A substance is analyzed and found to contain 85.7 % carbon and 14.3 % hydrogen by weight. A gase

A substance is analyzed and found to contain 85.7 % carbon...

Key Concepts

Isomerism

Isomerism refers to the phenomenon where compounds with the same molecular formula have different structural arrangements of atoms. This concept is crucial when there are multiple valid Lewis structures for a compound, as different isomers (such as chain isomers, position isomers, or cyclic versus acyclic forms) exhibit different chemical and physical properties.

Empirical Formula Determination

This concept involves using the percentage composition by mass of an element in a compound to determine the simplest whole?number ratio of atoms present. By converting mass percentages to moles using atomic masses, one can derive the empirical formula, which represents the simplest ratio of the elements in the substance.

Molar Mass Determination from Gas Density

This concept entails using the measured density of a gas along with its molar volume at standard conditions to calculate the molar mass of the compound. The ideal gas law relates the volume, pressure and temperature of a gas, so knowing that one mole occupies a certain volume (e.g., 22.4 L at STP) allows for the determination of molar mass when the density is given.

Molecular Formula Determination

Once both the empirical formula and the molar mass of the compound are known, the molecular formula can be determined by comparing the empirical formula mass with the actual molar mass. The molecular formula is an integer multiple of the empirical formula that matches the experimentally determined molar mass.

Lewis Structures

Lewis structures are diagrams that depict the bonding between atoms of a molecule and the lone pairs of electrons that may exist. They provide a visual representation of the molecule's structure and allow for the identification of different bonding arrangements (isomers) consistent with the molecular formula derived from the experimental data.

Transcript

00:01 This is now moving into the more advanced problems of the chapter that combined several concepts into one.

00:07 So the end of this problem is still going to be the lewis structures, as most of this chapter has focused on.

00:13 But to start, we're pulling on from previous chapter's experience.

00:16 Firstly, knowing that we have 85 .7 % carbon by mass and 14 .3 % hydrogen.

00:26 Now from that, we can get a empirical formula of our compound, try and figure out what this thing is.

00:34 The density, we're going to manipulate the density to get the molar mass.

00:38 The whole goal is to get an actual formula, a molecular formula of this thing.

00:44 From there, we can try to draw it.

00:47 So let's start with taking percent by weight and turning that into an empirical formula.

00:53 If it's been a little bit, this can be something of a reverse.

00:56 View.

00:57 Because these are mass percents, we can make an assumption that we have 100 grams of this stuff.

01:03 It's not going to make any difference at all later on.

01:06 We're just going to make that assumption.

01:08 The reason we do that is because then we can assume that we have 85 .7 grams of carbon and 14 .3 of carbon, or hydrogen.

01:21 This gives us a number we can actually work with.

01:24 So what we want to do is convert these to moles, and then from that we can get an empirical formula.

01:30 So i've added a periodic table.

01:32 We'll use the valence electron part of the periodic table at the end, but for right now, we're interested in the molar masses.

01:39 So first up, we need that of carbon.

01:43 Carbon has a molar mass of 12 .01, and that's going to be in grams per mole.

01:50 We want the grams to be on the bottom because we need them to cancel.

01:55 Grams cancel with grams.

01:57 Okay.

01:58 And if you divide those two together, you should get 7 .135, and that is moles of carbon.

02:11 We're then going to do the same thing using the hydrogen.

02:16 Molar mass of hydrogen is 1 .008.

02:21 Again, grams per mole.

02:23 We'll put the grams on the bottom, so they cancel with the one mole on top.

02:29 Grams will cancel.

02:32 14 .3 divided by 1 .008 gives you 14 .1 .1 .1.

02:38 1 .9 moles of hydrogen.

02:45 Now, from comparing these, we can get an empirical formula.

02:48 These have the right ratios to each other.

02:51 And maybe you can look at this and easily see, well, there appears to be twice as much hydrogen.

02:55 So it's probably ch2.

02:57 But just to do this completely all the way through, in order to get these numbers easier to read, we divide both of them by the smallest number.

03:11 And what it does is guarantee that one of these will come out as one.

03:15 So that's at least one whole number.

03:18 And the other one will come out as something very similar.

03:22 You get 1 .988, which as long as it's to one decimal point, like it's 0 .9, you can go ahead and round that up to two moles of hydrogen.

03:35 So that means that our empirical formula, meaning our simplified formula, is ch2.

03:48 We're interested in the molecular formula and we'll get there, but now we know it's going to be some multiple of this.

03:53 So maybe it's c -h -2.

03:55 Maybe it's c2h4.

03:58 Maybe it's c3 -h -6.

04:00 We don't know quite enough yet to know which it is, but this is a good start.

04:05 Okay.

04:06 Next, we're going to take advantage of what they were saying about they gave us, the sample has a density of 1 .875, or 1 .87 grams per liter, and one mole of it occupies a volume of 22 .4 liters.

04:23 We're going to take advantage of this to get the molar mass.

04:28 We're going to take our, they said, 1 .87 grams per 1 liter.

04:34 Now think for a second what molar mass, what units those are.

04:38 We've already used molar masses up higher.

04:40 Those are the numbers you pull off the periodic table.

04:42 And they're in grams per mole.

04:44 So if we want to get the molar mass, we need grams divided by moles.

04:50 Well, the density has the grams part.

04:54 What they said about moles and 22 .4 liters can get us the rest of the way like this.

05:08 The liters are going to cancel and you are left in grams divided by moles, which is molar mass.

05:18 Multiply those together...