Planning and Designing Chemistry Lab- To investigate the rate of reaction when calcium carbonate powder and lumps reacts with an acid

Planning and Designing Chemistry Lab- To investigate the rate of reaction when calcium carbonate powder and lumps reacts with an acid

Problem:

You are provided with calcium carbonate in the form of lumps and as a powdered substance, both dilute hydrochloric acid and sulphuric acid are available to you. Plan and design an investigation that you would carry out to compare the rate of reactions of both forms of calcium carbonate with an acid

Hypothesis: Powdered calcium carbonate will react faster with an acid than calcium carbonate lumps

Aim:

To investigate the rate of reaction when calcium carbonate powder and lumps reacts with an acid.

Apparatus/Materials:

thistle funnel, tap, gas syringe, flask, hydrochloric acid, sulphuric acid, calcium carbonate (lumps and powder) stop-watch,

Method:

  1. Set up apparatus as shown in diagram below
  2. Measure 10g of CaCO3 lumps and add it to the flask. Add 50cm3 of HCl to the flask. Record the volume of gas produced every 30sec
  3. Repeat this step for the powdered CaCO3

Please Note: HCl was used in preference to the H2SO4 as, CaCO3 when mixed with H2SO4 forms a slightly soluble salt and thus the reaction will end quickly.

Chemistry diagram 6

Variables:

  • Controlled - volume of acid, concentration of acid, and mass of CaCO3 lumps and powder
  • Manipulated: sample of CaCO3 lumps and CaCO3 powder
  • Responding – volume of gas produced

 Expected Results:Chemistry diagram

Chemistry diagram 7

Interpretation of Results:

The graph shows how much CO2 gas evolved from the experiment. The difference in shapes is because the powder caused the reaction to go at a faster rate than the lumps. The reactions finished at the same time because the masses were the same.

If at the end of the experiment the CaCO3 powder will react at a faster rate than the CaCO3 lumps then the hypothesis would be acceptable.

Sources of Error / Limitations:

Some of the gas escaped whilst being measured this skewing the results.

Assumption: Calcium carbonate reacts faster in the powdered form.

 

 

Planning and Design Chemistry Lab – Plan and Design an experiment that can be used to determine the contents of four bottles found in your home

Planning and Design Chemistry Lab – Plan and Design an experiment that can be used to determine the contents of four bottles found in your home

Problem:

You venture in a cellar one day and found four stoppered bottles. The labels were found on the floor next to the bottles and were marked: dilute sulphuric acid, sodium carbonate, calcium hydroxide and distilled water. If you were handed a pack of blue litmus paper and a rack of test tube, plan and design an experiment that can be used to determine the contents of each bottle. Label the bottles A, B, C and D before you begin your investigation.

Hypothesis:

Bottle A contains Sodium carbonate (Na2CO3), Bottle B contains (H2O), Bottle C contains Sulphuric acid (H2SO4) and Bottle D contains Ca(OH)2.

Aim:

To determine if the contents in Bottle A is Na2CO3, Bottle B is H2O, Bottle C is H2SO4 and Bottle D is Ca(OH)2 using litmus paper

Apparatus / Materials:

blue litmus paper, test-tube, test-tube rack and solutions in bottle

Method:

  1. Label the test-tube A-D
  2. Pour out 5cm of each solution in its corresponding test-tube.
  3. Place a blue litmus paper in each test-tube and observe for a color change to red
  4. The solution that changed the litmus paper to red would be the H2SO4, which is the acid.
  5. Label the corresponding solution as H2SO4.
  6. The test-tubes that were not identified will now be tested.
  7. Use the H2SO4 to change 3 litmus papers to red.
  8. Place the now red litmus papers in the remaining solutions. Observe for a color change to blue. Two of your solution will change to blue. The one that does not change is the water
  9. Label the corresponding bottle as H2O, and set it aside.
  10. The remaining two (2) solutions are the calcium hydroxide and the sodium carbonate.
  11. Add 2 cm of the H2SO4 to the 2 unknown test-tube and note your observations.
  12. The test-tube that showed signs of bubbling/effervescence would be the carbonate, Na2CO3.
  13. Label the corresponding bottle as Na2CO3 and the final bottle as hydroxide.

Expected Results:

Chemistry diagram 3

This test showed that one of the solutions turned the blue litmus paper red. This was identified as the sulphuric acid (H2SO4).

Chemistry diagram 4

 

This test would show two solutions changing the litmus paper to blue, the one that was not changed would be the water. Use the acid to test the remaining two unknowns to differentiate between them.

Chemistry diagram 5

Variables

 

  • Manipulated- Test carried out on the solutions
  • Controlled- Amount of solutions used
  • Responding- reaction with the litmus paper and the acid

 

Treatment of results

 

Label the bottles corresponding to the observations made when each was tested with litmus paper and the sulphuric acid.

 

Interpretation of results

 

If the contents in Bottle A reacted with sulphuric acid then it should be the sodium carbonate, if the contents of Bottle B did not change the color of the red litmus paper then it was the water. If the contents of Bottle C changed the blue litmus paper to red then it would be the acid and if the contents of Bottle D did not react with the acid then it would be the hydroxide.

 

The data collected would then support the hypothesis. If the contents of each of the bottles reacted differently with the blue litmus paper then the data collected does not support your hypothesis.

 

Planning and Designing Chemistry Lab- Different brands of permanent markers contain the same dyes in their black inks; however the markers are insoluble in water

Planning and Designing Chemistry Lab- Different brands of permanent markers contain the same dyes in their black inks; however the markers are insoluble in water

Problem:

Different brands of permanent markers contain the same dyes in their black inks; however the markers are insoluble in water.

Hypothesis:

Different brands of markers contain the same dyes in their black ink

Aim:

To investigate if four (4) different brands of permanent markers contains the same dyes in their black inks using paper chromatography.

Apparatus:

4 different brands of markers, alcohol, filter paper, scissors, ruler, beakers

Method

1. Fill a beaker with 30ml alcohol

2. Cut a rectangular strip of filter paper

3. Using one of the makers place a dot 1cm from the end of the filter paper.

4. Partially immerse the tip of the filter paper closest to the ink spot in the beaker with the alcohol.

5. Record your observations.

6. Allow the filter paper to dry.

7. Repeat the procedure for the remaining three (3) markers and ensure that the dot is of the same size for each marker.

8. Tabulate results and calculate the retention factor (Rf) of each of the markers

Variables

  • Manipulated- different brands of marker
  • Controlled- the volume of alcohol and the size of the dots
  • Dependent/Responding – the retention factor

Chemistry diagram 2

All four brands of marker will contain the same dyes therefore the hypothesis is true

Assumption:

The four brands of markers contain the same types of dye.

 

Percentage of water in a hydrate- Copper(II) sulphate pentahydrate (CuSO4 ·5 H2O)

 Percentage of water in a hydrate- Copper(II) sulphate pentahydrate (CuSO4 ·5 H2O)

Aim / Objective:

To determine the percentage of water in a hydrate

Introduction: 

Many pure substances combine with water in a fixed mole ratio to yield compounds called hydrates. For example, copper sulphate combines with water to form crystalline, CuSO4. 5H2O, which is a stable compound at normal atmospheric conditions. All pure samples of this hydrate show the same percentage of water by analysis. Thus, this hydrated compound obeys the law of constant composition. Upon heating a sample of such a hydrate, it may lose all its water of hydration and revert to the anhydrous salt. Substances which have adsorbed water on the surface do not show constant composition and therefore are not hydrates. An example of this would be common table salt, NaCl, which becomes very sticky on humid, summer days. In these cases, the percentage of water is not constant for all samples of a particular compound, and the water is not chemically bonded as part of the crystal structure.

Other examples of hydrates are Nickel (II) sulfate hexahydrate – (NiSO4 ·6 H2O), lithium perchlorate trihydrate ( LiClO4 . 3H2), aluminum potassium sulfate dodecahydrate – (AlK(SO4)2 ·12 H2O) and magnesium carbonate pentahydrate – (MgCO3 ·5 H2O.)

In this experiment the percentage of water in a hydrate will be determined in an known hydrate. Water is removed from the hydrate by heating an accurately weighed hydrate sample until the residue has reached a constant weight. The percentage of water in the sample is calculated by using the weight of water lost and the initial hydrate sample weight multiplied by 100.

Materials/ Apparatus:

Porcelain crucible and cover, clay triangle, tripod stand, Bunsen burner, flint, tongs,1.000g  Copper(II) sulphate pentahydrate (CuSO4 ·5 H2O(s) , watch glass

 Method / Procedure:

  1. Clean and dry a porcelain crucible and cover.
  2. Place the empty crucible on a covered crucible on a clay triangle supported by a ring on a ring stand.
  3. Heat the crucible and cover in the hottest flame of the Bunsen burner for 5 minutes. Ensure that a dull red glow is observed on the crucible and cover.
  4. Cool the crucible and cover to room temperature for approximately 15 minutes.
  5. Using crucible tongs transfer the crucible and cover to a watch glass and weigh them to the nearest 0.001g.
  6. Add 1g of the 1.000g CuSO4 ·5 H2O(s) to the crucible and weigh the covered crucible to the nearest 0.001g.
  7. Place the covered crucible on the clay triangle with the cover slightly opened.
  8. Heat the crucible gently for a few minutes. Continue to heat for 15 minutes.
  9. Then allow the crucible to cool on the triangle after removing the flame until it reaches room temperature.
  10. Transfer it to the watch glass and weigh the covered crucible to the nearest 0.001g.
  11. Reheat the crucible and contents for about 5minutes. Cool, and then weigh again.
  12. Repeat this heating, cooling and weighing sequence for a total of two readings .
  13. Tabulate results, and complete calculations for the percentage of water in the copper(II) sulphate.

Suggested Results: 

Table showing the results of the heating and cooling of Copper(II) sulphate pentahydrate (CuSO4 ·5 H2O) to find the percentage hydrate in water

Table showing the results of the heating and cooling of Copper(II) sulphate pentahydrate (CuSO4 ·5 H2O) to find the percentage hydrate in water

Use the results from your experiment. Your teacher may request that you use a evaporating dish instead of a crucible and cover which would may the results above slightly different but the concept is the same.

percentage hydrate snippet

Calculations/Answers/Discussion:

  1. Write the formula for the reaction

>>>CuSO4 ·5 H2O(s)  +  HEAT =  CuSO4 (s)  +  5 H2O (g)

2. Calculate the experimental measurement of the percent hydration:

  •       Mass of hydrate before heating =  1.000g
  •       Mass of hydrate after heating = 0.6400g
  •       Difference- mass of water lost = 0.3600g

Experimental Measurement of percent hydrate

  • (0.3600g/1.000g) x 100=36%

3. Calculate the theoretical percentage hydration from the formula.

percentage hydrate snippet2

 

 

 

 

 

 

 

4.Using the theoretical value and the experimental values calculate the percent error

percentage hydrate snippet3

 

 

 

 

 

 

Source of Error/ Limitations/ Assumptions: 

- Allowing the crucible to cool to room temperature before weighting as if not cooled then convection currents will lower the mass and resulting in incorrect results

Effect of Concentration on the rate of reaction of magnesium ribbon in hydrochloric acid

Effect of concentration on rate of reaction of magnesium ribbon in hydrochloric acid

Aim / Objective: 

To investigate the effect of concentration on rate of reaction of magnesium ribbon in hydrochloric acid.

Abstract:  

The rate of reaction was determined by measuring the time required for a given amount of magnesium metal to be consumed by Hydrochloric acid (HCl) solution of varying concentrations.

Introduction: 

The rate of a chemical reaction is the time required for a given quantity of reactant(s) to be changed to product(s). The unit of time may be seconds, minutes, hours, days or years.

The rate is affected by several factors, some of which are listed as follows:

(1) Nature of the reactants, i.e., one metal may react vigorously with acid while another does not react.

(2) The particle size of the reactants, i.e., a lump of coal burns slowly but powdered coal may explode.

( 3 ) Temperature increases in general increase the rate of reaction, i.e., a 2O°C rise in temperature doubles the reaction rate.

( 4 ) Catalysts affect the rate by using or allowing a different pathway for the reaction to follow.

(5) Concentration affects the rate of reaction, i.e., if the concentration of one of the reactants is doubled and is an integral part of the reaction then rate increases appropriately.

Some reactions are fast and other reactions are slow. The rate of a specific reaction can be found only by experiment.

Apparatus/Materials:

Magnesium ribbon , ruler, scissors, analytical balance, sandpaper, hydrochloric acid, measuring cylinder, graduated cylinder, distilled water, glass stirring rod.

Method / Procedure:

  1. Clean a 25cm length magnesium ribbon lightly by using sandpaper to remove the surface oxide layer. Cut the clean ribbon into five equal pieces of 6cm using a ruler.
  2. Weigh the five pieces together and determine the mass of one piece assuming all six are the same.
  3. Make up 30mL of each of the following five solutions of hydrochloric acid (HCl) with water: 2.0M, 1.5M, 1.0M, 0.5M and 0.25M. To do this, calculate the volumes of 3.0 mol L-1 stock hydrochloric acid solution and water that must be mixed using the dilution equation: (number mL stock HCL solution) x [HCL]= (3.0mL diluted solution) x [HCL] diluted
  4. Measure the calculated volume of stock hydrochloric acid solution using 10mL graduated cylinder. Pour this acid into a 50-mL graduated cylinder and then dilute to the 30-mL mark by carefully adding water from a bottle.
  5. Make up the other for solutions in a similar way. Pour each diluted solution into a 50-mL beaker.
  6. Drop a piece of magnesium into the 2.0 mol L-1 acid solution and start timing the reaction. Stir gently at first using the glass stirring rod to make sure the metal does not stick to the sides of the beaker.
  7.  Measure the time elapsed when the reaction stops and record the time.
  8. Repeat the procedure with the other four acid solutions.
  9. **Your instructor may require you to plot a graph

Suggested Results:

Table showing the effect of varying concentration of HCl on rate of reaction of Magnesium ribbon

Table showing the effect of varying concentration of HCl on rate of reaction of Magnesium ribbon

Discussion:

Write a balanced equation for the reaction

>>> Mg(s) + 2HCl(aq)      =         MgCl2(aq)   + H2(g)

 

Write the ionic equation for the reaction

>>>Mg(s) + 2H+(aq)           =        Mg2+(aq)  + H2(g)

 

Identify the variables in the experiment

  • The manipulated variables were the HCl and the water
  • The responding variable was time
  • The controlled variable was the magnesium ribbon

Based on your experimental data, make a general statement about the effect of concentration of reactants on time and reaction rate

>>>Concentration affects the rate of reaction. Therefore over time as the concentration of HCl increased then the rate of the reaction also increased.

Source of Error/ Limitations/ Assumptions:

  •  Inaccurate timing. Time may be lost during the experiment between the times taken to notice the cross has disappeared to the actual stopping of the watch.
  • Inaccurate measurement of reactants will affect the overall rate of reaction.

Reactivity series of metals|Reaction with hydrochloric acid and sulphuric acid (Mg, Zn. Al, Fe, Pb, Cu)

Reactivity series of metals|Reaction with hydrochloric acid and sulphuric acid (Mg, Zn. Al, Fe, Pb, Cu)

Aim / Objective: 

To investigate the reaction of some metals: Mg, Zn. Al, Fe, Pb, Cu, with hydrochloric acid and sulphuric acid.

Apparatus/ Materials:

metals of magnesium (Mg) turnings, zinc (Zn) granules, aluminum (Al) turnings, iron (Fe) filings, lead (Pb) foil, copper (Cu), 6 test-tubes, test-tube rack, dilute hydrochloric acid (HCl), dilute sulphuric acid (H2SO4), splints, Bunsen burner and spatula.

Method / Procedure:

  1. Half fill a test- tube with dilute HCl.
  2. Add a spatula full of magnesium turnings to the acid, place cork stopper in the mouth of test-tube and observe for effervescence of gas.
  3. Record your observations. Heat gently under the Bunsen burner if no reaction is taking place or if it is too slow.
  4. Remove cork-stopper and place a lighted splint at the mouth of test-tube. Note the reaction.
  5. Repeat steps 1-4 with the remaining metals and dilute HCl. Then repeat the same procedure with dilute sulphric acid.
  6. Tabulate your results.

Suggested Results: 

reactivity table 1 reactivity table 2

 

Discussion:

reactivity series equations

reactivity series results

 

 

 

 

Source of Error/ Limitations/ Assumptions:  

  • Overheating of test-tube.
  • Allowing some gas to escape when HCl and H2SO4 is added

Oxidizing and Reducing Agents

 Identification of oxidizing and reducing agents using solutions of acidified potassium dichromate (VI), hydrogen peroxide, acidified potassium iodide, acidified potassium manganate (VII), bleach

Aim / Objective: 

To identify oxidizing and reducing agents using solutions of acidified potassium dichromate (VI), hydrogen peroxide, acidified potassium iodide, acidified potassium manganate (VII), bleach.

Introduction: 

An oxidizing agent, or oxidant, is one that gains electrons and is reduced in a chemical reaction. They are also known as electron acceptors, the oxidizing agent is normally in one of its higher possible oxidation states because it will gain electrons and be reduced. Some examples of oxidizing agents include potassium nitrate, halogens and nitric acid.

A reducing agent, or reductant, loses electrons and is oxidized in a reaction.  A reducing agent is usually in one of its lower possible oxidation states and is the electron donor. A reducing agent is oxidized because it loses electrons in the redox reaction. Examples of reducing agents include formic acid, sulfite compounds and earth metals.

Materials/ Apparatus:

Hydrogen peroxide (H2O2), acidified iron (II) Sulphate (FeSO4), acidified potassium iodide (KI), acidified potassium dichromate(VI) (K2Cr2O7) acidified potassium manganate (VII) (KMnO4), sulphuric acid (H2SO4), bleach (NaClO), 7 test tubes, test tube rack, glass stirring rod, metal boiling tube holder, droppers, sodium hydroxide (NaOH).

Method / Procedure:

  1. Label test tubes 1-7 and place in test tube rack
  2. Half fill test tube 1 with acidified KI and use a dropper to add acidified potassium dichromate (VI) (K2Cr2O7) to the solution. Note color change.
  3. Half fill test tube 2 with acidified KI and use a dropper to add acidified potassium manganate (VII) (KMnO4) to the solution. Note color change.
  4. Half fill test tube 3 with acidified KI and use a dropper to add NaClO to the solution. Note color change.
  5. To test tube 4 half fill with hydrogen peroxide (H2O2) and use a dropper to add acidified KMnO4 to the solution. Note color change.
  6. Half fill test tube 5 with KI and use a dropper to add H2O2 to the solution. Note color change.
  7. Half fill test tube 6 with H2O2 and use a dropper to add acidified K2Cr2O7 to the solution. Note color change.
  8. To the final test tube half fill with FeSO4 and use a dropper to add acidified K2Cr2O7 to the test tube. Add sodium hydroxide (NaOH) to the test tube and note color change.
  9. Tabulate results

Suggested Results:

 

oxidizing and reducing agents 1 oxidizing and reducing agents 2 oxidizing and reducing agents 3 oxidizing and reducing agents 4

 

Hope this one helps, remember to keep reading and studying as the weeks go by and your exams will be great.

Qualitative Analysis on Compound X in order to identify cations and anions present

Qualitative Analysis on Compound X in order to identify cations and anions present

Aim / Objective: 

To perform qualitative analysis tests on compound X in order to identify cations and anions present.

Introduction: 

Qualitative analysis is a technique that is used to separate and detect cations and anions in a sample. Anions are atoms or groups of atoms that have gained an electron or electrons. The atoms that form ions most easily are the Group 17 or VII atoms, also called halides: Fluorine (F), Chlorine (Cl), Bromine (Br) and Iodine (I). These form anions with a -1 charge. Oxygen (O), Sulphur (S), Nitrogen (N) and Phosphorus (P) also form anions. Most anions are composed from multiple atoms, and are called polyatomic ions.

Cations are atoms that have lost an electron to become positively charged. For example: Sodium (Na)  has one valence electron, one electron in its outer energy level, so it tends to lose one electron, and to become an ion with a +1 charge

Materials/ Apparatus:

Bunsen burner, red litmus paper, blue litmus paper, distilled water, test tube rack, test tubes, glass stirring rod, metal test tube holder, dilute nitric acid (HNO3), sodium hydroxide (NaOH), potassium iodide (KI), barium chloride (BaCl2), ammonia solution, lead nitrate (Pb(NO3)2), silver nitrate (AgNO3), dilute hydrochloric acid (HCl), splint, dropper, compound X.

Method / Procedure:

  1. Heat solid compound X in a dry test tube.
  2. Test the gas evolved with red and blue litmus paper and a glowing splint.
  3. Record the observations.
  4. To a solution of compound X in a test tube add distilled water and dilute HNO3, then add AgNO3 and note observations.
  5. In a test tube mix solution X with distilled water, then add dilute HNO3 and BaCl2 to the mixture. Record the observations.
  6. In a test tube mix solution X with distilled water and heat the test tube moderately. Add dilute HNO3 and Pb(NO3) to the heated mixture. Record the observations.
  7. Mix a small portion of solution X with distilled water in a test tube. Use a dropper to add NaOH in excess to the mixture and heat test tube moderately. Record your observations.
  8. Mix a small portion of solution X with distilled water in a test tube. Add excess ammonia to the mixture and note the results.
  9.  Half fill a test tube with solution X and add KI(aq). Record the observations.
  10. Half fill a test tube with solution X, add NaOH(aq) and heat the test tube moderately. Dip the glass stirring rod with the mixture in a beaker of HCl. Test the gas evolved with litmus paper. Record the observations.
  11. Tabulate results

Suggested Results: 

qualitative analysis 1 qualitative analysis 2

 

Conclusion:

  1. The cations present In compound X were Al3+ and NH4+
  2. The anions present were SO42- and Cl-

 

Separating a mixture of Oil and Water

Separating a mixture of Oil and Water using a Separating Funnel Aim / Objective: To separate a mixture of cooking oil and water. Abstract: In this experiment, an immiscible mixture of oil and water was separated using a separating funnel. Materials/ Apparatus: measuring cylinder, retort stand, beaker, clamp, separating funnel, conical flask, cooking oil and …

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