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Question- When coal is burned, the sulfer it contains is converted into sulfur dioxide. The SO2 is a serious pollutant, so it needs to be removed before it escapes from the stack of a coal fired plant. One way to remove the SO2 is to add limestone, which contains calcium carbonate, CaCO3, to the coal before it is burned. The heat of the burning coal converts the CaCO3 to calcium oxide, CaO. The calcium oxide reacts with the sulfur dioxide in the following reaction: 2CaO+2SO2+O2 > 2CaSO4
The solid calcium sulfate does not escape from the stack as the gaseous sulfur dioxide would. What mass of calcium sulfate forms for each 1.00 Mg of SO2 removed by this technique?
Rework your solution from the beginning and check each step carefully.
Assuming that seawater has a total ion concentration (a.k.a colligative molarity) of 1.10 Mc, calculate how many liters of seawater are needed to produce 25.9 L of fresh water at 20 °C with an applied pressure of 65.0 bar.
The enthalpy change for the oxidation of styrene, C8H8, is measured by calorimetry. C8H8(l) + 10O2(g) ? 8CO2(g) + 4HOH(l) ?rH = -4395.0 kJ/mol How much heat in kJ is required to oxidize 25.0 g of styrene?
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How many glucose molecules does a fat cell, which has a diameter of 2.47 micrometers, contain when its internal glucose concentration is 1.47 mM
analyze the societal implications of using this process
What molarity of aqueous calcium chloride solution can be expected to show the same conductivity as 0.0050 M aluminum sulfate at 12,500 units.
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Compute the pH of the subsequent solutions. In every case ignore the effects of ionic strength as well as assume that a(i)=c(i)
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