3) A company needs to verify the amount of nitrogen in their new "QuikGreen" fertilizer product. To test the nitrogen content, a 5.01 g sample of fertilizer is dissolved in 100.0 mL of water, and powdered aluminum and sodium hydroxide added. The mixture is left to react and then heated to remove any dissolved NH3 which is transferred to a flask containing 50.0 mL of 0.989 M hydrochloric acid where it undergoes a neutralization reaction. The HCI solution is then titrated to methyl red endpoint using 11.74 mL of 0.251 M sodium hydroxide. a) How many grams of nitrate were in the original sample?

3) A company needs to verify the amount of nitrogen in their new "QuikGreen" fertilizer product. To test the nitrogen content, a 5.01 g sample of fertilizer is dissolved in 100.0 mL of water, and powdered aluminum and sodium hydroxide added. The mixture is left to react and then heated to remove any dissolved NH3 which is transferred to a flask containing 50.0 mL of 0.989 M hydrochloric acid where it undergoes a neutralization reaction. The HCI solution is then titrated to methyl red endpoint using 11.74 mL of 0.251 M sodium hydroxide. a) How many grams of nitrate were in the original sample?

Chemistry: Principles and Reactions

8th Edition

ISBN:9781305079373

Author:William L. Masterton, Cecile N. Hurley

Publisher:William L. Masterton, Cecile N. Hurley

Chapter4: Reactions In Aqueous Solution

Section: Chapter Questions

Problem 77QAP: A 300.0-g sample of a solid is made up of a uniform mixture of NaNO3, MgCl2, and BaCl2. A 100.0-g...

See similar textbooks

Similar questions

On Easter Sunday, April 3, 1983, nitric acid spilled from a tank car near downtown Denver, Colorado. The spill was neutralized with sodium carbonate: 2HNO3(aq)+Na2CO3(aq)2NaNO3(aq)+H2O(l)+CO2(g) a. Calculate H for this reaction. Approximately 2.0 104 gal nitric acid was spilled. Assume that the acid was an aqueous solution containing 70.0% HNO3 by mass with a density of 1.42 glcm3. What mass of sodium carbonate was required for complete neutralization of the spill, and what quantity of heat was evolved? (Hf for NaNO3(aq) = 467 kJ/mol) b. According to The Denver Post for April 4, 1983, authorities feared that dangerous air pollution might occur during the neutralization. Considering the magnitude of H, what was their major concern?
4.112 A metallurgical firm wishes to dispose of 1300 gallons of waste sulfuric acid whose molarity is 1.37 M. Before disposal, it will be reacted with calcium hydroxide (slaked lime), which costs $0.23 per pound. (a) Write the balanced chemical equation for this process. (b) Determine the cost that the firm will incur from this use of slaked lime.
Potatoes can be peeled commercially by soaking them in a 3-M to 6-M solution of sodium hydroxide, then removing the loosened skins by spraying them with water. Does a sodium hydroxide solution have a suitable concentration if titration of 12.00 mL of the solution requires 30.6 mL of 1.65 M HCI to reach the end point?
A 300.0-g sample of a solid is made up of a uniform mixture of NaNO3, MgCl2, and BaCl2. A 100.0-g sample of the mixture is dissolved in water and treated with an excess of KOH. The precipitate from the reaction has a mass of 13.47 g. The remaining 200.0-g sample is also dissolved in water and treated with an aqueous solution of AgNO3. The resulting precipitate has a mass of 195.8 g. What are the masses of NaNO3, MgCl2, and BaCl2 in the 300.0-g sample?
On the basis of the general solubility rules given in Table 4.1, predict the identity of the precipitate that forms when the following aqueous solutions are mixed. If no precipitate forms, indicate which rules apply.
The vanadium in a sample of ore is converted to VO2+. The VO2+ ion is subsequently titrated with MnO4 in acidic solution to form V(OH)4+ and manganese(II) ion. The unbalanced titration reaction is MnO4(aq)+VO2++(aq)+H2O(l)V(OH)4+(aq)+Mn2+(aq)+H+(aq) To titrate the solution, 26.45 mL of 0.02250 M MnO4 was required. If the mass percent of vanadium in the ore was 58.1 %, what was the mass of the ore sample? Hint: Balance the titration reaction by the oxidation states method.
Oyster beds in the oceans require chloride ions for growth. The minimum concentration is 8 mg/L (8 parts per million). To analyze for the amount of chloride ion in a 50.0-mL sample of water, you add a few drops of aqueous potassium chromate and then titrate the sample with 25.60 mL of 0.001036 M silver nitrate. The silver nitrate reacts with chloride ion, and, when the ion is completely removed, the silver nitrate reacts with potassium chromate to give a red precipitate. (a) Write a balanced net ionic equation for the reaction of silver nitrate with chloride ions. (b) Write a complete balanced equation and a net ionic equation for the reaction of silver nitrate with potassium chromate, indicating whether each compound is water-soluble or not. (c) What is the concentration of chloride ions in the sample? Is it sufficient to promote oyster growth?
What is the solid that precipitates, and how much of it forms, when an excess of sodium chloride solution is mixed with 10.0 mL of a 2.10 M silver nitrate solution?
In most of its ionic compounds, cobalt is either Co(II) or Co(III). One such compound, containing chloride ion and waters of hydration, was analyzed, and the following results were obtained. A 0.256-g sample of the compound was dissolved in water, and excess silver nitrate was added. The silver chloride was filtered, dried, and weighed, and it bad a mass of 0.308 g. A second sample of 0.416 g of the compound was dissolved in water, and an excess of sodium hydroxide was added. The hydroxide salt was filtered and heated in a flame, forming cobalt(III) oxide. The mass of cobalt(III) oxide formed was 0.145 g. a. What is the percent composition, by mass, of the compound? b. Assuming the compound contains one cobalt ion per formula unit, what is the formula? c. Write balanced equations for the three reactions described.
Magnesium metal (a component of alloys used in aircraft and a reducing agent used in the production of uranium, titanium, and other active metals) is isolated from sea water by the following sequence of reactions: Mg2+(aq)+Ca(OH)2(aq)Mg(OH)2(s)+Ca2+(aq)Mg(OH)2(s)+2HCl(aq)MgCl2(s)+2H2O(l)MgCl2(l)electrolysisMg(s)+Cl2+Cl2(g) Sea water has a density of 1.026 g/cm3 and contains 1272 parts per million of magnesium a5 Mg2+(aq) by mass. What mass, in kilograms, of Ca(OH)2; is required to precipitate 99.9% of the magnesium in 1.00103 L of sea water?
ummarize the simple solubility rules for ionic compounds. How do we use these rules in determining the identity of the solid formed in a precipitation reaction? Give examples including balanced complete and net ionic equations.