Ethene is converted to ethane by the reaction C2 H4(g)+H2(g) catalyst⟶ C2 H6(g) C2 H4 flows into

Ethene is converted to ethane by the reaction C2 H4(g)+H2(g)...

Ethene is converted to ethane by the reaction $$\mathrm{C}_{2} \mathrm{H}_{4}(g)+\mathrm{H}_{2}(g) \stackrel{catalyst}{\longrightarrow} \mathrm{C}_{2} \mathrm{H}_{6}(g)$$ $\mathrm{C}_{2} \mathrm{H}_{4}$ flows into a catalytic reactor at $25.0$ atm and $300 .^{\circ} \mathrm{C}$ with a flow rate of $1000 .$ Umin. Hydrogen at $25.0$ atm and $300 .^{\circ} \mathrm{C}$ flows into the reactor at a flow rate of $1500 .$ L/min. If $15.0 \mathrm{kg}$ $\mathrm{C}_{2} \mathrm{H}_{6}$ is collected per minute, what is the percent yield of the reaction?

Key Concepts

Reaction Stoichiometry

This concept involves using the balanced equation to understand the exact mole ratio in which reactants combine and products form. It is fundamental in determining how much of a product can theoretically be formed from given quantities of reactants under ideal conditions.

Limiting Reactant

The limiting reactant is the substance that is completely consumed first in a chemical reaction, thereby determining the maximum amount of product that can be produced. Identifying the limiting reactant is essential for calculating theoretical yields in any reaction.

Molar Flow Conversion

In processes involving gases at known pressures and temperatures, converting flow rates (typically given in volume per time) into moles per time is accomplished using principles from the ideal gas law. This conversion is necessary to relate actual flow rates of reactants to the stoichiometric ratios in the reaction.

Percent Yield

Percent yield is a measure of the efficiency of a reaction, defined as the ratio of the actual yield obtained to the theoretical yield predicted by stoichiometry, multiplied by 100. It quantifies how effectively reactants are converted into the desired product under given reaction conditions.

Transcript

00:02 Hi there.

00:03 In problem number 76, we are reacting ethene with hydrogen to make methane.

00:15 So c2h4 reacts with h2 and produces c2h6.

00:25 These are all gases.

00:27 I'm just leaving the subscripts off for now.

00:30 Just to keep things moving along here.

00:33 And we are told that this is being delivered.

00:37 At a rate of 1 ,000 liters per minute, and the hydrogen is being supplied at a rate of 1 ,500 liters per minute, and experimentally, we are producing 15 kilograms per minute.

01:12 Of the effing.

01:15 Additionally, we know that all of these gases are at 25 atmospheres, and the temperature is 300 degrees celsius, converting that straight away to kelvin by adding 273.

01:38 Gives me 573 kelvin.

01:44 All right.

01:46 This is a percent yield problem.

01:50 So we need to calculate the theoretical.

01:57 Try spelling that again.

02:03 We need to calculate the theoretical amount of c2h6 produced per minute so that we can take experimental, divided by theoretical to give percent yield.

02:14 Another little twist to this problem is it's also a limiting reactant problem.

02:19 But if we look at these two gases that are reacting, they are both reacting at the same temperature and the same pressure.

02:27 Therefore, since the pressure and temperature are constant, the volume can be used to look at to see which one is our limiting reactant.

02:38 Because it is a one -to -one ratio.

02:41 So for every 1 ,000 liters of c2h4 that's reacting, we would need 1 ,000 liters of h2.

02:48 But we have more than that.

02:49 So that means that the h2 is in excess, and the c2h4 is our limiting.

03:00 All right.

03:00 Before i can do stoichiometry with my limiting reactant, i am going to figure out how many, i need to figure out how many moles per minute i have of that, so that then i can go through normal stoichiometry.

03:12 I'm going to use the gas law equation, the ideal gas law, pv equals nrt, but i'm going to solve for number of moles.

03:19 So a number of moles are going to be equal to pv divided by rt...

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