12:06 A *A * ll 68%iZoom vLeave• RECBreakoutRoomsGROUP WORKAn iron ore was analyzed by dissolving a 1.1324 gsample in concentrated HCI.The resulting solution was diluted with water and theiron (III) was precipitated as the hydrous oxideFe,O3.xH,O by the addition of NH3. After filtration andwashing, the residue was ignited at a high temperatureto give 0.5394 g of pure Fe,03 (159.69g/mol).Calculate (a) % Fe (55.847g/mol) and (b) %Fe,04(231.54g/mol) in the sample.•. 28UnmuteStart VideoShareParticipantsMoreII

Answered: Image /qna-images/answer/17f8a7d6-01ab-4658-a942-5dd31a481219.jpg

GROUP WORK An iron ore was analyzed by dissolving a 1.1324 g sample in concentrated HCI. The resulting solution was diluted with water and the iron (III) was precipitated as the hydrous oxide Fe,03.xH,O by the addition of NH3. After filtration and washing, the residue was ignited at a high temperature to give 0.5394 g of pure Fe,O3 (159.69g/mol). Calculate (a) % Fe (55.847g/mol) and (b) % Fe,04(231.54g/mol) in the sample.

GROUP WORK An iron ore was analyzed by dissolving a 1.1324 g sample in concentrated HCI. The resulting solution was diluted with water and the iron (III) was precipitated as the hydrous oxide Fe,03.xH,O by the addition of NH3. After filtration and washing, the residue was ignited at a high temperature to give 0.5394 g of pure Fe,O3 (159.69g/mol). Calculate (a) % Fe (55.847g/mol) and (b) % Fe,04(231.54g/mol) in the sample.

Chemistry

10th Edition

ISBN:9781305957404

Author:Steven S. Zumdahl, Susan A. Zumdahl, Donald J. DeCoste

Publisher:Steven S. Zumdahl, Susan A. Zumdahl, Donald J. DeCoste

Chapter1: Chemical Foundations

Section: Chapter Questions

Problem 1RQ: Define and explain the differences between the following terms. a. law and theory b. theory and...

See similar textbooks

Similar questions

The ferrous ammonium sulfate solution that you tested was prepared by dissolving 40.0 g of solid(NH4)2Fe(SO4)2•6H2O in 1.00 liter of solution. This substance is often impure.a. Calculate the theoretical percent Fe in a pure sample of (NH4)2Fe(SO4)2•6H2O.
3. You start with a 3.865 g sample of a salt which has the formula Cox Cly · zH₂O, where x, y, and z are unknown. After gently heating the salt (as you did in lab with a similar compound), the mass decreased to 2.108 g. The remaining salt was then dissolved in water to form a clear solution. After adding an excess of aluminum metal, the solution fizzed, and a new blueish precipitate formed. After collecting and drying the blue precipitate, its mass was measured to be 0.957 g. Determine the empirical formula of the original compound.
A solution is prepared by dissolving 2.95 g Co2(SO4)3 in enough water to make 1.00 L of red solution; the red color is due to the presence of the cobalt ions. To this solution was added 1.52 g of magnesium powder. A redox reaction occurred, in which the magnesium atoms lose all of their valence electrons and the red color of the solution fades. (a) Write the total reaction that occurs and the net ionic equation for the redox process. (Assume that one of the products of the redox reaction is a metal.) (b) Calculate the initial concentration of cobalt ions, the final concentration of cobalt ions, and the final concentration of magnesium ions.
A 4.908-g sample of a petroleum product was burned in a tube furnace, and the SO2 produced was collected in 2% H2 O2. Reaction: SO2 (g) + H2O2 → H2SO4 A 27.00-mL portion of 0.00826 M NaOH was introduced into the solution of H2 SO4, following which the excess base was back-titrated with 15.15 mL of 0.01165 M HCl. Calculate the sulfur concentration in the sample in parts per million. Concentration ppm
16.792 g of Ba(OH)2 (171.342 g.mol-1) is weighed out and carefully dissolved to make a 500.00cm3 solution. A 25.00 cm3 of this solution is pipette then transfer into a 100.00 mL volumetric flask, filling up to the mark with water. A 40.00 mL sample of this diluted solution is then reacted with a 0.753 M solution of HNO3. Ba(OH)2(aq) + 2HNO3(aq) → Ba(NO3)2(aq) + 2H2O(l) The moles of Ba(OH)2 transferred to the beaker is:
To determine the amount of iron in a dietary supplement, a random sample of 23 tablets weighing a total of 847 g was ground into a fine powder. A 0.5046 g sample of the powdered tablets was dissolved and treated to precipitate the iron as Fe(OH)3(106.95g/mol). The precipitate was collected, rinsed, and ignited to a constant weight as Fe2O3.(159.69g/mol) yielding 0.5445 grams. Report the iron content of the dietary supplement as g FeSO4.7H₂O (277.91g/mol) per tablet. Unit should not be written in the box.
Manganese commonly occurs in nature as a mineral. The extraction of manganese from the carbonite mineral rhodochrosite. involves a two-step process. In the first step, manganese (II) carbonate and oxygen react to form manganese (IV) oxide and carbon dioxide: 2 MNCO,(s)-0,(g)→2 MnO,(s)-2 C0 (g) In the second step, manganese (IV) oxide and aluminum react to form manganese and aluminum oxide: 3 MnO, (s)- 4 Al(s)-3 Mn(s)–2 Al,O,(s) Write the net chemical equation for the production of manganese from manganese (II) carbonate, oxygen and aluminum. Be sure your equation is balanced.
A 57.0 mLmL sample of a 0.108 MM potassium sulfate solution is mixed with 37.0 mLmL of a 0.106 MM lead(II)(II) acetate solution and the following precipitation reaction occurs: K2SO4(aq)+Pb(C2H3O2)2(aq)→2KC2H3O2(aq)+PbSO4(s). Find the theoretical yield and percent yield.
It is known that when mixed in solution aqueous, silver ions will react with chloride ions to form insoluble silver chloride. Upon dissolving a 5.000 g solid mixture of AgNO3 and KNO3 in a concentrated KCl solution, 1.582 grams of a white precipitate, AgCl(s), formed. Determine the percent, by weight, of AgNO3 in the original mixture.
A stock solution of Cu(C2H302)2 has a concentration of 0.270 M. If 6.00 mL of this solution is transferred to a 50.00 mL flask and then diluted with distilled water, what is the final concentration of Cu2+ ions?
A student investigated the stoichiometry of the reaction of zinc (Zn) with HCI solution and reported the following data: When 0.2158 g of Zn reacted with 10.00 mL of 1.000M HCI, 82.062 g of water was displaced. A total of 36.00 mL of 9.501 x 10 M NaOH solution was required to titrate the HCl remaining in the reaction mixture at the end of the reaction. The room temperature was 27.0°C and the barometric pressure was 777 torr. The gram atomic mass of Zn is 65.38 g mol. The density of water at 27.0 °C is 0.9965 g mLand its vapor pressure is 27.0 torr. R = 8.21 X.10 2 L- atm K mol 1 Calculate the number of millimoles of HCI remaining in the flask at the end of the reaction
The reaction of hydroxylamine and Fe(III) produces nitrous oxide, Fe(II) and water: 2H2NOH + 4Fe3+  N2O(g) + 4Fe2+ + 4H+ + H2O. A 50 ml sample of hydroxylamine was treated with Fe(III) solution and the released Fe(II) was titrated with 0.0325 M K2Cr2O7 solution, which consumed 19.83 ml. Calculate the hydroxylamine concentration of the sample.