To analyze your findings and evaluate your methods.
To reflect on how your experiment could be improved in the future.
x1 Test Tube Rack
x5 Empty Test Tubes
x1 Bunsen Burner
x1 Test Tube Holder
x1 Eye Protection
x1 Test Tube with unknown Solution
x1 Beaker with HCL
If I am testing for halide ions using 2cm3 ±0.05 of AgNO3, then a colorful precipitate is going to form but no bubbles will be present.
If I am testing for sulphate ions (SO42) using 2cm3 ±0.05 of 0.1 mol/dm3 HCL and BaCl2, then a precipitate is going to form but no bubbles will be present.
If I am testing for carbonates (CO32-), using 2cm3 ±0.05 of 0.1 mol/dm3 HCL, then many bubbles will form but no precipitate.
If I am testing for metal ions (Cu2+, Fe2+, Fe3+) using 3cm3 ±0.05 of 0.1 mol/dm3 NaOH, then a dark colored precipitate will form.
If I am testing for Ammonium Ions (NH4+) using 2cm3±0.05 0.1 mol/dm3 NaOH and heat, then the litmus paper will turn the color green.
If I am testing for Na+, K+, Ca2+, Cu2+ using a nichrome wire with HCL over a Bunsen burner, then colored flames will be produced.
Part 2: If I am given test tube number nine with an unknown compound, then the compound will be formed by Na+ and SO42- (Sodium Sulfate).
Part 1- Testing to identify non-metal ions
- Test for Halide Ions Cl-, Br-, I-
- Find the Cl- Solution
- Using a dropper add 2cm3 ±0.05 Silver Nitrate AgNO3
Ag+(aq) + Cl−(aq) --> AgCl(s)
Procedure (continued from Part 1):
AgCl is insoluble in water and so it forms a precipitate
- Record the colour of the precipitate on the table
- Repeat with Br- and I- ions
- Test for Sulphate Ions SO42-
- Find the SO42- solution
- Using a dropper add 2cm3 ±0.05 of 0.1 mol/dm3 HCL
- Using a dropper add 2cm3 ±0.05 of Barium Chloride
Ionic equation is:
Ba2+(aq) + SO42−(aq) BaSO4(s)
- Record the color of the precipitate on the table
- Test for Carbonates CO32-
- Find the test tube with CO32-
- Add 2cm3 ±0.05 of 0.1 mol/dm3 HCL
- Observe the solution, if bubbles made the test is positive.
Equation of the reaction: 2 HCl(aq) + Na2CO3(s) --> 2 NaCl(aq) + H2O(l) + CO2(g)
Ionic equation: 2 H+(aq) + CO32−(aq) --> H2O(l) + CO2(g)
Part 2- Tests to Identify Metal Ions:
- Test for Cu2+, Fe2+, Fe3+
- Find the Cu2+ solution
- Using a dropper add 3cm3 ±0.05 of 0.1 mol/dm3 NaOH
- Record in the table the colour of the precipitate
The ionic equation for copper(II) ions reacting with an alkali is: Cu2+(aq) + 2 OH−(aq) --> Cu(OH)2(s)
4)Repeat using the Fe2+ and Fe3+ solutions
- Test for Ammonium Ions NH4+
- Find the NH4+ solution
- Using a dropper add 2cm3±0.05 0.1 mol/dm3 NaOH
- Heat gently over the Bunsen burner
- Test vapors with indicator paper (litmus paper)- Remember to use eye protection!
- Record the color change of litmus paper on a table.
3) Flame Tests for Na+, K+, Ca2+, Cu2+ (these are in communal beakers)
IMPORTANT ONLY USE A LITTLE SOLID WE DO NOT WANT THE AIR FULL OF TOXIC IONS
- Dip the nichrome wire in a beaker containing HCL
- Find the beaker with Na+
- Dip the spatula in Na+, so little adheres (Not too much)
- Hold the spatula in Bunsen flame- Important use eye protection!
- Observe the colours of the flame and record on the table.
- Repeat for K+, Ca2+, Cu2+
Procedure - Part 2: Testing for an unknown Ionically Bonded Substance:
- Choose individually a test tube with an unknown ionically bonded substance.
- Get 5 empty test tubes and divide equally the unknown ionically bonded substance in each of the 5 test tubes.
- Perform a series of tests to identify the non-metal component:
- Using a dropper add 2cm3 ±0.05 Silver Nitrate AgNO3 to the test tube to test for Cl-, Br-, I-. Is the result positive or negative?
- In another test tube, add 2cm3 ±0.05 of 0.1 mol/dm3 of HCL and 2cm3 ±0.05 of Barium Chloride to test for SO42. Is the result positive or negative?
- In another test tube, add 2cm3 ±0.05 0.1 mol/ of HCL to test for CO32-. Is the result positive or negative?
- Perform a series of tests to identify the metal component:
- In another test tube, add 3cm3 ±0.05 of 0.1 mol/dm3 NaOH to test for Cu2+, Fe2+, Fe3+. Is the result positive or negative?
- Using the test tube from Part 4 Step 1 containing 3cm3 ±0.05 of 0.1 mol/dm3 NaOH, heat the test tube over a Bunsen burner and hold litmus paper on the top. If the litmus paper turns green, then the test for NH4+ is positive. If not, then it is negative.
- Dip a microspatula in a beaker containing 0.1 mol/dm3 of HCL and then add the unknown ionic compound from the fifth test tube on the microspatula.
- Place the microspatula over the Bunsen burner to test for Na+, K+, Ca2+, Cu2+ and observe a color change in the flame. Is the result positive or negative?
- See on your data table which metal and non-metal resulted positive and identify your unknown ionic compound by writing a chemical formula.
- Once you have identified your substance you can clean up and leave. Leave the solutions in the test tubes for disposal.
Tables 2- Testing for my compound- TEST TUBE NUMBER 9
My first data table shows the results of part 1 of the experiment of tests for non metal ions. The first three ions tested were Cl-, Br-, I- where I tested for halide ions using silver nitrate. As we can see on my data table, a white precipitate indicates a chloride, a light green precipitate indicates a bromide, and a light/pale yellow color indicates an iodide. In testing for all halide ions, no bubbles were formed. The next test was for sulphate ions (SO42- ) using 0.1 mol/dm3 of HCL and barium chloride. As we can see on the data table, a white dense precipitate indicates a sulphate. Again, no bubbles were formed. The last test for non metal ions was for carbonates (CO32-) by adding 0.1 mol/dm3 of HCL. For this test, many white bubbles formed meaning that when we add dilute hydrochloric acid to a carbonate, carbon dioxide gas is released. In fact, we can see this in the equation:
2 H+(aq) + CO32−(aq) H2O(l) + CO2(g)
My second data table shows the results for part 1 of the experiment regarding metal ions. The first three positively charged ions (cations) were Cu2+, Fe2+, Fe3+ which all formed a dark colored precipitate with dilute sodium hydroxide. Copper (II) ion formed a dark precipitate on the top, the Fe(II) formed a bronze precipitate with black pieces on top and the Fe(III) formed a darker orange precipitate than Fe(II). We can say that the higher the positive charge, the darker the precipitate is. The next test was for ammonium ions NH4+ where sodium hydroxide and heat were applied and ammonia gas was released that turned the litmus paper green. This meant that the litmus paper had been exposed to an alkaline source and so there was a basic solution because the pH greater than 7 (pH of 9). The last test was the flame test for Na+, K+, Ca2+, Cu2+ where dilute hydrochloric acid and heat were applied. The sodium ion had an orange flame, the potassium had a lilac color, that calcium had a brick red color and the copper(II) had a green an blue flame. From this data, we can say that the color of the flame becomes lighter as we go down a group (comparing Na+ and K+) but become darker across a period (comparing K+ and Ca2+). This color phenomenon is due to oscillation or swinging of free electrons in any of the atomic orbitals which radiate the energy in form of color.
The second part of my data shows the results for the unknown compound “9” which was a light yellow powdery compound. The first table shows the tests for non-metal ions, and we can see that my substance reacted negatively with AgNO3 and 0.1 mol/dm3 of HCL, but positively with HCL and BaCl2, meaning that my substance contained SO42- as a nonmetal ion. My substance turned a light clear yellow with AgNO3 but no precipitate formed. Since the solution was clear and no precipitate formed, none of the halide ions would be present. With 0.1 mol/dm3 of HCL, my substance turned a clear yellow (resembling lemonade) and again no precipitate formed. This meant that my substance was not a carbonate. When testing for metal ions in my unknown substance, all the flame tests were negative as there was no change in color. When adding 0.1 mol/dm3 NaOH the unknown substance turned a light dark orange with a precipitate. This matched the reaction of Fe3+, so my substance contained Fe(III). I finally had a metal ion that reacted positively and a non metal that reacted positively so my final ionic compound was:
Fe3+SO4-2 --> Fe2(SO4)3 : called Iron(III) Sulfate.
Fortunately, my hypotheses were correct. I had hypothesized that if I was testing for halide ions using 2cm3 ±0.05 of AgNO3, then a colorful precipitate would form but no bubbles would be present. Indeed, this is precisely what occurred. For Cl-, Br-, I-, a colorful precipitate did form, respectively of the colors white, light green, light yellow and no bubble formed. Secondly I had hypothesized that if I was testing for sulphate ions (SO42) using 2cm3 ±0.05 of 0.1 mol/dm3 HCL and BaCl2, then a precipitate would form but no bubbles will be present. This is exactly what occurred and a white, dense precipitate formed with no bubbles. Thirdly, I hypothesized that if I was testing for carbonates (CO32-), using 2cm3 ±0.05 of 0.1 mol/dm3 HCL, then many bubbles would form but no precipitate. In fact, when 2cm3 ±0.05 of dilute hydrochloric acid was added, many white bubbles formed but no precipitate. In my fourth hypothesis I inferred that if I was testing for metal ions (Cu2+, Fe2+, Fe3+) using 3cm3 ±0.05 of 0.1 mol/dm3 NaOH, then a dark colored precipitate would form. In fact, three dark precipitates formed where Cu2+ had a dark precipitate on top, Fe2+ had a bronze precipitate with small black pieces and Fe3+ had a darker orange precipitate. In my fifth hypothesis I inferred that if I was testing for ammonium ions (NH4+) using 2cm3±0.05 0.1 mol/dm3 NaOH and heat, then the litmus paper would turn the color green. This is precisely what occurred as there was an alkaline solution and the pH was green showing a basic solution. In my last hypothesis for part 1 of the experiment I had said that if I was testing for Na+, K+, Ca2+, Cu2+ using a nichrome wire with HCL over a Bunsen burner, then colored flames would be produced. To my amazement, many colored flames were produced and sodium had an orange flame, potassium a lilac flame, calcium a brick red flame and copper (II) a green and blue flame.
For part 2 of the experiment, I had hypothesized that if I was given test tube number nine with an unknown compound, then the compound will be formed by Na+ and SO42- (Sodium Sulfate). I hypothesized this because I knew that sulfur had a yellow colour (thinking this was the color of sulfate) and mixed with sodium which is white would yield to a light yellow- which was the color of my unknown compound. But my hypothesis turned out to be partially correct as the tests indicated that Sodium Sulfate was present but no sodium (as it did not react with the flame) but Iron(III). So the compound was Iron(III) Sulfate.
I learned many important things from this experiment, being that I learned how to discover the identity of an unknown substance based on my observations on various metal and non-metal ions, I improved my observational skills and I learned that the different flame colors are due to oscillation or swinging of free electrons in any of the atomic orbitals which radiate the energy in form of color.
Evaluation and Improvements
There were many sources of error in this experiment, the first being that the solutions may have been old and so the reaction with the metal or non-metal ions would not have been so evident, affecting our results. Especially if silver nitrate had been exposed to a lot of sunlight, this would have caused the solution to become old and not act effectively. A possible source of error is that some people may have used different quantities of solutions rather than respecting the decided quantity 2cm3±0.05, and this may have caused more evident or less evident reactions altering our results. Another source of error was that people did not clean and dip the nichrome wire well enough in the hydrochloric acid causing some of the previous substance to remain. This would have caused the flames to be of a different color causing our results to be inaccurate.
One way I would modify this experiment is that I would add dilute ammonia solution to the silver chloride and silver bromide as the colors of the precipitates are difficult to distinguish. By adding ammonia solution, the silver chloride would dissolve but silver bromide and silver iodide would not as they are insoluble and this would allow us to distinguish them. Another modification to the lab is that I would insert the test for a (non-metal) nitrate ion (NO3-) and I would add 2cm3±0.05 of sodium hydroxide solution to the nitrate. Then I would add powdered aluminum to the nitrate ion and this would cause the formation of many bubbles and ammonia gas will be given off. If a litmus paper were put on top, it would turn blue meaning that it is a basic solution. Another way I would modify the experiment is that I would use the flame test to compare a calcium compound and a magnesium compound. The magnesium compound would have no effect in the flame while the calcium compound would create a brick red flame.
"Chemystery: Atoms and Molecules: Ionic Compounds." Oracle ThinkQuest Library. N.p., n.d. Web. 9 Nov. 2009. <http://library.thinkquest.org/3659/atommole/ionic.html>.
Yoder, Claude H.. Ionic Compounds: Applications of Chemistry to Mineralogy. New York: Wiley- Interscience, 2006. Print.
Here's what a teacher thought of this essay
This is a very complex GCSE experiment, using aspects of chemistry that would not usually be encountered until AS level. It is very detailed, uses excellent results tables and makes valid, repeatable conclusions. Overall, this piece of work is 5*