Lab 6: Determination of Empirical Formula
- Page ID
- 514168
\( \newcommand{\vecs}[1]{\overset { \scriptstyle \rightharpoonup} {\mathbf{#1}} } \)
\( \newcommand{\vecd}[1]{\overset{-\!-\!\rightharpoonup}{\vphantom{a}\smash {#1}}} \)
\( \newcommand{\id}{\mathrm{id}}\) \( \newcommand{\Span}{\mathrm{span}}\)
( \newcommand{\kernel}{\mathrm{null}\,}\) \( \newcommand{\range}{\mathrm{range}\,}\)
\( \newcommand{\RealPart}{\mathrm{Re}}\) \( \newcommand{\ImaginaryPart}{\mathrm{Im}}\)
\( \newcommand{\Argument}{\mathrm{Arg}}\) \( \newcommand{\norm}[1]{\| #1 \|}\)
\( \newcommand{\inner}[2]{\langle #1, #2 \rangle}\)
\( \newcommand{\Span}{\mathrm{span}}\)
\( \newcommand{\id}{\mathrm{id}}\)
\( \newcommand{\Span}{\mathrm{span}}\)
\( \newcommand{\kernel}{\mathrm{null}\,}\)
\( \newcommand{\range}{\mathrm{range}\,}\)
\( \newcommand{\RealPart}{\mathrm{Re}}\)
\( \newcommand{\ImaginaryPart}{\mathrm{Im}}\)
\( \newcommand{\Argument}{\mathrm{Arg}}\)
\( \newcommand{\norm}[1]{\| #1 \|}\)
\( \newcommand{\inner}[2]{\langle #1, #2 \rangle}\)
\( \newcommand{\Span}{\mathrm{span}}\) \( \newcommand{\AA}{\unicode[.8,0]{x212B}}\)
\( \newcommand{\vectorA}[1]{\vec{#1}} % arrow\)
\( \newcommand{\vectorAt}[1]{\vec{\text{#1}}} % arrow\)
\( \newcommand{\vectorB}[1]{\overset { \scriptstyle \rightharpoonup} {\mathbf{#1}} } \)
\( \newcommand{\vectorC}[1]{\textbf{#1}} \)
\( \newcommand{\vectorD}[1]{\overrightarrow{#1}} \)
\(\newcommand{\ket}[1]{\left| #1 \right>} \)
\( \newcommand{\bra}[1]{\left< #1 \right|} \)
\( \newcommand{\braket}[2]{\left< #1 \vphantom{#2} \right| \left. #2 \vphantom{#1} \right>} \)
\( \newcommand{\qmvec}[1]{\mathbf{\vec{#1}}} \)
\( \newcommand{\op}[1]{\hat{\mathbf{#1}}}\)
\( \newcommand{\expect}[1]{\langle #1 \rangle}\)
\( \newcommand{\vectorDt}[1]{\overrightarrow{\text{#1}}} \)
\( \newcommand{\vectE}[1]{\overset{-\!-\!\rightharpoonup}{\vphantom{a}\smash{\mathbf {#1}}}} \)
\( \newcommand{\vecs}[1]{\overset { \scriptstyle \rightharpoonup} {\mathbf{#1}} } \)
\( \newcommand{\vecd}[1]{\overset{-\!-\!\rightharpoonup}{\vphantom{a}\smash {#1}}} \)
\(\newcommand{\avec}{\mathbf a}\) \(\newcommand{\bvec}{\mathbf b}\) \(\newcommand{\cvec}{\mathbf c}\) \(\newcommand{\dvec}{\mathbf d}\) \(\newcommand{\dtil}{\widetilde{\mathbf d}}\) \(\newcommand{\evec}{\mathbf e}\) \(\newcommand{\fvec}{\mathbf f}\) \(\newcommand{\nvec}{\mathbf n}\) \(\newcommand{\pvec}{\mathbf p}\) \(\newcommand{\qvec}{\mathbf q}\) \(\newcommand{\svec}{\mathbf s}\) \(\newcommand{\tvec}{\mathbf t}\) \(\newcommand{\uvec}{\mathbf u}\) \(\newcommand{\vvec}{\mathbf v}\) \(\newcommand{\wvec}{\mathbf w}\) \(\newcommand{\xvec}{\mathbf x}\) \(\newcommand{\yvec}{\mathbf y}\) \(\newcommand{\zvec}{\mathbf z}\) \(\newcommand{\rvec}{\mathbf r}\) \(\newcommand{\mvec}{\mathbf m}\) \(\newcommand{\zerovec}{\mathbf 0}\) \(\newcommand{\onevec}{\mathbf 1}\) \(\newcommand{\real}{\mathbb R}\) \(\newcommand{\twovec}[2]{\left[\begin{array}{r}#1 \\ #2 \end{array}\right]}\) \(\newcommand{\ctwovec}[2]{\left[\begin{array}{c}#1 \\ #2 \end{array}\right]}\) \(\newcommand{\threevec}[3]{\left[\begin{array}{r}#1 \\ #2 \\ #3 \end{array}\right]}\) \(\newcommand{\cthreevec}[3]{\left[\begin{array}{c}#1 \\ #2 \\ #3 \end{array}\right]}\) \(\newcommand{\fourvec}[4]{\left[\begin{array}{r}#1 \\ #2 \\ #3 \\ #4 \end{array}\right]}\) \(\newcommand{\cfourvec}[4]{\left[\begin{array}{c}#1 \\ #2 \\ #3 \\ #4 \end{array}\right]}\) \(\newcommand{\fivevec}[5]{\left[\begin{array}{r}#1 \\ #2 \\ #3 \\ #4 \\ #5 \\ \end{array}\right]}\) \(\newcommand{\cfivevec}[5]{\left[\begin{array}{c}#1 \\ #2 \\ #3 \\ #4 \\ #5 \\ \end{array}\right]}\) \(\newcommand{\mattwo}[4]{\left[\begin{array}{rr}#1 \amp #2 \\ #3 \amp #4 \\ \end{array}\right]}\) \(\newcommand{\laspan}[1]{\text{Span}\{#1\}}\) \(\newcommand{\bcal}{\cal B}\) \(\newcommand{\ccal}{\cal C}\) \(\newcommand{\scal}{\cal S}\) \(\newcommand{\wcal}{\cal W}\) \(\newcommand{\ecal}{\cal E}\) \(\newcommand{\coords}[2]{\left\{#1\right\}_{#2}}\) \(\newcommand{\gray}[1]{\color{gray}{#1}}\) \(\newcommand{\lgray}[1]{\color{lightgray}{#1}}\) \(\newcommand{\rank}{\operatorname{rank}}\) \(\newcommand{\row}{\text{Row}}\) \(\newcommand{\col}{\text{Col}}\) \(\renewcommand{\row}{\text{Row}}\) \(\newcommand{\nul}{\text{Nul}}\) \(\newcommand{\var}{\text{Var}}\) \(\newcommand{\corr}{\text{corr}}\) \(\newcommand{\len}[1]{\left|#1\right|}\) \(\newcommand{\bbar}{\overline{\bvec}}\) \(\newcommand{\bhat}{\widehat{\bvec}}\) \(\newcommand{\bperp}{\bvec^\perp}\) \(\newcommand{\xhat}{\widehat{\xvec}}\) \(\newcommand{\vhat}{\widehat{\vvec}}\) \(\newcommand{\uhat}{\widehat{\uvec}}\) \(\newcommand{\what}{\widehat{\wvec}}\) \(\newcommand{\Sighat}{\widehat{\Sigma}}\) \(\newcommand{\lt}{<}\) \(\newcommand{\gt}{>}\) \(\newcommand{\amp}{&}\) \(\definecolor{fillinmathshade}{gray}{0.9}\)The purpose of this experiment is to:
- Determine the empirical formula of a crystalline copper chloride hydrate.
INTRODUCTION
Chemical formulas are a fundamental way to represent the composition of chemical compounds. The empirical formula is the simplest whole-number ratio of atoms in a compound. It is essential to distinguish this from the molecular formula, which specifies the number of atoms of each element in a molecule.
Many ionic compounds incorporate water molecules into their crystal structure, forming hydrates. Hydrates are crystalline compounds that incorporate water molecules into their structure in a specific ratio relative to the anhydrous salt. This water of hydration is an integral part of the crystal structure and can be removed by heating the hydrate. Without the water molecules, the resulting compound is known as an anhydrous compound.
In this experiment, you will investigate copper chloride hydrate, which is known for its green-blue color. The water molecules are removed by carefully heating the hydrate, which results in a color change to brown as the anhydrous copper chloride is formed. The mass difference between the original hydrate and the anhydrous compound allows us to determine the Mass of water lost.
We will use a single-displacement reaction with aluminum as the reducing agent to determine the ratio of copper to chloride in the anhydrous compound. The single displacement reaction involves displacing one element in a compound by another element. In our experiment, aluminum metal will reduce copper ions in solution, forming copper metal and aluminum ions. The copper metal is carefully separated, dried, and weighed. By determining the Mass of water lost and the Mass of copper and chloride in the original compound, we can calculate the moles of each component. By selecting the number of moles of water, copper, and chloride, we can establish this ratio and thus determine the empirical formula of the copper chloride hydrate. The empirical formula represents the simplest whole-number ratio of elements in a compound. From these values, the empirical formula of the copper chloride hydrate can be determined by finding the simplest whole-number ratio of the moles of each element and water, as shown in the example below.
A student heated a sample of a nickel(II) sulfate hydrate and collected the following data:
- Mass of empty crucible and lid: 25.668 g
- Mass of crucible, lid, and nickel sulfate hydrate: 28.225 g
- Mass of crucible, lid, and anhydrous nickel(II) sulfate: 27.389 g
Determine the empirical formula of the nickel(II) sulfate hydrate.
Solution
- Calculate the Mass of the hydrated salt:
o Mass of hydrated salt = Mass of crucible, lid, and hydrate - Mass of empty crucible and lid
o Mass of hydrated salt = 28.225 g - 25.668 g = 2.557 g
- Calculate the Mass of the anhydrous salt:
- Mass of anhydrous salt = Mass of crucible, lid, and anhydrous salt - Mass of empty crucible and lid
- Mass of anhydrous salt = 27.389 g - 25.668 g = 1.721 g
- Calculate the Mass of water lost:
- Mass of water lost = Mass of hydrated salt - Mass of anhydrous salt
- Mass of water lost = 2.557 g - 1.721 g = 0.836 g
- Calculate the moles of anhydrous nickel(II) sulfate (NiSO₄):
- First, find the molar Mass of NiSO₄:
- Ni: 58.69 g/mol
- S: 32.07 g/mol
- O: 16.00 g/mol x 4 = 64.00 g/mol
- Molar mass NiSO₄ = 58.69 g/mol + 32.07 g/mol + 64.00 g/mol = 154.76 g/mol
- Moles of NiSO₄ = Mass of anhydrous salt / Molar mass of NiSO₄
- Moles of NiSO₄ = 1.721 g / 154.76 g/mol = 0.01112 mol
- [In your experiment, you will find the Mass and moles of copper and chlorine from experimental data obtained by a displacement reaction]
- First, find the molar Mass of NiSO₄:
- Calculate the moles of water (H₂O):
- Molar mass of H₂O = (1.008 g/mol x 2) + 16.00 g/mol = 18.016 g/mol
- Moles of H₂O = Mass of water lost / Molar Mass of H₂O
- Moles of H₂O = 0.836 g / 18.016 g/mol = 0.04640 mol
- Determine the mole ratio of anhydrous salt to water:
- Divide both mole values by the smallest number of moles (0.01112 mol in this case):
- NiSO₄: 0.01112 mol / 0.01112 mol = 1
- H₂O: 0.04640 mol / 0.01112 mol = 4.17 ≈ 4
- Divide both mole values by the smallest number of moles (0.01112 mol in this case):
- Write the empirical formula of the hydrate:
- The ratio is approximately 1:4, so the empirical formula is NiSO₄ · 4H₂O
Therefore, the empirical formula of the nickel(II) sulfate hydrate is NiSO₄ · 4H₂O.
1. Always wear safety goggles.
2. Hydrochloric acid is a strong acid that can cause severe burns. Handle it with extreme care and avoid contact with skin and eyes. In case of contact, immediately flush the affected area with copious amounts of water and seek medical attention.
3. The reaction between aluminum wire and hydrochloric acid produces flammable hydrogen gas. Ensure there are no open flames in the vicinity during this step.
4. Handle glassware carefully when it is hot, during heating, and after.
5. Use caution when heating the crucible and handling hot glassware. Allow the crucible and glassware to cool down before handling.
6. Dispose of chemicals properly according to your instructor's guidelines.
7. Clean and return all glassware and equipment to their designated places.
EQUIPMENT AND CHEMICALS NEEDED
| Equipment | Equipment | Chemicals |
|---|---|---|
| 50 mL beaker | Crucible with lid | Deionized water |
| Stirring rod | Clay triangle | Aluminum wire |
| Watch glass | Büchner funnel, Filtration flask, and filtration trap flask | Copper chloride hydrate |
| Iron ring | Ring stand | 95% ethanol |
| Hot plate | Bunsen burner assembly | 6 M HCl |
| Dropper | Filter paper | - |
EXPERIMENTAL PROCEDURE
A: Determination of the Mass of Water by Dehydration of Copper Chloride Hydrate
1) Record the Mass of a clean, dry crucible and its lid in the data table.
2) Add approximately 1.00 g of copper chloride hydrate to the crucible. Weigh the crucible, lid, and hydrate. Record the Mass in the data table.
3) Set up a Bunsen burner with a clay triangle supported by a ring stand and an iron ring as shown in Figure 1. Place the crucible with the hydrate on the clay triangle and gently heat it with a small flame.
Figure 1: Ring stand with clay triangle and crucible with lid
4) It is essential that the hydrate is not overheated, or it will decompose, and black particles will appear within the brown powder. Adjust the Bunsen burner to give a small flame and, holding the base of the burner (not the barrel of the burner), move the burner back and forth
5) Continue heating until the green crystals turn brown, indicating dehydration. Heat for an additional 2 minutes after all crystals have turned brown.
6) Turn the Bunsen burner off and cover the crucible with the lid to prevent the moisture in the air from rehydrating as the compound cools down. Allow it to cool to room temperature for 15 minutes, and then check if any green crystals remain; reheat the crucible if necessary.
7) When no green crystals remain, accurately measure the Mass of the cooled crucible with lid and the anhydrous compound and record the Mass.
B: Determination of the Mass of Copper by Displacement Reaction
8) Transfer the anhydrous copper chloride compound to a 50 mL beaker. Rinse the crucible with two 5 mL portions of deionized water and add these portions to the beaker with the anhydrous compound. Swirl the beaker to completely dissolve the solid, which will change to a green solution.
9) Measure the indicated length of aluminum wire (approximately 0.25 g),(length depends on the gauge/thickness of the wire provided). Coil the wire into a loose spiral and add it to the solution, ensuring it is completely covered.
10) Some hydrogen gas will be given off, and the color of the aluminum will slowly change.
11) Allow the reaction to proceed for about 30 minutes, or until the solution becomes colorless as the copper metal gets deposited on the aluminum wire.
12) Add five drops of 6 M HCl to dissolve insoluble aluminum salts.
13) Use a stirring rod to remove copper from the aluminum wire. Rinse the wire with deionized water, adding the rinsing to the beaker. Remove the aluminum wire and place it on a paper towel to be discarded into the metal waste later.
14) Assemble a Büchner funnel and the filtration flask, attach them to a second filtration trap flask, and then connect the trap flask to the vacuum with a hose.
15) Obtain a dry filter paper to fit the Büchner funnel and record its Mass.
16) Place the filter paper into the funnel. Wet the filter and filter paper in the Büchner funnel with a few milliliters of deionized water.
17) Turn on the vacuum suction and carefully pour the copper solution through the filter. Ensure that all the copper is transferred onto the filter paper. Wash the copper with deionized water.
18) Turn off the suction and add 10 mL of 95% ethanol to the copper on the filter paper. After 1 minute, turn the suction back on for 5 minutes. Turn off the suction and carefully remove the filter paper with the copper.
19) Place the filter paper and copper on a watch glass and dry them on a hot plate at a low setting.
20) Once dry, weigh the filter paper and copper. Record the Mass.
CHEMICAL DISPOSAL
Dispose of the filtrate into a NON-HALOGENATED ORGANIC waste jar and place the aluminum wire and recovered copper powder in a metal solids waste jar.
CALCULATIONS
1. Moles of Water of Hydration
\[\text{mass of water of hydration} = \text{(mass of copper chloride hydrate)(g)} - \text{(mass of anhydrous copper chloride)(g)} \nonumber \]
\[\text{moles of water(mol)} = \frac{\text{mass of water (g)}}{\text{molar mass of water (g/mol)}} \nonumber \]
2. Moles of Copper
\[\text{mass of copper} = \text{(mass of copper and filter paper)(g)} - \text{(mass of filter paper)(g)} \nonumber \]
\[\text{moles of copper(mol)} = \frac{\text{mass of copper (g)}}{\text{molar mass of copper (g/mol)}} \nonumber \]
3. Moles of Chlorine
\[\text{mass of chlorine} = \text{(mass of copper chloride hydrate)(g)} - \text{(mass of water of Hydration)(g)} - \text{(mass of copper)(g)} \nonumber \]
\[\text{moles of chlorine(mol)} = \frac{\text{mass of chlorine (g)}}{\text{molar mass of chlorine (g/mol)}} \nonumber \]
4. Calculate the empirical formula of the copper chloride hydrate.
a. Divide the number of moles each by the smallest number of moles.
b. Determine the smallest whole-number ratio of moles.
PRE-LAB QUESTIONS Name ____________________________________
1. Define the following terms:
o Hydrate:
o Anhydrous compound:
o Empirical formula:
2. Provided one has determined the relative ratio of the number of moles in an unknown compound, one can determine the empirical formula. In an experiment, a student determined the Mass of tungsten (W) to be 4.23 g and found the Mass of an oxide of tungsten, after heating tungsten in air, to be 5.34 g. Determine the empirical formula of this oxide of tungsten. Show your work.
- What is the empirical formula of a compound composed of 60% carbon, 5% hydrogen, and 35% oxygen? Show your work.
4. A compound contains 40.0 % carbon, 6.7 % hydrogen, and 53.3 % oxygen. Its molar Mass is 120.0 g/mol. What is its empirical formula? What is its molecular formula? Show your work.
- Why is it essential to avoid overheating the copper chloride hydrate? What visual cues might indicate overheating?
DATA AND OBSERVATIONS
Name _________________________ Lab Partner(s) ____________________________________
A: Determination of the Mass of Water by Dehydration of Copper Chloride Hydrate
|
Mass (g) |
|
|---|---|
|
1. Empty the crucible with the lid |
|
|
2. Crucible, lid, and copper chloride hydrate |
|
|
3. Copper chloride hydrate (2. – 1.) |
|
|
4. Crucible, lid, and anhydrous copper chloride |
|
|
5. Anhydrous copper chloride |
|
|
6. Water of hydration |
Calculate the Mass of anhydrous copper chloride formed (4. – 1.). (Show your work.):
Calculate the Mass of water lost by the copper chloride hydrate (2. – 4.). (Show your work.):
B: Determination of the Mass of Copper by Displacement Reaction
|
Mass (g) |
|
|---|---|
|
7. Filter Paper |
|
|
8. Filter Paper and Copper |
|
|
9. Dried Copper |
Calculate the Mass of Copper recovered (9 = 8 – 7). (Show your work.):
C: Calculation of the Mass of Chlorine in the Copper Chloride Hydrate
\[\text{mass of chlorine} = \text{(mass of copper chloride hydrate)(g)} - \text{(mass of water of Hydration)(g)} - \text{(mass of copper)(g)} \nonumber \]
|
Mass (g) |
|
|---|---|
| 10. Chlorine |
Calculate the Mass of Chlorine in the original hydrate (8 – 7). (Show your work.):
D: Calculation of the Empirical Formula of Copper Chloride Hydrate
| Moles | Show your work | |
|---|---|---|
| 11. Water | ||
| 12. Copper | ||
| 13. Chlorine |
Determine the simplest whole-number ratio of moles of water, copper, and chlorine. (Show your work.):
The Empirical Formula of Copper Chloride Hydrate:
POST-LAB QUESTIONS
1. Explain how overheating the hydrated compound would affect the calculated empirical formula. Which values would be higher or lower than expected?
- What is the purpose of the aluminum wire in this experiment? Explain the chemical reactions involved.
- If the filtration process was incomplete and the student lost some copper particles, how would this affect the calculated copper-to-chloride ratio in the empirical formula?
- A student obtained the following data:
- Mass of hydrated salt: 5.00 g
- Mass of anhydrous salt: 3.98 g
- Mass of copper recovered: 1.25 g
Calculate the empirical formula of the copper chloride hydrate based on these data. Show your work.