Assume that you have 1.39 mol of H2 and 3.44 mol of N2 How many grams of ammonia (NH ) can you m

Assume that you have 1.39 mol of H2 and 3.44 mol of N2 How...

Assume that you have 1.39 $\mathrm{mol}$ of $\mathrm{H}_{2}$ and 3.44 $\mathrm{mol}$ of $\mathrm{N}_{2}$ How many grams of ammonia (NH_ ) can you make, and how many grams of which reactant will be left over?
$$3 \mathrm{H}_{2}+\mathrm{N}_{2} \longrightarrow 2 \mathrm{NH}_{3}$$

Key Concepts

Mass-Mole Conversion

Mass-mole conversion is the process of converting between the mass of a substance (in grams) and the amount of substance (in moles) using the substance's molar mass. This conversion is essential in stoichiometry to relate the theoretical yield of products to the amounts of reactants used.

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 formed. Once this reactant is exhausted, the reaction cannot proceed further, regardless of the excess of other reactants.

Mole Concept

The mole is a fundamental unit in chemistry used to quantify amounts of a chemical substance. It provides a bridge between the atomic scale and the macroscopic scale, allowing chemists to convert between the number of molecules or atoms and a measurable mass in grams.

Chemical Reaction Stoichiometry

Stoichiometry involves the quantitative relationships between reactants and products in a chemical reaction. It is based on the balanced chemical equation, which shows the proportion of moles of each reactant and product involved in the reaction.

Transcript

00:01 This is a limiting reactant problem, which you know because you were given quantities for both of the reactants.

00:09 So we're told that we start with 1 .39 moles of h2 and 3 .44 moles of n2.

00:24 Actually, i'm going to do this in a different color.

00:31 We don't know which of these is going to run out first.

00:34 Ultimately, we're being asked for how much product we're going to get.

00:37 There's a lot of different ways to approach these problems.

00:40 My personal favorite is for each one or for at least one of them, figure out how many moles of the other reactant are needed.

00:49 So let's pretend.

00:51 Forget that i have this much and two for a second and pretend that i'm going to use all the h2.

00:58 Using multiple -mole ratios, i can see that 1 .39 moles of h2 needs one mole of nitrogen.

01:11 For every three moles of h2 that react.

01:16 Notice that these units cancel, and i find that the 1 .39 needs a third, right, as many moles of n2.

01:28 This is equal to 0 .463.

01:32 So 1 .39 divided by 3 gives you 0 .463 moles of n2.

01:39 In other words, 1 .39 moles of h2, needs 0 .439 moles of n2.

01:53 Now we can take a look at how much we actually had.

02:00 We have much more than that, not but, but we have more than what it's needed.

02:14 So that means that h2 is going to be my limiting reactants because it needed 0 .43 moles of n2 and i had more n2 than i needed, so h2 is going to run out first...

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