What is the percent yield of Ho0 if 2.64 moles of HoO were produced when 5.68 mol of H2 was mixed with excess O2? (Hint: Percent yield can be calculated using masses or moles, as long as you are consistent). (Hint: Give your answer to 1 decimal place; if your answer is 51.2%, type in 51.2) 2 H2 + 02- 2 H20

What is the percent yield of Ho0 if 2.64 moles of HoO were produced when 5.68 mol of H2 was mixed with excess O2? (Hint: Percent yield can be calculated using masses or moles, as long as you are consistent). (Hint: Give your answer to 1 decimal place; if your answer is 51.2%, type in 51.2) 2 H2 + 02- 2 H20

Chemistry

9th Edition

ISBN:9781133611097

Author:Steven S. Zumdahl

Publisher:Steven S. Zumdahl

Chapter3: Stoichiometry

Section: Chapter Questions

Problem 172CP: The aspirin substitute, acetaminophen (C8H9O2N), is produced by the following three-step synthesis:...

See similar textbooks

Similar questions

Freon-12, CCl2F2, is prepared from CCl4 by reaction with HF. The other product of this reaction is HCl. Outline the steps needed to determine the percent yield of a reaction that produces 12.5 g of CCl2F2 from 32.9 g of Cd4. Freon-12 has been banned and is no longer used as a refrigerant because it catalyzes the decomposition of ozone and has a very long lifetime in the atmosphere. Determine the percent yield.
The pictures below show a molecular-scale view of a chemical reaction between H2 and CO to produce methanol, CH3OH The box on the left represents the reactants at the instant of mixing, and the box on the right shows what is left once the reaction has gone to completion. D Was there a limiting reactant in this reaction? If so, what was it? Write a balanced chemical equation for this reaction. As usual, your equation should use the smallest possible whole number coefficients for all substances.
In Example 4.2, you found that a particular mixture of CO and H2 could produce 407 g CH3OH. CO(g) + 2H2(g) CH3OH (l) If only 332 g of CH3OH is actually produced, what is the percent yield of the compound?
Fig. 5-5 illustrates a schematic diagram of a combustion device used to analyze organic compounds. Given that a certain amount of a compound containing carbon, hydrogen, and oxygen is combusted in this device, explain how the data relating to the mass of CO2 produced and the mass of H2O produced can be manipulated to determine the empirical formula.
Consider the following diagram, where atom X is represented by a square and atom Y is represented by a circle. (a) Write the equation for the reaction represented by the diagram. (b) If each circle stands for a mole of Y and each square a mole of X, how many moles of X did one start with? How many moles of Y? (c) Using the same representation described in part (b), how many moles of product are formed? How many moles of X and Y are left unreacted?
Small quantities of oxygen gas can be generated in the laboratory by the decomposition of hydrogen peroxide. The unbalanced equation for the reaction is H2O2(uz/)-? H2O(/) + O2(g) Calculate the mass of oxygen produced when 10.00 g of hydrogen peroxide decomposes.
A weighed sample of iron (Fe) is added to liquid bromine (Br2) and allowed to react completely. The reaction produces a single product, which can be isolated and weighed. The experiment was repeated a number of times with different masses of iron but with the same mass of bromine (see graph below). (a) What mass of Br2 is used when the reaction consumes 2.0 g of Fe? (b) What is the mole ratio of Br2 to Fe in the reaction? (c) What is the empirical formula of the product? (d) Write the balanced chemical equation tor the reaction of iron and bromine. (e) What is the name of the reaction product? (f) Which statement or statements best describe the experiments summarized by the graph? (i) When 1.00 g of Fe is added to the Br2, Fe is the limiting reagent. (ii) When 3.50 g of Fe is added to the Br2, there is an excess of Br2. (iii) When 2.50 g of Fe is added to the Br2, both reactanu are used up compietely. (iv) When 2.00 g of Fe is added to the Br2, 10.8 g of product is formed. The percent yield must therefore be 20.0%.
What is meant by a limiting reactant in a particular reaction? In what way is the reaction “limited”? What does it mean to say that one or more of the reactants are present in excess? What happens to a reaction when the limiting reactant is used up?
4.45 If 21 g of H2S is mixed with 38 g of O2 and 31 g of SO2 forms, what is the percentage yield? 2H2S+3O22SO2+2H2O
Quicklime, CaO, is formed when calcium hydroxide is heated. Ca(OH)2(s) → CaO(s) + H2O(ℓ) The theoretical yield is 65.5 g but only 36.7 g quicklime I is produced. Calculate the percent yield.
Cyanogen gas, C2N2, has been found in the gases of outer space. It can react with fluorine to form carbon tetrafluoride and nitrogen trifluoride. C2N2(g)+7F2(g)2CF4(g)+2NF3(g)(a) How many moles of fluorine react with 1.37 mol of cyanogen? (b) How many moles of CF4 are obtained from 13.75 mol of fluorine? (c) How many moles of cyanogen are required to produce 0.8974 mol of NF3? (d) How many moles of fluorine will yield 4.981 mol of nitrogen trifluoride?
One mol of ammonia reacts with 1.00 mol of oxygen to form nitrogen oxide and water according to the reaction 4NH3(g)+5O2(g)4NO(g)+6H2O(l)State which statements are true about the reaction and make the false statements true. (a) All the oxygen is consumed. (b) 4.00 mol NO are produced. (c) 1.50 mol H2O are produced. (d) The description of the experiment does not provide enough information to determine percent yield. (e) Three moles of water are produced for every two moles of NO obtained.