When a mixture of 10.0 g of acetylene (C 2 H 2 ) and 10.0 g of oxygen (O 2 ) is ignited, the resulting combustion reaction produces CO 2 and H 2 O. (c) How many grams of CO 2 and H 2 O are present after the reaction is complete?

Welcome back everyone in this example, we're told that ethanol is a fuel made from renewable sources and produces less carbon dioxide than traditional fossil fuels. We need to calculate what massive carbon dioxide gas is produced when 13.5 g of ethanol undergoes complete combustion with 45 g of oxygen. So we want to recall that in a combustion reaction, we have our re agent which in this case is ethanol, which we recall is the formula C two H five oh H. It's reacting with oxygen and in combustion reactions, our products are always carbon dioxide and water. And as far as our labels are, ethanol is going to be a liquid and the rest of our re agents are going to have the gaseous label. Our next step is to ensure that this equation is balanced. And when we look at the molds of each of our Atoms on both sides of our equation, we're going to have to place a coefficient of three in front of our 02, a coefficient of two in front of our CO2 and a coefficient of three in front of our water so that we can ensure we have a balanced equation. Now, the prompt is asking us to figure out the mass of carbon dioxide that is produced is essentially asking us for our theoretical yield of our carbon dioxide. So we're going to need to find the theoretical yield of carbon dioxide based on both of our reactant. So, first we're going to find the theoretical yields for the mass of carbon dioxide that is produced. And sorry, that's CO2. So for the massive carbon dioxide that is produced from ethanol. So we want to make note of our molar ratio between our ethanol And our carbon dioxide. And we would get that from our balanced equation where we see we have a coefficient of one in front of ethanol and then we have a coefficient of two in front of our carbon dioxide. So we have a 1-2 molar ratio. So we're going to begin our calculation with our massive ethanol from the prompt given us 13.5 g of ethanol and we want to cancel out our units of grams of ethanol to get two moles. And so we're going to recall our molar mass of ethanol from the periodic table which we see as equal to 46.7 g of ethanol For one mole of ethanol. So now we're able to cancel out our units of grams of ethanol. And now we want to focus on going from moles of ethanol to moles of our product carbon dioxide. And we would get that from our molar ratio, which we noted to the left where we said we have one mole of ethanol for one mole of carbon dioxide. So now we're able to cancel out our moles of ethanol. And now we want to get two g of carbon dioxide. So we would recall our molar mass of carbon dioxide from the periodic table, which we understand for one mole of carbon dioxide is 43.99 g of CO2. So now we can cancel out our units of moles of carbon dioxide And we're left with our final unit being grams of carbon dioxide. And this is going to produce a mass from our ethanol of our carbon dioxide being 25.8 g of CO2. So now we want to compare this to the theoretical yields for our mass of carbon dioxide from our second reactant being oxygen. So making note of our molar ratio between oxygen and carbon dioxide, we see from our balanced equation that that ratio is a two 23 Molar ratio. And so beginning with our mass of oxygen from the prompt, we have a mass of 45.0 g of oxygen. We're going to multiply to get two moles of oxygen by recalling our molar mass from the product table, which we understand is 32 g of oxygen for one mole of oxygen. So now we can cancel out our grams of oxygen. And now we want to get from moles of oxygen. Two moles of C. 02. So, using our ratio that we took note of, we said, we have two moles of oxygen for three moles of C. 02. So now we can cancel out moles of 02 and focus on going from moles of C 022 g of C 02. So we're calling from the periodic table, our molar mass of C. 02. We see that we have one mole of C. 02 for 43.99 grams of C. 02. And now we can cancel our moles of C. 02. We're left with grams of C. 02. And this gives us a mass equal to 41.2 g of C. 02 produced from our re agent oxygen. And because we can see that we have only 25.8 g of C. 02, which is less than 41.2 g of C. 02. We would say that therefore our ethanol which makes that lower amount of carbon dioxide. So our ethanol is our limiting reactant and the reaction the reaction produces the mass of carbon dioxide being 25.8 g of C. 02. And so for our final answer, we have that highlighted in yellow here that our mass of carbon dioxide is 25.8 g produced from our ethanol. So actually we should just highlight the mass of C. 02 produced. So what's highlighted in yellow is our final answer. I hope that everything I explained was clear. If you have any questions, leave them down below and I'll see everyone in the next practice video

Ch.3 - Chemical Reactions and Reaction Stoichiometry

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Brown 15th Edition

Ch.3 - Chemical Reactions and Reaction Stoichiometry

Problem 103c3

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Chapter 3, Problem 103c3

When a mixture of 10.0 g of acetylene (C₂H₂) and 10.0 g of oxygen (O₂) is ignited, the resulting combustion reaction produces CO₂ and H₂O. (c) How many grams of CO₂ and H₂O are present after the reaction is complete?

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Determine the molar mass of acetylene (C2H2) and oxygen (O2).

Calculate the moles of acetylene and oxygen initially present using their respective molar masses.

Write the balanced chemical equation for the combustion of acetylene: 2 C2H2 + 5 O2 ightarrow 4 CO2 + 2 H2O.

Using stoichiometry, determine the limiting reactant by comparing the mole ratio from the balanced equation with the moles of each reactant calculated.

Calculate the moles of CO2 produced based on the moles of the limiting reactant, then convert these moles of CO2 to grams using the molar mass of CO2.

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Key Concepts

Here are the essential concepts you must grasp in order to answer the question correctly.

Stoichiometry

Stoichiometry is the branch of chemistry that deals with the quantitative relationships between the reactants and products in a chemical reaction. It allows us to calculate the amounts of substances consumed and produced in a reaction based on balanced chemical equations. In this case, understanding the stoichiometric ratios of acetylene and oxygen will help determine how much carbon dioxide is produced.

Combustion Reactions

Combustion reactions are exothermic reactions that occur when a substance reacts with oxygen, producing heat and light. In the case of acetylene (C2H2), it combusts in the presence of oxygen to form carbon dioxide (CO2) and water (H2O). Recognizing the products of the combustion of acetylene is essential for calculating the mass of CO2 produced.

Molar Mass

Molar mass is the mass of one mole of a substance, typically expressed in grams per mole (g/mol). It is crucial for converting between grams of a substance and moles, which are used in stoichiometric calculations. For this problem, knowing the molar masses of acetylene, oxygen, carbon dioxide, and water will facilitate the conversion of the initial masses into moles, allowing for the determination of the final mass of CO2 produced.

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Molar Mass Concept

Related Practice

Write a balanced chemical equation for the reaction that occurs when (b) silver(I) oxide decomposes into silver metal and oxygen gas when heated (c) propanol, C₃H₇OH(l) burns in air (d) methyl tert-butyl ether, C₅H₁₂O(l), burns in air.

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Textbook Question

(b) Menthol, the substance we can smell in mentholated cough drops, is composed of C, H, and O. A 0.1005-g sample of menthol is combusted, producing 0.2829 g of CO₂ and 0.1159 g of H₂O. What is the empirical formula for menthol? If menthol has a molar mass of 156 g/mol, what is its molecular formula?

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Textbook Question

Nitrogen 1N22 and hydrogen 1H22 react to form ammonia 1NH32. Consider the mixture of N2 and H2 shown in the accompanying diagram. The blue spheres represent N, and the white ones represent H. (d) If so, how many of which type are left over?

How many N2 molecules are left over?

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Textbook Question

An element X forms an iodide (XI₃) and a chloride (XCl₃). The iodide is quantitatively converted to the chloride when it is heated in a stream of chlorine: 2 XI₃ + 3 Cl₂ → 2 XCl₃ + 3 I₂ If 0.5000 g of XI₃ is treated with chlorine, 0.2360 g of XCl₃ is obtained. (a) Calculate the atomic weight of the element X. (b) Identify the element X.

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Textbook Question

Several brands of antacids use Al1OH23 to react with stomach acid, which contains primarily HCl: Al(OH)₃(s) + HCl(aq) → AlCl₃(aq) + H₂O(l) (b) Calculate the number of grams of HCl that can react with 0.500 g of Al(OH)₃. (c) Calculate the number of grams of AlCl₃ and the number of grams of H₂Oformed when 0.500 g of Al(OH)₃ reacts.

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Textbook Question

Write balanced chemical equations to correspond to each of the following descriptions: (d) When solid mercury(II) nitrate is heated, it decomposes to form solid mercury(II) oxide, gaseous nitrogen dioxide, and oxygen. (e) Copper metal reacts with hot concentrated sulfuric acid solution to form aqueous copper(II) sulfate, sulfur dioxide gas, and water.

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When a mixture of 10.0 g of acetylene (C2H2) and 10.0 g of oxygen (O2) is ignited, the resulting combustion reaction produces CO2 and H2O. (c) How many grams of CO2 and H2O are present after the reaction is complete?

Verified Solution

Key Concepts

Stoichiometry

Combustion Reactions

Molar Mass

When a mixture of 10.0 g of acetylene (C₂H₂) and 10.0 g of oxygen (O₂) is ignited, the resulting combustion reaction produces CO₂ and H₂O. (c) How many grams of CO₂ and H₂O are present after the reaction is complete?