules is written after the salt formula and is separated by a large, raised dot.

CaSO4 ● H2O                         MgSO4 ● 7H2O                                     Na2CO3 ● 10H2O

Heating the hydrate provides energy to remove the water molecules.  The salt without water is called an anhydrate. The following is an example but is not the dehydration that you will perform.

MgSO4 ● 7H2O  à    MgSO4             +          7 H2O (g)

Hydrate                       Anhydrate                   Water of hydration

Through careful measurement, you will determine the water of hydration associated with the salt CuSO4 by separating the anhydrate from the water.

Procedure Part two:

Materials: Crucible, clay triangle, crucible tongs, hydrate of CuSO4, iron ring and stand, Bunsen burner, ceramic square.

  1. Scrape out your crucible into the inorganic solid waste and heat for 2-3 minutes. Let it cool and weigh carefully.
  2. Fill the crucible about 1/3 full of a hydrate of CuSO4. Record the mass of the hydrate and crucible.
  3. Record the appearance of the hydrate.
  4. Set the crucible and hydrate on the clay triangle and iron ring. Heat gently for five (5) minutes.  You may need to gently stir the compound using a metal spatula. Increase the intensity of the flame, and heat strongly for ten (10) minutes.  The bottom of the crucible may glow a dull red. Remove the crucible to the ceramic square to cool.  Record the mass.
  5. Heat the crucible and contents a second time for five (5) minutes. Cool and reweigh.  If the mass from the second heating is within 0.05 g of the mass obtained after the first heating, you have completely dehydrated the salt.  If not, heat again until you have agreement between the final masses.  Use the final mass for your calculation.  Do not average the masses.
  6. Record the appearance of the hydrate after heating.
  7. Place the compound on a watch glass and slowly add water dropwise to the compound using a disposable pipet. Record your observations.

The following example data table may be helpful in recording your data.

Data Table 2: Formula of a Hydrate of CuSO4
First Heating Second Heating
Mass of dry crucible (g) xxxxxxxxxxxxxx
Mass of crucible and hydrate (g) xxxxxxxxxxxxxx
Mass of crucible and anhydrate (g)
Mass of hydrate (g)
Mass of anhydrate (g)
Mass of water (g)
Moles anhydrate
Moles water

 

Data:

Calculate the formula of the hydrate: This calculation is similar to an empirical formula calculation. You must, however, find a ratio of one (1) for the anhydrate and a whole number for the water. In other words, consider dividing by the smallest number of moles.

 

 

Post Lab Questions:

  1. Through experimentation you found a 9.0 gram of water sample from Nampa contained 1.0 g hydrogen and 8.0 g of oxygen. Another water sample, 13.5 g total mass, taken from Crouch contained 1.5 g hydrogen and 12.0 g oxygen. Show how this illustrates the Law of Definite Composition.

 

  1. Two different chemicals, A and X, were reacted. The first time they were combined, the ratio was AX1.51. The second time they were reacted, the ratio was AX3.32. Give at least two possible reasons for the difference. (The Law of Definite Composition is valid in both cases). Consider what ASSUMPTIONS you are making when doing an experiment and whether these assumptions are valid.

 

 

  1. An old car rusts in place. If none of the rust is removed by weathering, will the car gain or lose mass?  Defend your answer.

 

  1. What practical applications would knowing the percent composition of an element in a compound have in chemistry or another discipline?

 

  1. If the law of definite composition is obeyed every time, how is it possible to obtain two different formulas for copper oxide?

 

Your Report:

  1. Complete the pre-lab and take the pre-lab quiz in your Blackboard shell before attending lab.
  2. Use a lab notebook to record observations and make calculations.
  3. Create an informal lab document to turn in at the beginning of the next lab. Your lab report should include clearly labeled and tabulated or graphed raw data, any formulas or equations, unknown number or letter (if applicable), conclusive data such as chemical or hydrate formula etc. with clear example calculations or sample calculations*. Discuss results relevant to your findings. Think about the main point of the lab, or the results that you worked for and be sure to include it or them. Refer to How to Write an Informal Lab Report as you write your report.
  4. Include numbered responses to any post lab questions. You do not need to re-write the questions but will use complete sentences or a short paragraph as appropriate. Be specific. If you describe an error, for example, you must describe the direction the data would skew and why.
  5. Submit your document before your next lab appointment under the assignment tab on your laboratory Blackboard shell.

 

 

References:

Davila, R. In Lab 4: Law of Definite Composition College of Southern Idaho; College of Southern Idaho; 2020.

Proust’s Law of Constant Proportion. https://chem.libretexts.org/Bookshelves/Physical_and_Theoretical_Chemistry_Textbook_Maps/Supplemental_Modules_(Physical_and_Theoretical_Chemistry)/Atomic_Theory/Dalton’s_Atomic_Theory/Proust’s_Law_of_Constant_Proportion (accessed Jun 22, 2020).

Adapted from: Timberlake, K. In Chemistry laboratory manual: an introduction to general, organic, and biological chemistry, 9th ed.; Pearson/Benjamin Cummings: San Francisco, CA, 2006; pp 119–128.