Week 1: Chemical Analysis
VCE Chemistry Week 1 - Chemical Analysis (1)
Horizen Education February 8, 2013
1
1.1
Review: Gravimetric and Volumetric Analysis
Gravimetric Analysis
Solubilities Always soluble: Usually soluble: Exceptions:
Na+ , K+ , NO3− , CH3 COO− , NH+ 4 Cl− , I− , Br−
i. Sodium carbonate and Silver nitrate
ii. Iron (II) sulfate and Lead (II) nitrate
iii. Sodium nitrate and Nickel sulfate
iv. Potassium hydroxide and Copper (II) nitrate
v. Sodium sul de and Cadmium sulfate
Ho
Write down the ionic chemical equation when the following pairs of solutions are mixed together. (Take note of their solubilities of the products formed)
Question 1.
riz
Remember that all ionic compounds are soluble to an extent; those classi ed …show more content…
as insoluble really just have extremely low solubility.
en
It is assumed knowledge that AgCl(s) and BaSO4 (s) are not soluble. (Must know for U3)
Ed
1
uc
ati
This form of analysis focuses on measuring the mass of a precipitate to nd the concentration of an ion in a sample. Gravimetric analysis involves forming an ionic precipitate with the ion to be analysed, and then drying and weighing the precipitate to nd the amount of ion in the original sample.
on
Finding ion concentration
Step 1: Weigh sample. Step 2: Dissolve sample in distilled water, and then lter the waste solid out. Step 3: Add an excess of solution to precipitate the ions being analysed Step 4: Filter this precipitate Step 5: Wash the precipitate with deionised water to remove any unwanted ions Step 6: Dry the precipitate at 100◦ C Step 7: Weigh Step 8: Repeat 6-7 until at constant mass.
Example 1. A certain barium halide exists as the hydrated salt BaX2 · 2 H2 O, where X is the halogen. The barium content of the salt can be determined by gravimetric methods. A sample of the halide (0.2650 g) was dissolved in water (200 cm3) and excess sulfuric acid added. The mixture was then heated and held at boiling for 45 minutes. The precipitate (barium sulfate) was ltered o, washed and dried. Mass of precipitate obtained = 0.2533 g. Determine the identity of X.
Conditions
When designing the experiment there are several considerations. The known solution must be able to create a precipitate with a component of the unknown chemical which can be measured. This precipitate must have:
A known formula so that it can be used in calculations involving the reaction equation. A low solubility to ensure that the majority of it can be collected by the lter paper when ltering.
This will minimise the amount of error in measurements.
Be stable when heated so that it can be dried easily.
Ho
2
riz
en
Ed
uc
ati
on
Not form a precipitate with any other ions present in the sample to ensure that measurements are
accurate. As nitrates are soluble in any ionic compound, they are often used in gravimetric analysis.
Possible errors:
Reaction did not precipitate fully and hence a smaller mass of precipitate was formed The precipitate was not rinsed or dried properly Transfer errors led to a smaller amount of precipitate formed Errors in rinsing the precipitate
A sample of lake water near an industrial chemical plant was being analysed for toxic Barium ions by adding 0.050M Silver Sulfate standard solution and reacting to form a precipitate.
Example 2.
a) Predict what would happen to the calculated concentration of Barium ions if: i) The concentration of the AgSO4 standard solution was actually 0.040M. ii) The precipitate was not rinsed properly after ltration. b) What errors could have caused a i) Higher than expected concentration ii) Lower than expected concentration
Note:
Percentage Composition
Percentage composition questions involving a chemical reaction seek to determine the amount, in grams, of a certain chemical found in 100g of a sample. The steps to solve these questions are always the same, and they are as follows: a) Find the reaction formula if required b) Calculate the amount, in moles, of the pure compound c) Find the mass, in grams, of the chemical being analysed d) Divide the mass of the chemical being analysed by the mass of the original sample/mixture Make sure you do not use the mass of the original mixture/sample in your initial calculations! You should only ever need to use the mass of your original sample in the nal step to calculate percentage composition.
A 2.056g sample of Buttery Nuts peanut butter was dissolved, ltered, and reacted with Silver Chloride solution to form a precipitate. The precipitate was then heated to constant mass and found to weigh 0.0452g. Using this information, nd the percentage composition of sodium ions in the original sample.
Example 3.
Ho
riz
3
en
The most commonly used precipitates are AgCl and BaSO4 An excess is added to ensure all ions react. The precipitate is washed to wash o any ions that may have remained on the lter paper. It is then dried to ensure that the no water remains on the lter paper which would in uence the result.
Ed
uc
ati
on
Question 2. The following components are used in common fertilisers as a source of nitrogen. Find the percentage composition, by mass, of nitrogen if a 3.456g sample of fertilizer is found to contain:
a) 1.750g of Ammonium (NH+ ) 4 b) 2.546g of Ammonium Nitrate (NH4 NO3 ) c) 1.075g of Urea (CO(NH2 )2 ) Excessive salt intake in the diet can contribute to high blood pressure and heart disease. The salt content of a 14.96 g sample of powdered chicken soup was determined by dissolving it in water to make a volume of 250.0 mL. A 20.00 mL volume of this stock solution was pipetted into a conical ask and excess silver nitrate was added. The silver chloride precipitate that formed was then ltered, washed and dried. Its mass was 0.246 g. a) Write an ionic equation for the formation of the silver chloride precipitate.
Question 3.
c) Determine the amount of sodium chloride in the 20.00 mL volume of stock solution in the conical ask. d) Calculate the amount of sodium chloride in 250.0 mL of the stock solution. e) What mass of sodium chloride was in the sample?
Question 4. A 3.245g Berokka tablet is treated with nitric acid to produce 600mL of carbon dioxide, measured at SLC. If magnesium carbonate is the active ingredient in the tablet, calculate the percentage composition of magnesium carbonate in the original sample.
Question 5. The reaction involved in thermite is given by the equation: Fe2 O3 + 2 Al − − 2 Fe + Al2 O3 . → If 300g of an impure source of Iron(II) Oxide was reacted with excess aluminium powder to form 179g of Al2 O3 , calculate the percentage composition of iron in the original source of Iron(II) Oxide.
Question 6. A 0.400g sample of silver sulfate (Ag2 SO4 ) and a 0.350g sample of aluminium sulfate (Al2 SO4 ) were dissolved and reacted with barium chloride (BaCl2 ), nd the mass of precipitate formed. Concentrations
A patient in the emergency ward of a hospital who is suering from hyponatremia has a sodium ion concentration in the blood of 0.110M and has a total blood volume of 4.1L. What volume of sodium chloride saline solution (0.9%) would need to be added to the blood to increase the sodium ion concentration to 0.142M?
Question 7.
Example 4.
Express the following concentrations in terms of gL1 4
Ho
riz
en
Ed
uc
ati
b) Calculate the amount, in mol, of silver chloride that was produced. Assume all the chloride in the powdered soup came from sodium chloride (common salt).
on
a) 0.0036 M Ca(OH)2 solution b) 6.3 × 10−5 M Pb(NO3 )2 solution c) 17.5%(w/v) glucose (C6 H12 O6 ) in grape drink
1.2 Volumetric Analysis
A student wishes to analyse the concentration of a sodium hydroxide solution by titrating it with hydrochloric acid from a burette. What would happen to the nal concentration if:
Example 5.
a) She rinsed the pipette with water b) She rinsed the burette with water c) The conical ask was rinsed with hydrochloric acid d) The actual concentration of hydrochloric acid was lower than expected A 0.6417g sample of impure magnesium hydroxide is dissolved in 150mL of 0.1950M HCl solution. An average titre of 19.85mL of a 0.1015M NaOH solution is used to neutralise the excess acid.
Example 6.
i. Write down the chemical equations describing the steps above.
iii. The NaOH solution used in the experiment was in fact lower than 0.1015M, is the value calculated in part ii. greater or lesser than the actual value? Justify your answer.
Question 8. The odour of vinegar is due to acetic acid CH3 COOH , which reacts with sodium hydroxide. If 4.15mL of vinegar needs 52.15mL of 0.115M N aOH to reach the equivalence point in the titration.
Ho
5
riz
ii. Calculate the percentage by mass of magnesium hydroxide in the sample.
en
Ed
uc
ati
on
i. How many grams of acetic acid are there in a 250mL sample of this vinegar?
ii. The value calculated in part i. is higher than the true value. State two possible reason why this error occurred.
Question 9.
c) What amount of the metal carbonate was present in each 20.0mL aliquot? e) Determine the molar mass of the metal carbonate, and hence determine metal X. Depending on the reactants in the acid-base titration, we must choose a suitable indicator that changes colour close to the equivalence point. The following table outlines the dierent indicators and the pH range in which they reach their end point. Titration curves: Table 1: Acid-Base Indicators Name Thymol blue Methyl orange Bromophenol blue Methyl red Bromothymol blue Phenol red Phenolphthalein pH range 1.2-2.8 3.1-4.1 3.0-4.6 4.2-6.3 6.0-7.6 6.8-8.4 8.3-10.0 Acid Colour red red yellow red yellow yellow colourless Base Colour yellow yellow blue yellow blue red red
2 × 10˘2 2 × 10−4 6 × 10−5 8 × 10−6 1 × 10−7 1 × 10−8 5 × 10−10
Ho
Choosing an Indicator
riz
en
d) What amount of metal carbonate was present in the original sample?
Ed
6
b) What amount of sulfuric acid, in mol, was needed to reach the endpoint?
uc
a) Write an equation between the metal carbonate and sulfuric acid.
ati
A solution of a metal carbonate with chemical formula X2 CO3 was prepared by dissolving 2.80g of the solid in 250mL of water. 20.00mL aliquots were titrated with a 0.150M H2 SO4 solution, using methyl orange as the indicator. The average titre was found to be 10.8mL. (Source: Heinemann Chemistry)
on
Ka
Strong acid + strong base: equivalence pH 7
Strong acid + weak base: equivalence pH 3-6
Strong base + weak acid: equivalence pH 8-11
Weak base + weak acid: equivalence unknown (use back titration for analysis)
i. A larger volume of NaOH(aq) is needed to reach the equivalence point in the titration of HNO3
v. The titration curves will essentially be the same after passing the equivalence point vi. Methyl red would be a suitable indicator for both titrations
1.3 Back Titration
Some acids and bases are too weak to analyse using conventional methods. These weak acids and bases must undergo a process called back titration to determine their …show more content…
concentration.
Question 11.
Limestone is known to contain calcium carbonate as an impurity. A sample of limestone was analysed using back titration. A 1.0g sample of limestone was weighed and made into a 50.0mL solution in a conical ask. An excess of 0.395M hydrochloric acid, exactly 50.0mL, was added and allowed to fully react with the CaCO3 . After this reaction, the remaining HCl was titrated with a standard 0.0489M solution of Sodium Hydroxide. The titre was found to be 22.32mL. (Source: Heinemann
Chemistry)
a) Write balanced chemical equations for the two reactions that occur. b) Determine the moles of the hydrochloric acid in excess after the reaction with the limestone. c) Calculate the amount of HCl initially added to the limestone solution. 7
Ho
iv. The pH at the beginning of the two titrations will be the same
riz
iii. Phenolphthalein would be suitable indicator for both titrations
en
ii. The pH at the equivalence point in the HNO3 , titration will be lower than the pH at the equivalence point in the CH3 COOH titration
Ed
uc
Question 10.
Determine whether each of the following statements concerning the titration of nitric and acetic acid, both at 500 ml of 0.1 M with 0.05 M solution of NaOH(aq) are true or false.
ati
on
d) Calculate the amount of CaCO3 in the original limestone sample. e) What is the percentage composition of calcium in the limestone? f) In this experiment, the entire 1.0g sample of limestone was reacted. How could you improve the precision of the experiment? Explain. Question 12. A 0.5130g aspirin tablet which was dissolved in 95% ethanol required 27.98mL of 0.1M NaOH for neutralization. Then an additional 42.78mL of 0.1M NaOH was added and the solution was heated to hydrolyze the acetylsalicylic acid. After the reaction mixture was cooled, the excess base was back-titrated with 14.29mL of 0.1056M HCl. CH3 COOC6 H4 COOH(s) + 2 NaOH(aq) −→ CH3 COONa(aq) + HOC6 H4 COONa(aq) + H2 O(l) − NaOH(aq) + HCl(aq) −→ NaCl(aq) + H2 O(l) − a) Calculate the amount of HCl required to hydrolyze the acetylsalicylic acid. b) Calculate the amount of NaOH initially added to the aspirin solution.
1.4
The most important formula to remember is pH=log10 [H3 O+ ] Example 7. Calculate the pH of the solution that results when 40.0 mL of 0.100 M NaOH is: a) Diluted with 200.0 mL of distilled water. b) Mixed with 20.0 mL of 0.200 M HCl solution. c) Mixed with 20.0 mL of 0.150 M HNO3 Question 14. 100 cm3 of a magnesium hydroxide solution required 4.5 cm3 of sulphuric acid (of concentration 0.100 mol dm-3) for complete neutralisation. [atomic masses: Mg = 24.3, O = 16, H = 1) a) Give the equation for the neutralisation reaction. b) Calculate the moles of sulphuric acid neutralised. c) Calculate the moles of magnesium hydroxide neutralised. d) Calculate the concentration of the magnesium hydroxide in mol dm-3 (molarity). e) Calculate the concentration of the magnesium hydroxide in g cm-3. Question 15. A 0.484 g sample of impure NH4 Br is treated with 25.00 mL of 0.2050 M NaOH and heated to drive o the NH3 . The unreacted NaOH in the reaction mixture after heating required 9.95 mL of 0.0525 M H2 C2 O4 to neutralize. How many grams of NH4 Br were in the original sample? Question 16. Calculate how many mL of 0.450M NaOH solution must be added to 40mL of pH 1.2 HNO3 to neutralise. 8
Acid/base and pH
Ho
riz
en
A 0.4755 g sample containing ammonium oxalate (NH4 )2 C2 O4 and inert compounds has dissolved in water and made alkaline with KOH. The liberated NH3 was distilled into 50 mL of 0.1007 N H2 SO4 . The excess H2 SO4 was back titrated with 11.13 mL of 0.1214 N NaOH. Calculate the %N and %(NH4 )2 C2 O4 in the sample. 2 Note: Oxalate = C2 O4
Question 13.
Ed
uc
e) What % w/w of the tablet is acetylsalicylic acid?
ati
d) How many grams of acetylsalicylic acid are in the tablet?
on
c) Calculate the amount of aspirin that reacted with NaOH.
Horizen Education February 8, 2013
1
1.1
Review: Gravimetric and Volumetric Analysis
Gravimetric Analysis
Solubilities Always soluble: Usually soluble: Exceptions:
Na+ , K+ , NO3− , CH3 COO− , NH+ 4 Cl− , I− , Br−
i. Sodium carbonate and Silver nitrate
ii. Iron (II) sulfate and Lead (II) nitrate
iii. Sodium nitrate and Nickel sulfate
iv. Potassium hydroxide and Copper (II) nitrate
v. Sodium sul de and Cadmium sulfate
Ho
Write down the ionic chemical equation when the following pairs of solutions are mixed together. (Take note of their solubilities of the products formed)
Question 1.
riz
Remember that all ionic compounds are soluble to an extent; those classi ed …show more content…
as insoluble really just have extremely low solubility.
en
It is assumed knowledge that AgCl(s) and BaSO4 (s) are not soluble. (Must know for U3)
Ed
1
uc
ati
This form of analysis focuses on measuring the mass of a precipitate to nd the concentration of an ion in a sample. Gravimetric analysis involves forming an ionic precipitate with the ion to be analysed, and then drying and weighing the precipitate to nd the amount of ion in the original sample.
on
Finding ion concentration
Step 1: Weigh sample. Step 2: Dissolve sample in distilled water, and then lter the waste solid out. Step 3: Add an excess of solution to precipitate the ions being analysed Step 4: Filter this precipitate Step 5: Wash the precipitate with deionised water to remove any unwanted ions Step 6: Dry the precipitate at 100◦ C Step 7: Weigh Step 8: Repeat 6-7 until at constant mass.
Example 1. A certain barium halide exists as the hydrated salt BaX2 · 2 H2 O, where X is the halogen. The barium content of the salt can be determined by gravimetric methods. A sample of the halide (0.2650 g) was dissolved in water (200 cm3) and excess sulfuric acid added. The mixture was then heated and held at boiling for 45 minutes. The precipitate (barium sulfate) was ltered o, washed and dried. Mass of precipitate obtained = 0.2533 g. Determine the identity of X.
Conditions
When designing the experiment there are several considerations. The known solution must be able to create a precipitate with a component of the unknown chemical which can be measured. This precipitate must have:
A known formula so that it can be used in calculations involving the reaction equation. A low solubility to ensure that the majority of it can be collected by the lter paper when ltering.
This will minimise the amount of error in measurements.
Be stable when heated so that it can be dried easily.
Ho
2
riz
en
Ed
uc
ati
on
Not form a precipitate with any other ions present in the sample to ensure that measurements are
accurate. As nitrates are soluble in any ionic compound, they are often used in gravimetric analysis.
Possible errors:
Reaction did not precipitate fully and hence a smaller mass of precipitate was formed The precipitate was not rinsed or dried properly Transfer errors led to a smaller amount of precipitate formed Errors in rinsing the precipitate
A sample of lake water near an industrial chemical plant was being analysed for toxic Barium ions by adding 0.050M Silver Sulfate standard solution and reacting to form a precipitate.
Example 2.
a) Predict what would happen to the calculated concentration of Barium ions if: i) The concentration of the AgSO4 standard solution was actually 0.040M. ii) The precipitate was not rinsed properly after ltration. b) What errors could have caused a i) Higher than expected concentration ii) Lower than expected concentration
Note:
Percentage Composition
Percentage composition questions involving a chemical reaction seek to determine the amount, in grams, of a certain chemical found in 100g of a sample. The steps to solve these questions are always the same, and they are as follows: a) Find the reaction formula if required b) Calculate the amount, in moles, of the pure compound c) Find the mass, in grams, of the chemical being analysed d) Divide the mass of the chemical being analysed by the mass of the original sample/mixture Make sure you do not use the mass of the original mixture/sample in your initial calculations! You should only ever need to use the mass of your original sample in the nal step to calculate percentage composition.
A 2.056g sample of Buttery Nuts peanut butter was dissolved, ltered, and reacted with Silver Chloride solution to form a precipitate. The precipitate was then heated to constant mass and found to weigh 0.0452g. Using this information, nd the percentage composition of sodium ions in the original sample.
Example 3.
Ho
riz
3
en
The most commonly used precipitates are AgCl and BaSO4 An excess is added to ensure all ions react. The precipitate is washed to wash o any ions that may have remained on the lter paper. It is then dried to ensure that the no water remains on the lter paper which would in uence the result.
Ed
uc
ati
on
Question 2. The following components are used in common fertilisers as a source of nitrogen. Find the percentage composition, by mass, of nitrogen if a 3.456g sample of fertilizer is found to contain:
a) 1.750g of Ammonium (NH+ ) 4 b) 2.546g of Ammonium Nitrate (NH4 NO3 ) c) 1.075g of Urea (CO(NH2 )2 ) Excessive salt intake in the diet can contribute to high blood pressure and heart disease. The salt content of a 14.96 g sample of powdered chicken soup was determined by dissolving it in water to make a volume of 250.0 mL. A 20.00 mL volume of this stock solution was pipetted into a conical ask and excess silver nitrate was added. The silver chloride precipitate that formed was then ltered, washed and dried. Its mass was 0.246 g. a) Write an ionic equation for the formation of the silver chloride precipitate.
Question 3.
c) Determine the amount of sodium chloride in the 20.00 mL volume of stock solution in the conical ask. d) Calculate the amount of sodium chloride in 250.0 mL of the stock solution. e) What mass of sodium chloride was in the sample?
Question 4. A 3.245g Berokka tablet is treated with nitric acid to produce 600mL of carbon dioxide, measured at SLC. If magnesium carbonate is the active ingredient in the tablet, calculate the percentage composition of magnesium carbonate in the original sample.
Question 5. The reaction involved in thermite is given by the equation: Fe2 O3 + 2 Al − − 2 Fe + Al2 O3 . → If 300g of an impure source of Iron(II) Oxide was reacted with excess aluminium powder to form 179g of Al2 O3 , calculate the percentage composition of iron in the original source of Iron(II) Oxide.
Question 6. A 0.400g sample of silver sulfate (Ag2 SO4 ) and a 0.350g sample of aluminium sulfate (Al2 SO4 ) were dissolved and reacted with barium chloride (BaCl2 ), nd the mass of precipitate formed. Concentrations
A patient in the emergency ward of a hospital who is suering from hyponatremia has a sodium ion concentration in the blood of 0.110M and has a total blood volume of 4.1L. What volume of sodium chloride saline solution (0.9%) would need to be added to the blood to increase the sodium ion concentration to 0.142M?
Question 7.
Example 4.
Express the following concentrations in terms of gL1 4
Ho
riz
en
Ed
uc
ati
b) Calculate the amount, in mol, of silver chloride that was produced. Assume all the chloride in the powdered soup came from sodium chloride (common salt).
on
a) 0.0036 M Ca(OH)2 solution b) 6.3 × 10−5 M Pb(NO3 )2 solution c) 17.5%(w/v) glucose (C6 H12 O6 ) in grape drink
1.2 Volumetric Analysis
A student wishes to analyse the concentration of a sodium hydroxide solution by titrating it with hydrochloric acid from a burette. What would happen to the nal concentration if:
Example 5.
a) She rinsed the pipette with water b) She rinsed the burette with water c) The conical ask was rinsed with hydrochloric acid d) The actual concentration of hydrochloric acid was lower than expected A 0.6417g sample of impure magnesium hydroxide is dissolved in 150mL of 0.1950M HCl solution. An average titre of 19.85mL of a 0.1015M NaOH solution is used to neutralise the excess acid.
Example 6.
i. Write down the chemical equations describing the steps above.
iii. The NaOH solution used in the experiment was in fact lower than 0.1015M, is the value calculated in part ii. greater or lesser than the actual value? Justify your answer.
Question 8. The odour of vinegar is due to acetic acid CH3 COOH , which reacts with sodium hydroxide. If 4.15mL of vinegar needs 52.15mL of 0.115M N aOH to reach the equivalence point in the titration.
Ho
5
riz
ii. Calculate the percentage by mass of magnesium hydroxide in the sample.
en
Ed
uc
ati
on
i. How many grams of acetic acid are there in a 250mL sample of this vinegar?
ii. The value calculated in part i. is higher than the true value. State two possible reason why this error occurred.
Question 9.
c) What amount of the metal carbonate was present in each 20.0mL aliquot? e) Determine the molar mass of the metal carbonate, and hence determine metal X. Depending on the reactants in the acid-base titration, we must choose a suitable indicator that changes colour close to the equivalence point. The following table outlines the dierent indicators and the pH range in which they reach their end point. Titration curves: Table 1: Acid-Base Indicators Name Thymol blue Methyl orange Bromophenol blue Methyl red Bromothymol blue Phenol red Phenolphthalein pH range 1.2-2.8 3.1-4.1 3.0-4.6 4.2-6.3 6.0-7.6 6.8-8.4 8.3-10.0 Acid Colour red red yellow red yellow yellow colourless Base Colour yellow yellow blue yellow blue red red
2 × 10˘2 2 × 10−4 6 × 10−5 8 × 10−6 1 × 10−7 1 × 10−8 5 × 10−10
Ho
Choosing an Indicator
riz
en
d) What amount of metal carbonate was present in the original sample?
Ed
6
b) What amount of sulfuric acid, in mol, was needed to reach the endpoint?
uc
a) Write an equation between the metal carbonate and sulfuric acid.
ati
A solution of a metal carbonate with chemical formula X2 CO3 was prepared by dissolving 2.80g of the solid in 250mL of water. 20.00mL aliquots were titrated with a 0.150M H2 SO4 solution, using methyl orange as the indicator. The average titre was found to be 10.8mL. (Source: Heinemann Chemistry)
on
Ka
Strong acid + strong base: equivalence pH 7
Strong acid + weak base: equivalence pH 3-6
Strong base + weak acid: equivalence pH 8-11
Weak base + weak acid: equivalence unknown (use back titration for analysis)
i. A larger volume of NaOH(aq) is needed to reach the equivalence point in the titration of HNO3
v. The titration curves will essentially be the same after passing the equivalence point vi. Methyl red would be a suitable indicator for both titrations
1.3 Back Titration
Some acids and bases are too weak to analyse using conventional methods. These weak acids and bases must undergo a process called back titration to determine their …show more content…
concentration.
Question 11.
Limestone is known to contain calcium carbonate as an impurity. A sample of limestone was analysed using back titration. A 1.0g sample of limestone was weighed and made into a 50.0mL solution in a conical ask. An excess of 0.395M hydrochloric acid, exactly 50.0mL, was added and allowed to fully react with the CaCO3 . After this reaction, the remaining HCl was titrated with a standard 0.0489M solution of Sodium Hydroxide. The titre was found to be 22.32mL. (Source: Heinemann
Chemistry)
a) Write balanced chemical equations for the two reactions that occur. b) Determine the moles of the hydrochloric acid in excess after the reaction with the limestone. c) Calculate the amount of HCl initially added to the limestone solution. 7
Ho
iv. The pH at the beginning of the two titrations will be the same
riz
iii. Phenolphthalein would be suitable indicator for both titrations
en
ii. The pH at the equivalence point in the HNO3 , titration will be lower than the pH at the equivalence point in the CH3 COOH titration
Ed
uc
Question 10.
Determine whether each of the following statements concerning the titration of nitric and acetic acid, both at 500 ml of 0.1 M with 0.05 M solution of NaOH(aq) are true or false.
ati
on
d) Calculate the amount of CaCO3 in the original limestone sample. e) What is the percentage composition of calcium in the limestone? f) In this experiment, the entire 1.0g sample of limestone was reacted. How could you improve the precision of the experiment? Explain. Question 12. A 0.5130g aspirin tablet which was dissolved in 95% ethanol required 27.98mL of 0.1M NaOH for neutralization. Then an additional 42.78mL of 0.1M NaOH was added and the solution was heated to hydrolyze the acetylsalicylic acid. After the reaction mixture was cooled, the excess base was back-titrated with 14.29mL of 0.1056M HCl. CH3 COOC6 H4 COOH(s) + 2 NaOH(aq) −→ CH3 COONa(aq) + HOC6 H4 COONa(aq) + H2 O(l) − NaOH(aq) + HCl(aq) −→ NaCl(aq) + H2 O(l) − a) Calculate the amount of HCl required to hydrolyze the acetylsalicylic acid. b) Calculate the amount of NaOH initially added to the aspirin solution.
1.4
The most important formula to remember is pH=log10 [H3 O+ ] Example 7. Calculate the pH of the solution that results when 40.0 mL of 0.100 M NaOH is: a) Diluted with 200.0 mL of distilled water. b) Mixed with 20.0 mL of 0.200 M HCl solution. c) Mixed with 20.0 mL of 0.150 M HNO3 Question 14. 100 cm3 of a magnesium hydroxide solution required 4.5 cm3 of sulphuric acid (of concentration 0.100 mol dm-3) for complete neutralisation. [atomic masses: Mg = 24.3, O = 16, H = 1) a) Give the equation for the neutralisation reaction. b) Calculate the moles of sulphuric acid neutralised. c) Calculate the moles of magnesium hydroxide neutralised. d) Calculate the concentration of the magnesium hydroxide in mol dm-3 (molarity). e) Calculate the concentration of the magnesium hydroxide in g cm-3. Question 15. A 0.484 g sample of impure NH4 Br is treated with 25.00 mL of 0.2050 M NaOH and heated to drive o the NH3 . The unreacted NaOH in the reaction mixture after heating required 9.95 mL of 0.0525 M H2 C2 O4 to neutralize. How many grams of NH4 Br were in the original sample? Question 16. Calculate how many mL of 0.450M NaOH solution must be added to 40mL of pH 1.2 HNO3 to neutralise. 8
Acid/base and pH
Ho
riz
en
A 0.4755 g sample containing ammonium oxalate (NH4 )2 C2 O4 and inert compounds has dissolved in water and made alkaline with KOH. The liberated NH3 was distilled into 50 mL of 0.1007 N H2 SO4 . The excess H2 SO4 was back titrated with 11.13 mL of 0.1214 N NaOH. Calculate the %N and %(NH4 )2 C2 O4 in the sample. 2 Note: Oxalate = C2 O4
Question 13.
Ed
uc
e) What % w/w of the tablet is acetylsalicylic acid?
ati
d) How many grams of acetylsalicylic acid are in the tablet?
on
c) Calculate the amount of aspirin that reacted with NaOH.