The episode recaps foundational terms for chemical calculations.
The episode recaps foundational terms for chemical calculations. Starting by comparing the concept of moles to a dozen (:40) we define moles using Avogadro’s number (1:58). Connecting moles to grams we introduce Molar Mass (2:50) and contrast it to atomic mass (3:30). The episode discusses average atomic mass (4:38) and how the abundance of isotopes can be determined utilizing mass spectroscopy (5:25). Using Magnesium as an example the calculations for average atomic mass are described (6:00).
Question: What volume does 1 mole of a gas occupy? (7:22)
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Hi and welcome to the APsolute Recap: Chemistry Edition. Today’s episode will recap moles, molar mass and mass spectroscopy.
Lets Zoom Out:
Unit 1 - Atomic Structure and Properties
Topic - 1.1 and 1.2
Big idea - Scale, Proportion and Quantity
Introduction:
There are a lot of treats out there that our author Sarah likes, but one of her favorites are donuts. Can you blame her?
Because Sarah is also a person who loves to share baked goods, she orders donuts by the dozen and adds some donut holes too. The term “dozen” is commonly used instead of 12 - no matter if it is 12 donuts or 12 donut holes, which DOES make a huge difference! In chemistry, we have a similar term: mole - and no, I am not talking about the black-fured critter that is every gardener's nightmare. A mole means a certain quantity, 6.022 × 1023 to be more exact. Moles help us to connect the particle number to mass via Molar Mass. Today’s episode will recap moles, molar mass and mass spectroscopy.
Let’s zoom in:
A mole of a substance contains 6.022 x 1023 elementary units, like a dozen is always 12. It is much easier to say: 1 mole of a substance than 6.022 x 1023 atoms. What does elementary units mean? It is a term that indicates that it could be atoms, molecules or even electrons. 1 mole of electrons, for example, means 6.022 x 1023 electrons. This number is called Avogadro’s number - careful, it has nothing to do with an Avocado! It is named after Amedeo Avogardo, an Italian scientist.
I am sure you have heard this joke before: what is heavier, 1 kilogram of feathers or 1 kilogram of rocks? ---- They are both 1 kilogram. From a quantity perspective, we would have way more feathers than rocks, though. We can use this idea the other way around for chemistry. 1 mole of any substance, no matter if helium or argon, has 6.022 x 1023 elementary units, but the weight would be very different! 6.022 x 1023 atoms of helium on average weigh 4.00 grams, whereas 6.022 x 1023 atoms of argon weigh 39.95g. Wait, how do you know the mass?
We need a way to connect the number of particles, the moles and the mass of a substance. In a lab, we cannot count out the atoms and therefore not the moles. So, to be able to convert the moles, which again is pretty much just a word for how many particles, into mass we use Molar Mass. Molar mass indicates how much 1 mol of a substance weighs in grams. It’s unit is grams/mol. So we can use Molar Mass to convert from moles to grams and vice-versa using dimensional analysis. But, let’s take a closer look at how we know how much 6.022 x 1023 atoms weigh by tying this to atomic structure and atomic mass.
Atomic mass is the mass of an atom and, no surprise, measured in “atomic mass unit” or amu. One amu is 1.66 x 10-27 kg. Since atoms are made up of protons, neutrons and electrons, we can, in a simplified approach, add up their masses to get the atomic mass. One proton weighs 1.67 x 10-27 kg, a neutron is a tad heavier, but also around 1.67 x 10-27 kg and an electron weighs 9.1 x 10-31 kg. These are super TINY numbers so we convert them to amu. As you might have realized, a proton and a neutron are approximately equivalent to 1 amu. An electron is about 0.0005 amu. In our simplified approach, we can now say: helium consists of 2 protons and 2 neutrons so its atomic mass is 4 amu. Since electrons are only a fraction of this, they are usually not taken into account.
As you may have noticed, the numerical value for atomic mass and molar mass is the same. But how do we know the atomic mass or molar mass of an element? - We are using our trusted friend, the periodic table. Looking into the periodic table, we see Helium with a mass of 4.0026 and Argon with 39.948. Wait, why are there decimals? Are there pieces of protons and neutrons? Do electrons count after all? No, it is the average atomic mass, taking into account the abundance of isotopes. Remember that isotopes are atoms that have the same number of protons - which determines their “identity” - but different numbers of neutrons. The periodic table tells us the average, taking into account the natural abundance of elements.
To determine the abundance of isotopes in a sample, scientists use mass spectroscopy. The idea is to ionize the atoms forming positive ions, accelerating the ions and then deflecting them using a magnetic field. The degree of deflection is determined by their mass: the lighter they are, the more they are deflected. This degree of deflection will then be recorded. The output graph will show the relative abundance - or relative intensity - vs the mass/charge ratio. Since the charge is +1, this value gives you the mass of the cation.
Let’s use an example: the output of the mass spectroscopy of Magnesium shows that the sample of magnesium consists of 80% of Magnesium-24 - which means the isotope that has 12 protons and 12 neutrons, 10% of Magnesium-25 - which is the isotope with 12 protons and 13 neutrons and 10% of Magnesium-26, which is the isotope with 12 protons and 14 neutrons. To calculate the average atomic mass as a weighted average, we multiply the abundance as decimal with the mass of the isotope for each isotope. Then we add our values together. In our example, this would give us an average atomic mass of 24.3. In AP Chemistry you will only have to be able to interpret outputs of singly charged monatomic ions.
To recap……
A mole of a substance contains 6.022 x 1023 elementary units. The weight of 1 mole of a substance in g is called Molar Mass. We can use molar mass to convert grams to moles and moles to grams - and in extension grams to number of particles and vice-versa. The atomic mass unit is numerically equal to molar mass. The periodic table reports the average atomic mass, which can be experimentally determined using mass spectroscopy.
Coming up next on the Apsolute RecAP Chemistry Edition: Atomic Structure, electron configuration and PES
Today’s Question of the day is about moles and gases.
Question: What volume does 1 mole of a gas occupy?