Hi there! I’m Jeremy Krug, and welcome to my review of AP Chemistry’s Unit 1 – which covers Atomic Structure and Properties. Please like, subscribe, leave a comment, and share with your classmates! Let’s get started…. You have to be able to convert from moles to grams, and vice versa. To do this, use the atomic mass for an element – or for a compound, the sum of the atomic masses of the atoms in the compound. For example, to convert 10.00 grams of carbon dioxide to moles, you’ll use dimensional analysis – in your conversion factor put 1 mole on top, 44.01 grams on bottom (which is 12.01 plus 16.00 plus 16.00). And when you divide, you get an answer of 0.2272 moles. You should also be able to convert particles to moles. There are 6.022 x 1023 teeny tiny particles in a mole – atoms, molecules, ions, whatever it happens to be for that substance. So if you have those 0.2272 moles of carbon dioxide, you can convert to molecules, by putting 1 mole on the bottom of your conversion factor, and 6.022 x 1023 molecules on the top. When you cancel out the units and multiply, you get 1.368 x 1023 molecules. Another skill is being able to interpret a mass spectrometer graph for an element. For example, a graph like this shows the relative abundance of the isotopes of the element that’s being analyzed. So we have two isotopes, one of them has a mass of 107 atomic mass units, and the other one has a mass of 109 atomic mass units. The 107 isotope is more abundant; about 52% of all the atoms of this element have a mass of 107, while about 48% have a mass of 109. We can do some quick math and estimate that the average atomic mass of this element would be right around 108, maybe a tad less than that. So we can look at the periodic table and confidently say that this is the graph for SILVER. You need to be able to do that for any element’s mass spec graph that you might be given on the exam. You should be able to determine an empirical formula from a compound’s composition data. An empirical formula is the simplest whole number ratio formula for a compound. It’s like taking a formula and reducing it down to lowest terms. So for example, the molecular formula for oxalic acid is H2C2O4, but its empirical formula is HCO2. If we have a substance that contains 40.05 percent sulfur and 59.95 percent oxygen by mass, we would express those percents as grams, like we show here. Then, convert both of those masses to moles, using those elements’ respective atomic masses. Then, when we get the moles, we divide each of them by the smallest of those values. These are the relative subscripts of the elements, so the 1 and 3 tell us that the empirical formula is SO3. Now the law of definite proportions tells us that every sample of SO3, sulfur trioxide, no matter where it comes from, will always have that same proportion, 40.05 percent sulfur, and 59.95 percent oxygen. You need to understand that mixtures are not the same as pure substances. Often in the lab, we’re asked to analyze a sample that has a certain substance we’re interested in. Let’s say it’s potassium chloride. Now that sample might have some other impurities in it. We can weigh the sample, dissolve it in water, then analyze how much potassium or how much chloride is in there using a variety of methods. If we’re asked to compare samples containing multiple chlorides – such as a few different vials containing lithium chloride, sodium chloride, aluminum chloride, and who knows what else, our key here is to focus on the chloride, because it’s the ion they have in common. And if you have a sample of sodium chloride, you can calculate its percent mass and see that a pure sample of sodium chloride should be about 61% chloride. If the sample is only 20% chloride, then you know that the sodium chloride only makes up about 1/3 of what it would be if it were a pure sample; that means you have about 2/3 of the sample making up impurities. One of the key skills in understanding atomic structure is being able to write electron configurations for the elements. For example, the electron configuration for scandium would be 1s2 2s2 2p6 3s2 3p6 4s2 3d1. You need to be very good at writing these. You need to recognize that the electrons in the outermost energy level, or shell, are called valence electrons. In scandium, there are two. The s, p, and d represent sublevels, or subshells, and there are 7 sublevels in scandium. If we’re trying to compare the forces holding the electrons to the nucleus, we use Coulomb’s Law. Essentially, Coulomb’s Law states that there are two factors, CHARGE and DISTANCE, when determining the force holding two charged particles together. The GREATER the CHARGE, the stronger the attractive force, the LOWER the CHARGE, the weaker the attractive force. The GREATER the DISTANCE, the weaker the attractive force, and the LOWER the DISTANCE, the stronger the attractive force. So that means that the electrons that are farthest way from the nucleus, the VALENCE electrons, have the weakest attractions to the nucleus. So they’re the ones that can be removed the easiest. And these core electrons, the ones that are closest to the nucleus, are most difficult to remove. You can identify an atom using photoelectron spectroscopy. If you can write an electron configuration, you can interpret a PES graph like this. All you have to do is label the peaks from left to right with the sublevels in increasing energy, so 1s, 2s, 2p, 3s, 3p, and 4s. And the relative heights of the peaks correspond to number of electrons in each sublevel. So all these s sublevels each have 2, since they’re all the same height. And these two are three times taller, so they must have 6. So you can look at this and see that it ends with 4s2, so that means it’s calcium. Several trends in atomic properties can be predicted by looking at the periodic table. For example, ionization energy and electronegativity generally increase as you move to the right and top of the periodic table. And they decrease as you move toward the left and bottom of the table. Atomic radius is the opposite – atoms are smaller toward the top and right of the periodic table, and they are the largest toward the bottom and left. The periodic table helps predict these trends, but they DO NOT EXPLAIN THEM! Generally speaking, if you are asked for an explanation comparing atoms that are across from each other, left and right, their differences are due to a greater effective nuclear charge on the right, and lower effective nuclear charge on the left. When comparing atoms that are above or below each other, their differences are due to a greater distance of valence electrons from the nucleus toward the bottom of the table, and a lower distance from the valence electrons to the nucleus for the atom at the top. When you’re looking at ions, generally speaking, the more positively charged an ion, the smaller it is, and the more negatively charged the ion it is, the larger it will be. That’s because of Coulomb’s Law – an ion with more electrons than protons will allow the electrons to repel each other more and will be larger; an ion with more protons than electrons will have that nucleus pulling in those electrons as tightly as possible. You need to be comfortable with the number of valence electrons in atoms, and how this affects the compounds they form. There are specific patterns that help us to see how many valence electrons an atom has, and as you can see here, the group that an element is in tells us how many valence electrons it has. The octet rule helps us to predict the charge that ions of these atoms will have. So ions from Group 1 tend to have a +1 charge, Group 2 will be +2, and Group 13 will be +3. Ions from Group 17 tend to take a -1 charge, Group 16 will be -2, and Group 15 will be -3. So an ionic compound made up of magnesium and chloride will have +2 and -1 charges. So its formula would be MgCl2. And likewise, aluminum has a +3 charge, while sulfide has a -2 charge. So the formula for aluminum sulfide would be Al2S3. That was 10 minutes, and that’s Unit 1! I’m Jeremy Krug, join me soon for my Unit 2 review in 10 minutes, so we can review some more AP Chemistry together!